Edexcel Triple Science: Chemistry — Recall & Retrieval

1.1 Describe how the Dalton model of an atom has changed over time because of the discovery of subatomic particles

1.

How did Dalton's original model describe the structure of an atom?

A Dalton atom was a solid, indivisible sphere with no internal structure.

2.

Which subatomic particle was discovered that showed atoms were not indivisible?

The electron.

3.

How did the discovery of the electron change Dalton's model of the atom?

It showed that atoms contain smaller negatively charged particles, so atoms are divisible.

4.

Which subatomic particles were discovered after the electron?

Protons and neutrons.

5.

How did the discovery of the nucleus change the model of the atom?

It showed that most of the atom is empty space surrounding a small, dense, positively charged nucleus.

6.

Why is the modern atomic model different from Dalton's original model?

New experimental evidence led to the discovery of subatomic particles and a more detailed model of the atom.

1.2 Describe the structure of an atom as a nucleus containing protons and neutrons, surrounded by electrons in shells

1.

What three types of subatomic particle make up an atom?

Protons, neutrons and electrons.

2.

Which subatomic particles are found in the nucleus of an atom?

Protons and neutrons.

3.

Where are electrons located in an atom?

In shells around the nucleus.

4.

What charge does the nucleus of an atom have?

Positive.

5.

Why does an atom have no overall electrical charge?

The positive charge of the protons balances the negative charge of the electrons.

6.

How are electrons arranged around the nucleus of an atom?

Electrons occupy shells at different distances from the nucleus.

1.3 Recall the relative charge and relative mass of: a a proton b a neutron c an electron

1.

What is the relative charge of a proton?

+1

2.

What is the relative mass of a proton?

1

3.

What is the relative charge of a neutron?

0

4.

What is the relative mass of a neutron?

1

5.

What is the relative charge of an electron?

-1

6.

What is the relative mass of an electron?

1/1840

1.4 Explain why atoms contain equal numbers of protons and electrons

1.

Why does a neutral atom contain equal numbers of protons and electrons?

A neutral atom has no overall charge, so the positive and negative charges must balance.

2.

What charge does a proton have?

+1

3.

What charge does an electron have?

-1

4.

What would happen to the overall charge of an atom if it gained an electron?

It becomes negatively charged.

5.

What would happen to the overall charge of an atom if it lost an electron?

It becomes positively charged.

6.

How do equal numbers of protons and electrons make an atom electrically neutral?

Each proton contributes +1 and each electron contributes -1, so equal numbers give an overall charge of 0.

1.5 Describe the nucleus of an atom as very small compared to the overall size of the atom

1.

How does the size of an atomic nucleus compare with the overall size of an atom?

The nucleus is extremely small compared with the whole atom.

2.

What proportion of the overall size of an atom is occupied by its nucleus?

Only a tiny fraction of the atom's overall volume.

3.

Why is most of an atom considered to be empty space?

The nucleus occupies very little of the atom's volume.

4.

Which part of an atom occupies most of its overall volume?

The electron shells.

5.

How does the size of the nucleus compare with the distance of the electrons from the nucleus?

The nucleus is much smaller than the distances of the electrons from it.

6.

What does the small size of the nucleus tell us about the structure of an atom?

It shows that most of the atom is empty space.

1.6 Recall that most of the mass of an atom is concentrated in the nucleus

1.

Where is most of the mass of an atom concentrated?

In the nucleus.

2.

Which two subatomic particles account for almost all of an atom's mass?

Protons and neutrons.

3.

Why does the electron contribute very little to the mass of an atom?

An electron has a very small relative mass.

4.

How does the mass of a proton compare with the mass of an electron?

A proton has a mass of about 1840 times that of an electron.

5.

How does the mass of a neutron compare with the mass of an electron?

A neutron has a mass of about 1840 times that of an electron.

6.

Why is the nucleus much more massive than the surrounding electron shells?

The nucleus contains the much more massive protons and neutrons.

1.7 Recall the meaning of the term mass number of an atom

1.

What does the mass number of an atom represent?

The total number of protons and neutrons in the nucleus.

2.

How is the mass number of an atom calculated from its subatomic particles?

Mass number = number of protons + number of neutrons.

3.

Which two types of subatomic particle are included in the mass number?

Protons and neutrons.

4.

Why are electrons not included when calculating the mass number?

Electrons have negligible mass compared with protons and neutrons.

5.

How can the number of neutrons in an atom be calculated from its mass number and atomic number?

Number of neutrons = mass number − atomic number.

6.

What information does the mass number give about the nucleus of an atom?

It gives the total number of protons and neutrons in the nucleus.

1.8 Describe atoms of a given element as having the same number of protons in the nucleus and that this number is unique to that element

1.

What determines which element an atom belongs to?

The number of protons in its nucleus.

2.

What is unique about the number of protons in atoms of each element?

Each element has a unique number of protons.

3.

Why do all atoms of the same element have the same atomic number?

The number of protons defines the element.

4.

How can the number of protons identify an element?

The atomic number gives the number of protons.

5.

What would happen to an atom's identity if the number of protons in its nucleus changed?

It would become a different element.

6.

How does the number of protons distinguish one element from another?

Different elements have different numbers of protons.

1.9 Describe isotopes as different atoms of the same element containing the same number of protons but different numbers of neutrons in their nuclei

1.

What is an isotope?

Atoms of the same element with the same number of protons but different numbers of neutrons.

2.

Why do isotopes of the same element have the same chemical identity?

They contain the same number of protons.

3.

What is the same in the nuclei of all isotopes of an element?

The number of protons.

4.

What is different in the nuclei of isotopes of the same element?

The number of neutrons.

5.

How does changing the number of neutrons affect the mass number of an isotope?

It changes the mass number.

6.

Why can isotopes of the same element have different masses?

They contain different numbers of neutrons.

1.10 Calculate the numbers of protons, neutrons and electrons in atoms given the atomic number and mass number

1.

How can the number of protons in an atom be determined from its atomic number?

Number of protons = atomic number.

2.

How can the number of electrons in a neutral atom be determined from its atomic number?

Number of electrons = atomic number for a neutral atom.

3.

How can the number of neutrons in an atom be calculated from its mass number and atomic number?

Number of neutrons = mass number − atomic number.

4.

How many protons, neutrons and electrons are in an atom with an atomic number of 8 and a mass number of 16?

8 protons, 8 neutrons, 8 electrons.

5.

How many protons, neutrons and electrons are in an atom with an atomic number of 11 and a mass number of 23?

11 protons, 12 neutrons, 11 electrons.

6.

How many protons, neutrons and electrons are in an atom with an atomic number of 17 and a mass number of 35?

17 protons, 18 neutrons, 17 electrons.

1.11 Explain how the existence of isotopes results in relative atomic masses of some elements not being whole numbers

1.

Why are the relative atomic masses of some elements not whole numbers?

Most elements exist as mixtures of isotopes with different masses.

2.

How does the existence of isotopes affect the relative atomic mass of an element?

The relative atomic mass is a weighted mean of the masses of the isotopes.

3.

Why is the relative atomic mass of an element a weighted average?

Different isotopes occur in different abundances, so the average is weighted according to abundance.

4.

How does the abundance of each isotope affect the relative atomic mass?

More abundant isotopes have a greater effect on the relative atomic mass.

5.

Why does the relative atomic mass of chlorine differ from the mass number of either of its common isotopes?

Chlorine is a mixture of isotopes with different masses, so its relative atomic mass is between the mass numbers of its common isotopes.

6.

Why is the relative atomic mass of an element usually different from the mass number of an individual atom?

The relative atomic mass is an average for all naturally occurring isotopes, whereas the mass number refers to one atom.

1.12 Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes

1.

How is the relative atomic mass of an element calculated from its isotopes?

Relative atomic mass = Σ(isotope mass × abundance) / Σ(abundances)

2.

What information about isotopes is needed to calculate relative atomic mass?

The mass of each isotope and its relative abundance.

3.

Why must isotope masses be multiplied by their relative abundances when calculating relative atomic mass?

To account for how common each isotope is.

4.

How would you calculate the relative atomic mass of an element containing two isotopes with known masses and abundances?

Multiply each isotope mass by its abundance, add the results, then divide by the total abundance.

5.

An element has two isotopes with masses of 10 and 11 in abundances of 20% and 80%; what is its relative atomic mass?

(10 × 20 + 11 × 80) / 100 = 10.8

6.

An element has isotopes of masses 35 and 37 with abundances of 75% and 25%; what is its relative atomic mass?

(35 × 75 + 37 × 25) / 100 = 35.5

1.13 Describe how Dmitri Mendeleev arranged the elements, known at that time, in a periodic table by using properties of these elements and their compounds

1.

How did Dmitri Mendeleev arrange the elements in his periodic table?

He arranged elements in order of increasing relative atomic mass while grouping elements with similar properties.

2.

Which properties did Mendeleev use when arranging the elements?

The properties of the elements and their compounds.

3.

Why did Mendeleev consider the properties of compounds when developing his periodic table?

To identify recurring patterns in chemical behaviour.

4.

What did Mendeleev notice about the properties of elements when arranging them?

Properties repeated periodically.

5.

How did Mendeleev's arrangement allow elements with similar properties to be grouped together?

Similar elements were placed in the same groups.

6.

Why was Mendeleev's periodic table different from simply arranging elements alphabetically?

It grouped elements according to repeating chemical patterns and properties.

1.14 Describe how Dmitri Mendeleev used his table to predict the existence and properties of some elements not then discovered

1.

How did Mendeleev use gaps in his periodic table?

He left gaps where he predicted undiscovered elements should occur.

2.

Why did Mendeleev leave gaps in his periodic table?

To allow for elements that had not yet been discovered.

3.

How could Mendeleev predict the properties of undiscovered elements?

By using the properties and trends of neighbouring elements.

4.

What did the successful discovery of predicted elements provide evidence for?

Strong evidence that his periodic table was valid.

5.

Why was Mendeleev able to predict the properties of elements that had not yet been discovered?

The repeating patterns allowed properties to be predicted from surrounding elements.

6.

How did the discovery of elements that matched Mendeleev's predictions support his periodic table?

The discoveries of elements with properties matching his predictions supported his table.

1.15 Explain that Dmitri Mendeleev thought he had arranged elements in order of increasing relative atomic mass but this was not always true because of the relative abundance of isotopes of some pairs of elements in the periodic table

1.

Why did Mendeleev initially think he had arranged elements in order of increasing relative atomic mass?

He generally observed that relative atomic mass increased as he arranged the elements.

2.

Why is the order of some elements in the periodic table not exactly the order of their relative atomic masses?

Isotopes can give an element a relative atomic mass that does not fit the simple atomic-mass order.

3.

How can isotopes affect the relative atomic mass of an element?

Different isotopes have different masses and abundances.

4.

Why can an element with a higher atomic number have a lower relative atomic mass than the element before it?

The element with the higher atomic number can have isotopes that give it a lower average relative atomic mass.

5.

How did the existence of isotopes explain anomalies in Mendeleev's arrangement?

Isotopes explained why some elements had to be placed out of strict relative-mass order to preserve similar properties.

6.

Why is atomic number now used instead of relative atomic mass to order the modern periodic table?

Atomic number corresponds to proton number and gives the correct periodic order.

1.16 Explain the meaning of atomic number of an element in terms of position in the periodic table and number of protons in the nucleus

1.

What does the atomic number of an element represent?

The number of protons in the nucleus.

2.

How is an element's atomic number related to the number of protons in its nucleus?

Atomic number = number of protons.

3.

How does atomic number determine the position of an element in the periodic table?

It determines the element's position in increasing atomic number.

4.

Why does each element have a unique atomic number?

Each element has a unique number of protons.

5.

What is the atomic number of an element containing 12 protons?

12

6.

How can the atomic number be used to identify an element?

Match the atomic number to the element in the periodic table.

1.17 Describe that in the periodic table a elements are arranged in order of increasing atomic number, in rows called periods b elements with similar properties are placed in the same vertical columns called groups

1.

How are elements arranged across the modern periodic table?

In order of increasing atomic number.

2.

What are the horizontal rows of the periodic table called?

Periods.

3.

What are the vertical columns of the periodic table called?

Groups.

4.

Why are elements with similar properties placed in the same group?

They have similar outer-electron arrangements and therefore similar chemical properties.

5.

What does an element's period tell you about its electron arrangement?

It gives the number of occupied electron shells.

6.

What does an element's group tell you about the elements with which it has similar properties?

Elements in the same group have similar outer-shell electron arrangements and therefore similar properties.

1.18 Identify elements as metals or non-metals according to their position in the periodic table, explaining this division in terms of the atomic structures of the elements

1.

Where are metals generally found in the periodic table?

Mainly on the left and centre.

2.

Where are non-metals generally found in the periodic table?

Mainly on the right.

3.

How does the atomic structure of metals differ from that of non-metals?

Metals generally have fewer outer electrons and tend to lose electrons; non-metals generally have more outer electrons and tend to gain or share electrons.

4.

Why are elements on opposite sides of the periodic table generally different in their properties?

Their electron arrangements are different.

5.

How can an element's position in the periodic table be used to predict whether it is a metal or non-metal?

Use its position relative to the metal/non-metal boundary.

6.

How is the distinction between metals and non-metals related to their electron arrangements?

Their outer-shell electron arrangements determine how they gain, lose or share electrons.

1.19 Predict the electronic configurations of the first 20 elements in the periodic table as diagrams and in the form, for example 2.8.1

1.

How are electrons arranged in shells around the nucleus of an atom?

Electrons fill shells starting with the shell closest to the nucleus.

2.

What is the maximum number of electrons that can occupy the first electron shell?

2

3.

What is the maximum number of electrons that can occupy the second electron shell for the first 20 elements?

8

4.

What is the electronic configuration of an atom of sodium?

2.8.1

5.

What is the electronic configuration of an atom of chlorine?

2.8.7

6.

What is the electronic configuration of an atom of calcium?

2.8.8.2

1.20 Explain how the electronic configuration of an element is related to its position in the periodic table

1.

How is the number of occupied electron shells related to an element's period?

The number of occupied shells gives the period.

2.

How is the number of electrons in the outer shell related to an element's group?

For the first 20 elements, the number of electrons in the outer shell corresponds to the group number for Groups 1–7.

3.

Why do elements in the same group have similar chemical properties?

They have the same number of outer-shell electrons.

4.

How can an element's electronic configuration be used to determine its position in the periodic table?

Count the occupied shells for the period and the outer-shell electrons for the group.

5.

What does the electronic configuration 2.8.1 tell you about an element's position in the periodic table?

Group 1, period 3.

6.

What does the electronic configuration 2.8.7 tell you about an element's group and period?

Group 7, period 3.

1.21 Explain how ionic bonds are formed by the transfer of electrons between atoms to produce cations and anions, including the use of dot and cross diagrams

1.

How is an ionic bond formed between two atoms?

By transfer of electrons from one atom to another, forming oppositely charged ions that attract.

2.

Which type of atom loses electrons when forming an ionic bond?

Metals.

3.

Which type of atom gains electrons when forming an ionic bond?

Non-metals.

4.

What is a cation and how is it formed?

A positively charged ion formed when an atom loses electrons.

5.

What is an anion and how is it formed?

A negatively charged ion formed when an atom gains electrons.

6.

How can a dot-and-cross diagram show the formation of an ionic bond?

They show the outer electrons of the atoms and the transferred electron(s) using different symbols.

1.22 Recall that an ion is an atom or group of atoms with a positive or negative charge

1.

What is an ion?

An atom or group of atoms with an overall positive or negative charge.

2.

How does an atom become a positively charged ion?

By losing one or more electrons.

3.

How does an atom become a negatively charged ion?

By gaining one or more electrons.

4.

What is the difference between a cation and an anion?

A cation is positive; an anion is negative.

5.

Why does an ion have an overall electrical charge?

The number of protons and electrons is unequal.

6.

Can an ion consist of more than one atom?

Yes, a polyatomic ion can contain more than one atom.

1.23 Calculate the numbers of protons, neutrons and electrons in simple ions given the atomic number and mass number

1.

How can the number of protons in a simple ion be determined from its atomic number?

Number of protons = atomic number.

2.

How can the number of neutrons in a simple ion be calculated from its mass number and atomic number?

Number of neutrons = mass number − atomic number.

3.

How can the number of electrons in a positive ion be calculated from its atomic number and charge?

Electrons = atomic number − positive charge.

4.

How can the number of electrons in a negative ion be calculated from its atomic number and charge?

Electrons = atomic number + magnitude of negative charge.

5.

How many protons, neutrons and electrons are in Na⁺ with an atomic number of 11 and a mass number of 23?

11 protons, 12 neutrons, 10 electrons.

6.

How many protons, neutrons and electrons are in O²⁻ with an atomic number of 8 and a mass number of 16?

8 protons, 8 neutrons, 10 electrons.

1.24 Explain the formation of ions in ionic compounds from their atoms, limited to compounds of elements in groups 1, 2, 6 and 7

1.

How do Group 1 atoms form ions?

Lose 1 electron to form 1+ ions.

2.

How do Group 2 atoms form ions?

Lose 2 electrons to form 2+ ions.

3.

How do Group 6 atoms form ions?

Gain 2 electrons to form 2− ions.

4.

How do Group 7 atoms form ions?

Gain 1 electron to form 1− ions.

5.

Why do Group 1 and Group 2 atoms lose electrons when forming ions?

They have 1 or 2 outer electrons and lose them to obtain a full outer shell.

6.

Why do Group 6 and Group 7 atoms gain electrons when forming ions?

They need 2 or 1 electrons to complete their outer shell.

1.25 Explain the use of the endings –ide and –ate in the names of compounds

1.

What does the ending "-ide" generally indicate in the name of a compound?

Usually a compound containing a simple anion with no oxygen in the name, such as chloride or oxide.

2.

What does the ending "-ate" generally indicate in the name of a compound?

Usually a compound containing a polyatomic oxygen-containing ion, such as nitrate, carbonate or sulfate.

3.

What is the difference between an oxide and a nitrate in terms of their names?

Oxide contains O²−; nitrate contains NO3−.

4.

Which ending is used for compounds containing a simple non-metal ion such as chloride?

-ide

5.

Which ending is used in compounds containing ions such as nitrate, carbonate and sulfate?

-ate

6.

How can the ending of an ionic compound's name help identify the type of negative ion present?

It can indicate whether the negative ion is a simple ion or an oxygen-containing polyatomic ion.

1.26 Deduce the formulae of ionic compounds (including oxides, hydroxides, halides, nitrates, carbonates and sulfates) given the formulae of the constituent ions

1.

How can the formula of an ionic compound be deduced from the charges of its constituent ions?

Combine the ions in the simplest whole-number ratio that gives overall charge 0.

2.

Why must an ionic compound have an overall charge of zero?

Positive and negative charges must balance.

3.

What is the formula of the ionic compound formed from Na⁺ and Cl⁻?

NaCl

4.

What is the formula of the ionic compound formed from Mg²⁺ and Cl⁻?

MgCl2

5.

What is the formula of the ionic compound formed from Ca²⁺ and OH⁻?

Ca(OH)2

6.

What is the formula of the ionic compound formed from Al³⁺ and SO₄²⁻?

Al2(SO4)3

1.27 Explain the structure of an ionic compound as a lattice structure: a consisting of a regular arrangement of ions b held together by strong electrostatic forces (ionic bonds) between oppositely-charged ions

1.

What is the structure of an ionic compound?

A giant ionic lattice.

2.

What is meant by a lattice structure in an ionic compound?

A regular three-dimensional arrangement of ions.

3.

How are the ions arranged within an ionic lattice?

Positive and negative ions alternate in a regular arrangement.

4.

What holds oppositely charged ions together in an ionic lattice?

Strong electrostatic attractions between oppositely charged ions.

5.

Why does an ionic lattice contain both positive and negative ions?

Electrical neutrality requires equal total positive and negative charge.

6.

Why are the electrostatic forces between ions in an ionic lattice strong?

The ions have full charges and strong electrostatic attraction acts between neighbouring oppositely charged ions.

1.28 Explain how a covalent bond is formed when a pair of electrons is shared between two atoms

1.

How is a covalent bond formed?

When two atoms share a pair of electrons.

2.

What is shared between two atoms in a covalent bond?

A pair of electrons.

3.

Why do atoms form covalent bonds?

To obtain a more stable electron arrangement.

4.

Which types of atoms commonly form covalent bonds?

Non-metal atoms.

5.

How does sharing electrons help atoms achieve a more stable electron arrangement?

Sharing electrons helps fill their outer electron shells.

6.

How can a dot-and-cross diagram represent a covalent bond?

Show one shared pair of electrons between the bonded atoms.

1.29 Recall that covalent bonding results in the formation of molecules

1.

What type of bonding results in the formation of molecules?

Covalent bonding.

2.

What is a molecule?

A group of atoms joined by covalent bonds.

3.

How are atoms held together within a covalent molecule?

By shared pairs of electrons.

4.

Why are covalent substances able to form molecules?

Atoms share electrons and form discrete groups of bonded atoms.

5.

What type of particles make up a simple molecular covalent substance?

Molecules.

6.

How does covalent bonding differ from ionic bonding in terms of the particles formed?

Ionic bonding forms a giant ionic lattice; covalent bonding can form molecules.

1.30 Recall the typical size (order of magnitude) of atoms and small molecules

1.

What is the typical order of magnitude of the size of an atom?

About 10⁻¹⁰ m.

2.

What is the approximate size of an atom in metres?

About 10⁻¹⁰ m.

3.

What is meant by the order of magnitude of a particle's size?

The power of ten that gives the approximate size of the particle.

4.

How does the size of a small molecule compare with the size of an atom?

They are of a similar order of magnitude.

5.

Why are atoms and small molecules too small to see with the naked eye?

They are far too small to resolve with the naked eye.

6.

How can standard form be used to represent the size of atoms and small molecules?

For example, 1 × 10⁻¹⁰ m.

1.31 Explain the formation of simple molecular, covalent substances, using dot and cross diagrams, including: a hydrogen b hydrogen chloride c water d methane e oxygen f carbon dioxide

1.

How is the covalent bond in a hydrogen molecule formed?

Each hydrogen atom shares one electron to form one shared pair.

2.

How is the covalent bond in a hydrogen chloride molecule formed?

Hydrogen and chlorine share one pair of electrons.

3.

How are the covalent bonds in a water molecule formed?

Oxygen shares one pair with each hydrogen atom, forming two O-H bonds.

4.

How are the covalent bonds in a methane molecule formed?

Carbon shares one pair of electrons with each of four hydrogen atoms, forming four C-H bonds.

5.

How is the double covalent bond in an oxygen molecule represented?

The two oxygen atoms share two pairs of electrons, forming a double bond.

6.

How are the covalent bonds in a carbon dioxide molecule represented?

Carbon shares two pairs of electrons with each oxygen atom, forming two C=O double bonds.

1.32 Explain why elements and compounds can be classified as: a ionic b simple molecular (covalent) c giant covalent d metallic and how the structure and bonding of these types of substances results in different physical properties, including relative melting point and boiling point, relative solubility in water and ability to conduct electricity (as solids and in solution)

1.

What are the four main types of substance based on their structure and bonding?

Ionic, simple molecular covalent, giant covalent and metallic.

2.

How does the structure and bonding of ionic substances affect their melting and boiling points?

Strong electrostatic forces between ions require a lot of energy to overcome, giving high melting and boiling points.

3.

How does the structure and bonding of simple molecular covalent substances affect their melting and boiling points?

Only weak intermolecular forces need to be overcome, giving low melting and boiling points.

4.

How does the structure and bonding of giant covalent substances affect their physical properties?

Strong covalent bonds throughout the giant structure give high melting points and generally poor electrical conductivity, apart from structures such as graphite and graphene.

5.

How does metallic bonding allow metals to conduct electricity as solids?

Delocalised electrons can move through the metal lattice.

6.

How do the structures and bonding of different types of substances affect their solubility in water and ability to conduct electricity?

Different bonding and structures determine the strength of forces, particle mobility and therefore melting point, solubility and electrical conductivity.

1.33 Explain the properties of ionic compounds limited to: a high melting points and boiling points, in terms of forces between ions b whether or not they conduct electricity as solids, when molten and in aqueous solution

1.

Why do ionic compounds generally have high melting and boiling points?

Strong electrostatic forces between oppositely charged ions require a lot of energy to overcome.

2.

What type of forces hold the ions together in an ionic compound?

Strong electrostatic forces of attraction.

3.

Why do solid ionic compounds not conduct electricity?

The ions are fixed in position and cannot move.

4.

Why do molten ionic compounds conduct electricity?

The ions are free to move and carry charge.

5.

Why do ionic compounds conduct electricity when dissolved in water?

The ions are separated and can move through the solution.

6.

What happens to the ions in an ionic compound when the compound is molten or dissolved in water?

The ions become mobile when molten or dissolved.

1.34 Explain the properties of typical covalent, simple molecular compounds limited to: a low melting points and boiling points, in terms of forces between molecules (intermolecular forces) b poor conduction of electricity

1.

Why do simple molecular covalent substances generally have low melting and boiling points?

Only weak intermolecular forces need to be overcome.

2.

What forces have to be overcome when a simple molecular substance melts or boils?

Intermolecular forces.

3.

Why are intermolecular forces weaker than the covalent bonds within molecules?

Covalent bonds are much stronger than intermolecular forces.

4.

Why do simple molecular covalent substances generally not conduct electricity?

They contain no mobile charged particles.

5.

Why does having covalent bonds not allow simple molecular substances to conduct electricity?

The electrons are held in covalent bonds and are not free to move through the substance.

6.

Why does a simple molecular covalent substance remain a poor conductor when melted?

Melting does not produce enough mobile charged particles to conduct electricity.

1.35 Recall that graphite and diamond are different forms of carbon and that they are examples of giant covalent substances

1.

What are diamond and graphite?

Giant covalent forms of carbon.

2.

Why are diamond and graphite described as different forms of carbon?

They are allotropes containing only carbon atoms arranged differently.

3.

What type of substances are diamond and graphite?

Giant covalent substances.

4.

What is meant by an allotrope of carbon?

A different structural form of the same element.

5.

Which two giant covalent allotropes of carbon are specified in the Edexcel GCSE Chemistry specification?

Diamond and graphite.

6.

Why can diamond and graphite have different properties even though both contain only carbon atoms?

Their atoms are arranged and bonded differently.

1.36 Describe the structures of graphite and diamond

1.

What is the structure of diamond?

Each carbon atom is covalently bonded to four other carbon atoms in a giant three-dimensional structure.

2.

How many covalent bonds does each carbon atom form in diamond?

Four.

3.

What is the structure of graphite?

Layers of carbon atoms arranged in hexagons.

4.

How are carbon atoms arranged in the layers of graphite?

Carbon atoms form hexagonal layers.

5.

How many covalent bonds does each carbon atom form in graphite?

Three.

6.

How are the layers in graphite held together?

Weak forces between the layers.

1.37 Explain, in terms of structure and bonding, why graphite is used to make electrodes and as a lubricant, whereas diamond is used in cutting tools

1.

Why can graphite conduct electricity?

Each carbon atom has one delocalised electron that can move through the structure.

2.

Why is graphite suitable for use as an electrode?

It conducts electricity and is stable at high temperatures.

3.

Why can graphite be used as a lubricant?

Weak forces between layers allow the layers to slide over one another.

4.

Why can graphite layers slide over one another?

Only weak forces act between the layers.

5.

Why is diamond suitable for use in cutting tools?

It has a rigid giant structure with strong covalent bonds in all directions.

6.

Why is diamond harder than graphite?

Diamond has strong covalent bonds throughout its three-dimensional structure, whereas graphite has weak forces between layers.

1.38 Explain the properties of fullerenes including C60 and graphene in terms of their structures and bonding

1.

What is a fullerene?

Carbon structures made from molecules containing carbon atoms arranged in hollow cages or tubes.

2.

What is the structure of C60?

A spherical cage containing 60 carbon atoms.

3.

What is graphene?

A single layer of carbon atoms arranged in a hexagonal lattice.

4.

Why can graphene conduct electricity?

It contains delocalised electrons that can move through the structure.

5.

Why is graphene strong despite being only one atom thick?

Strong covalent bonds connect the carbon atoms throughout the sheet.

6.

How does the structure of graphene differ from the structure of C60?

Graphene is a single flat layer; C60 is a spherical molecular structure.

1.39 Describe, using poly(ethene) as the example, that simple polymers consist of large molecules containing chains of carbon atoms

1.

What is a polymer?

A large molecule made from many repeating units.

2.

What is the structure of a simple polymer such as poly(ethene)?

Long chains of carbon atoms with hydrogen atoms attached.

3.

What type of atoms form the main chain in poly(ethene)?

Carbon atoms.

4.

Why are polymers described as large molecules?

They contain many repeating units joined together.

5.

How are the carbon atoms arranged in the chain of poly(ethene)?

They form long, continuous chains.

6.

What monomer is used to produce poly(ethene)?

Ethene.

1.40 Explain the properties of metals, including malleability and the ability to conduct electricity

1.

What is meant by the malleability of a metal?

The ability to be hammered or bent into shape without breaking.

2.

Why are metals malleable?

Layers of metal ions can slide while metallic bonding is maintained.

3.

Why can metals conduct electricity?

Delocalised electrons can move through the structure.

4.

How does the structure of metallic substances allow electrons to move?

Delocalised electrons are free to move.

5.

What type of bonding occurs in metals?

Metallic bonding.

6.

How does metallic bonding explain the physical properties of metals?

Strong metallic bonding explains strength, while mobile delocalised electrons explain electrical conductivity and the ability of layers to slide explains malleability.

1.41 Describe the limitations of particular representations and models, to include dot and cross, ball and stick models and two- and three-dimensional representations

1.

What is one limitation of a dot-and-cross diagram when representing a chemical structure?

It may not show the actual three-dimensional shape or relative sizes accurately.

2.

What is one limitation of a ball-and-stick model when representing a chemical structure?

The balls and sticks do not represent the real sizes and distances between atoms accurately.

3.

Why can a ball-and-stick model give a misleading impression of the space occupied by atoms?

The atoms appear as separate spheres with exaggerated gaps.

4.

What is one limitation of representing a three-dimensional chemical structure in two dimensions?

It cannot fully show the three-dimensional arrangement.

5.

Why are models useful when studying chemical structures despite their limitations?

They simplify complex structures and make them easier to visualise.

6.

Why can no single model provide a completely accurate representation of a chemical structure?

Different models represent different aspects and therefore each has limitations.

1.42 Describe most metals as shiny solids which have high melting points, high density and are good conductors of electricity whereas most non-metals have low boiling points and are poor conductors of electricity

1.

What are the typical physical properties of most metals?

Shiny solids, high melting points, high density and good electrical conductivity.

2.

Why are most metals good conductors of electricity?

They contain delocalised electrons that can move through the metal.

3.

What are the typical physical properties of most non-metals?

Generally lower boiling points and poor electrical conductivity.

4.

How do the melting points of most metals compare with those of most non-metals?

Most metals have higher melting points.

5.

How does the density of most metals compare with that of most non-metals?

Most metals have higher densities.

6.

How does the electrical conductivity of most metals compare with that of most non-metals?

Metals are generally good conductors; non-metals are generally poor conductors.

1.43 Calculate: a relative formula mass given relative atomic masses b percentage by mass of an element in a compound given relative atomic masses

1.

How is the relative formula mass of a compound calculated from its relative atomic masses?

Add the relative atomic masses of all atoms in the formula.

2.

How is the percentage by mass of an element in a compound calculated?

Percentage by mass = (mass of element in one formula unit / Mr of compound) × 100

3.

What is the relative formula mass of H₂O?

Mr(H2O) = 2(1) + 16 = 18

4.

What is the relative formula mass of CaCO₃?

Mr(CaCO3) = 40 + 12 + 3(16) = 100

5.

How would you calculate the percentage by mass of oxygen in H₂O?

Percentage oxygen = (16 / 18) × 100 = 88.9%

6.

How would you calculate the percentage by mass of carbon in CO₂?

Percentage carbon = (12 / 44) × 100 = 27.3%

1.44 Calculate the formulae of simple compounds from reacting masses or percentage composition and understand that these are empirical formulae

1.

What is an empirical formula?

The simplest whole-number ratio of atoms of each element in a compound.

2.

How can reacting masses be used to determine the empirical formula of a compound?

Convert each reacting mass into moles, then find the simplest whole-number ratio.

3.

How can percentage composition be used to determine an empirical formula?

Treat the percentages as masses out of 100 g, convert each to moles, then simplify the ratio.

4.

Why are the masses of elements converted into amounts in moles when determining an empirical formula?

Moles allow the amounts of different elements to be compared directly.

5.

How are mole ratios used to obtain an empirical formula?

Divide all mole amounts by the smallest and multiply if necessary to obtain whole numbers.

6.

What does an empirical formula tell you about the atoms in a compound?

The simplest ratio of the atoms of each element.

1.45 Deduce: a the empirical formula of a compound from the formula of its molecule b the molecular formula of a compound from its empirical formula and its relative molecular mass

1.

How can the empirical formula of a compound be determined from its molecular formula?

Divide all subscripts in the molecular formula by their highest common factor.

2.

How can the molecular formula of a compound be determined from its empirical formula?

Multiplier = Mr of molecular formula / Mr of empirical formula.

3.

What information is needed to determine a molecular formula from an empirical formula?

The empirical formula and relative molecular mass.

4.

How is the multiplier between the empirical formula and molecular formula calculated?

Divide the molecular Mr by the empirical formula Mr.

5.

What is the empirical formula of C₆H₁₂O₆?

CH2O

6.

How can the relative molecular mass be used to determine the molecular formula of a compound?

Calculate the empirical formula Mr, divide the molecular Mr by it, then multiply every empirical-formula subscript by the resulting whole number.

1.46 Describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide

1.

How can an experiment be used to determine the empirical formula of magnesium oxide?

Weigh magnesium, heat it strongly in oxygen, allow it to cool, then reweigh the magnesium oxide and calculate the masses of magnesium and oxygen before converting them to moles and finding the simplest ratio.

2.

Why is magnesium weighed before it is heated in an experiment to determine the formula of magnesium oxide?

To determine the starting mass of magnesium.

3.

Why is magnesium heated in oxygen when determining the empirical formula of magnesium oxide?

To react the magnesium with oxygen and form magnesium oxide.

4.

What measurements are needed to calculate the empirical formula of magnesium oxide?

Mass of magnesium before heating and mass of magnesium oxide after heating.

5.

Why must magnesium oxide be heated to constant mass?

To ensure the reaction is complete and all possible oxygen has reacted.

6.

How can the measured masses of magnesium and oxygen be converted into a ratio of moles?

Calculate moles using n = mass / Ar, then divide by the smallest number of moles.

1.47 Explain the law of conservation of mass applied to: a a closed system including a precipitation reaction in a closed flask b a non-enclosed system including a reaction in an open flask that takes in or gives out a gas

1.

What does the law of conservation of mass state?

Mass is neither created nor destroyed in a chemical reaction.

2.

Why is mass conserved during a chemical reaction in a closed system?

No reactants or products can escape or enter the system.

3.

How does a precipitation reaction in a closed flask demonstrate conservation of mass?

The total mass before and after the precipitation reaction remains the same.

4.

Why can the measured mass of a non-enclosed system change during a chemical reaction?

Gas can enter or leave the system.

5.

Why can a reaction in an open flask appear to lose mass when a gas is produced?

The gas escapes from the flask, so the measured mass decreases.

6.

Why can a reaction in an open flask appear to gain mass when a gas from the surroundings is taken in?

Gas from the surroundings enters the flask, so the measured mass increases.

1.48 Calculate masses of reactants and products from balanced equations, given the mass of one substance

1.

How can the mass of an unknown reactant or product be calculated from a balanced chemical equation?

Convert the known mass to moles, use the mole ratio from the balanced equation, then convert the required moles back to mass.

2.

Why must a chemical equation be balanced before using it to calculate reacting masses?

The coefficients give the correct mole ratios between reactants and products.

3.

How are mole ratios obtained from a balanced chemical equation?

From the coefficients in the balanced equation.

4.

What information is needed to calculate the mass of an unknown substance from a balanced equation?

The mass of the known substance, its Ar/Mr, the balanced equation and the Ar/Mr of the unknown substance.

5.

How are moles converted into mass when calculating reacting masses?

Mass = moles × Mr, or mass = moles × Ar for atoms.

6.

How can the balanced equation 2Mg + O₂ → 2MgO be used to calculate the mass of magnesium oxide produced from a known mass of magnesium?

Use the 2:2 mole ratio for Mg:MgO, calculate moles of Mg from the known mass, then calculate the same number of moles of MgO and convert to mass.

1.49 Calculate the concentration of solutions in g dm⁻³

1.

What is meant by the concentration of a solution in g dm⁻³?

The mass of solute dissolved per cubic decimetre of solution.

2.

How is concentration in g dm⁻³ calculated?

Concentration = mass / volume

3.

What units are used for mass and volume when calculating concentration in g dm⁻³?

Mass in g and volume in dm³.

4.

How can the concentration of a solution be calculated from its mass of solute and volume?

concentration (g dm⁻³) = mass of solute (g) / volume of solution (dm³)

5.

How can the mass of solute be calculated when the concentration and volume are known?

mass = concentration × volume

6.

How can the volume of solution be calculated when the concentration and mass of solute are known?

volume = mass / concentration

1.50 Recall that one mole of particles of a substance is defined as: a the Avogadro constant number of particles (6.02 × 10²³ atoms, molecules, formulae or ions) of that substance b a mass of 'relative particle mass'

1.

What is a mole?

The amount of substance containing 6.02 × 10²³ particles.

2.

How many particles are present in one mole of a substance?

6.02 × 10²³

3.

What is the Avogadro constant?

6.02 × 10²³ mol⁻¹

4.

Which types of particles can be counted using the Avogadro constant?

Atoms, molecules, formula units or ions.

5.

What does the mass of one mole of a substance correspond to?

Its relative particle mass in grams.

6.

How is the relative particle mass related to the mass of one mole of a substance?

The mass of one mole in grams is numerically equal to the relative particle mass.

1.51 Calculate the number of: a moles of particles of a substance in a given mass of that substance and vice versa b particles of a substance in a given number of moles of that substance and vice versa c particles of a substance in a given mass of that substance and vice versa

1.

How can the number of moles in a given mass of a substance be calculated?

moles = mass / Mr

2.

How can the mass of a substance be calculated when the number of moles is known?

mass = moles × Mr

3.

How can the number of particles in a given number of moles be calculated using the Avogadro constant?

particles = moles × 6.02 × 10²³

4.

How can the number of moles be calculated when the number of particles is known?

moles = particles / (6.02 × 10²³)

5.

How can the number of particles in a given mass of a substance be calculated?

Calculate moles from mass, then multiply by 6.02 × 10²³.

6.

How can the mass of a substance be calculated when the number of particles is known?

Calculate moles from particles, then multiply by Mr.

1.52 Explain why, in a reaction, the mass of product formed is controlled by the mass of the reactant which is not in excess

1.

What is meant by the limiting reactant in a chemical reaction?

The reactant that is completely used up first.

2.

Why does the limiting reactant control the maximum amount of product that can form?

Once it is used up, no further product can form.

3.

What happens to the limiting reactant when a reaction goes to completion?

It is completely used up.

4.

What happens to a reactant that is present in excess when the limiting reactant is completely used up?

It remains unreacted.

5.

How can the limiting reactant be identified from the amounts of reactants used in a reaction?

Calculate the moles of each reactant and compare them with the stoichiometric ratio in the balanced equation.

6.

Why does increasing the amount of an excess reactant not increase the amount of product once the limiting reactant is used up?

The limiting reactant has already been completely consumed, so extra excess reactant cannot react.

1.53 Deduce the stoichiometry of a reaction from the masses of the reactants and products

1.

Aluminium reacts with oxygen according to the equation 4Al + 3O₂ → 2Al₂O₃. What mass of aluminium oxide can be formed when 135 g of aluminium is completely reacted with oxygen?

135 g Al = 5.0 mol. 4Al → 2Al2O3, so Al2O3 = 2.5 mol. Mr(Al2O3) = 102. Mass = 2.5 × 102 = 255 g.

2.

Calcium reacts with oxygen according to the equation 2Ca + O₂ → 2CaO. What mass of calcium oxide is formed when 20 g of calcium reacts completely with oxygen?

20 g Ca = 20 / 40 = 0.50 mol. CaO = 0.50 mol. Mr(CaO) = 56. Mass = 0.50 × 56 = 28 g.

3.

Magnesium reacts with oxygen according to the equation 2Mg + O₂ → 2MgO. What mass of magnesium oxide is formed when 12 g of magnesium reacts completely with oxygen?

12 g Mg = 12 / 24 = 0.50 mol. MgO = 0.50 mol. Mr(MgO) = 40. Mass = 0.50 × 40 = 20 g.

4.

Carbon reacts with oxygen according to the equation C + O₂ → CO₂. What mass of carbon dioxide is formed when 6 g of carbon reacts completely with oxygen?

6 g C = 6 / 12 = 0.50 mol. CO2 = 0.50 mol. Mr(CO2) = 44. Mass = 0.50 × 44 = 22 g.

5.

Iron reacts with sulfur according to the equation Fe + S → FeS. What mass of iron sulfide is formed when 28 g of iron reacts completely with sulfur?

28 g Fe = 28 / 56 = 0.50 mol. FeS = 0.50 mol. Mr(FeS) = 88. Mass = 0.50 × 88 = 44 g.

6.

A reaction produces 44 g of carbon dioxide from 12 g of carbon. Use these masses to determine the simplest mole ratio of carbon to carbon dioxide in the reaction.

12 g C = 1 mol and 44 g CO2 = 1 mol. Simplest ratio = 1:1.

2.1 Describe the arrangement, movement and the relative energy of particles in each of the three states of matter: solid, liquid and gas

1.

How are particles arranged in a solid?

Closely packed in a fixed, regular arrangement.

2.

How do particles move in a solid?

Vibrate about fixed positions.

3.

How are particles arranged in a liquid?

Close together but irregularly arranged.

4.

How do particles move in a liquid?

Move around and past one another.

5.

How are particles arranged in a gas?

Far apart and randomly arranged.

6.

How does the energy of particles compare between solids, liquids and gases?

Particle energy generally increases from solid to liquid to gas.

2.2 Recall the names used for the interconversions between the three states of matter, recognising that these are physical changes: contrasted with chemical reactions that result in chemical changes

1.

What is the name of the change from a solid to a liquid?

Melting.

2.

What is the name of the change from a liquid to a gas?

Boiling or evaporation.

3.

What is the name of the change from a gas to a liquid?

Condensation.

4.

What is the name of the change from a liquid to a solid?

Freezing.

5.

What is the name of the change from a solid directly to a gas?

Sublimation.

6.

What is the name of the change from a gas directly to a solid?

Deposition.

2.3 Explain the changes in arrangement, movement and energy of particles during these interconversions

1.

What happens to the arrangement and movement of particles when a solid melts?

Particles become less ordered, move more freely and gain energy.

2.

What happens to the energy of particles when a liquid evaporates?

The particles gain energy.

3.

What happens to the arrangement and movement of particles when a gas condenses?

Particles move closer together, become more ordered and lose energy.

4.

What happens to the energy of particles when a liquid freezes?

The particles lose energy.

5.

What happens to the particles when a solid sublimes?

Particles gain energy and move directly from a solid arrangement to widely separated gas particles.

6.

Why do particles become more widely separated when a substance changes from a liquid to a gas?

Gas particles have greater energy and move more freely, overcoming the forces keeping them close together.

2.4 Predict the physical state of a substance under specified conditions, given suitable data

1.

How can melting point data be used to determine whether a substance is solid or liquid at a given temperature?

Below the melting point it is solid; above the melting point it is liquid unless the boiling point has been exceeded.

2.

How can boiling point data be used to determine whether a substance is liquid or gas at a given temperature?

Below the boiling point it is liquid; above the boiling point it is gas, provided it is above the melting point.

3.

A substance has a melting point of 20°C and a boiling point of 80°C. What state is it in at 10°C?

Solid.

4.

A substance has a melting point of −10°C and a boiling point of 60°C. What state is it in at 25°C?

Liquid.

5.

A substance has a melting point of 50°C and a boiling point of 120°C. What state is it in at 150°C?

Gas.

6.

A substance has a melting point of −5°C and a boiling point of 95°C. What state is it in at −20°C?

Solid.

2.5 Explain the difference between the use of 'pure' in chemistry compared with its everyday use and the differences in chemistry between a pure substance and a mixture

1.

What does the term "pure" mean in chemistry?

A substance containing only one element or one compound.

2.

How can the everyday meaning of "pure" differ from its meaning in chemistry?

Everyday use can mean clean, uncontaminated or not mixed with unwanted material; in chemistry it has a precise meaning.

3.

What is the difference between a pure substance and a mixture?

A pure substance contains only one substance; a mixture contains two or more substances.

4.

Why does a pure substance contain only one substance?

Because only one substance is present.

5.

Why can a mixture contain more than one element or compound?

Different elements and compounds can be physically mixed without reacting.

6.

Why can the components of a mixture retain their individual chemical properties?

They have not necessarily chemically reacted and can retain their own properties.

2.6 Interpret melting point data to distinguish between pure substances which have a sharp melting point and mixtures which melt over a range of temperatures

1.

How does the melting point of a pure substance differ from that of a mixture?

A pure substance has a sharp melting point; a mixture usually melts over a range.

2.

What does a sharp melting point indicate about the purity of a substance?

It suggests the substance is pure.

3.

What does melting over a range of temperatures indicate about a substance?

It suggests the substance is a mixture or impure.

4.

A sample melts sharply at 78°C. What does this suggest about the sample?

It is likely to be a pure substance.

5.

A sample begins melting at 70°C and finishes melting at 76°C. What does this suggest about the sample?

It is likely to be a mixture or impure.

6.

How can melting point data be used to distinguish between a pure substance and a mixture?

Compare whether melting occurs sharply at one temperature or over a range.

2.7 Explain the types of mixtures that can be separated by using the following experimental techniques: a simple distillation b fractional distillation c filtration d crystallisation e paper chromatography

1.

What type of mixture can be separated using simple distillation?

A solution containing a soluble solid and a solvent.

2.

What type of mixture can be separated using fractional distillation?

A mixture of miscible liquids with different boiling points.

3.

What type of mixture can be separated using filtration?

An insoluble solid mixed with a liquid.

4.

What type of mixture can be separated using crystallisation?

A solution containing a dissolved solid.

5.

What type of mixture can be separated using paper chromatography?

A mixture of soluble substances such as dyes or inks.

6.

How does the type of mixture determine which separation technique should be used?

Choose the method according to properties such as solubility, boiling point and physical state.

2.8 Describe an appropriate experimental technique to separate a mixture, knowing the properties of the components of the mixture

1.

How would you choose an appropriate technique for separating a mixture?

Identify the physical properties that differ, then choose the appropriate separation method.

2.

How would you separate an insoluble solid from a liquid?

Filtration.

3.

How would you separate a soluble solid from a solution to obtain the solid?

Crystallisation.

4.

How would you separate two miscible liquids with different boiling points?

Fractional distillation.

5.

How would you obtain a pure solvent from a solution containing a dissolved solid?

Simple distillation.

6.

Why must the physical properties of the components be considered when choosing a separation technique?

Separation techniques rely on differences in physical properties.

2.9 Describe paper chromatography as the separation of mixtures of soluble substances by running a solvent (mobile phase) through the mixture on the paper (the paper contains the stationary phase), which causes the substances to move at different rates over the paper

1.

What is the purpose of the solvent in paper chromatography?

To dissolve and carry the soluble substances up the paper.

2.

What is meant by the mobile phase in paper chromatography?

The solvent that moves through the paper.

3.

What is meant by the stationary phase in paper chromatography?

The paper.

4.

Why do different substances move different distances during paper chromatography?

They have different attractions for the mobile and stationary phases.

5.

Why must the substances being separated be soluble in the solvent?

They must dissolve in the mobile phase to move.

6.

How does paper chromatography separate the components of a mixture?

The solvent carries the substances at different rates, separating them into spots.

2.10 Interpret a paper chromatogram: a to distinguish between pure and impure substances b to identify substances by comparison with known substances c to identify substances by calculation and use of Rf values

1.

How can a paper chromatogram be used to determine whether a substance is pure or impure?

One spot indicates a pure substance; more than one spot indicates an impure substance or mixture.

2.

How can a paper chromatogram be used to identify an unknown substance by comparison with known substances?

Compare the positions of spots with those from known substances.

3.

What does an Rf value represent in paper chromatography?

Rf = distance travelled by substance / distance travelled by solvent front

4.

How is an Rf value calculated from a paper chromatogram?

Rf = distance travelled by substance / distance travelled by solvent front

5.

A substance travels 4 cm and the solvent front travels 8 cm. What is the Rf value?

Rf = 4 / 8 = 0.50

6.

Why can substances with the same Rf value under identical conditions be identified as the same substance?

Under the same conditions, the same substance has the same Rf value.

2.11 Core Practical: Investigate the composition of inks using simple distillation and paper chromatography

1.

How can simple distillation be used to investigate the composition of an ink?

Heat the ink and collect the solvent that distils, leaving the dissolved substances behind.

2.

How can paper chromatography be used to investigate the composition of an ink?

Place ink on chromatography paper, place the paper in a suitable solvent and compare the separated spots.

3.

What information does the chromatogram provide about the dyes present in an ink?

The number and positions of the dye spots present.

4.

How can known dye samples be used to identify dyes in an unknown ink?

Compare the spots or Rf values with known dyes.

5.

Why is a pencil used to mark the starting line in paper chromatography rather than ink?

Ink would dissolve and move with the solvent, whereas pencil does not.

6.

What variables should be controlled when comparing the chromatograms of different inks?

Solvent, paper type, spot size, starting-line position, solvent depth and temperature.

2.12 Describe how: a waste and ground water can be made potable, including the need for sedimentation, filtration and chlorination b sea water can be made potable by using distillation c water used in analysis must not contain any dissolved salts

1.

What processes are used to make waste water or ground water potable?

Sedimentation, filtration and chlorination.

2.

What is the purpose of sedimentation when producing potable water?

To allow larger suspended particles to settle.

3.

What is the purpose of filtration when producing potable water?

To remove smaller insoluble particles.

4.

Why is chlorine added during the production of potable water?

To kill microorganisms.

5.

How can sea water be made potable using distillation?

Distil the seawater so that water evaporates and condenses separately from dissolved salts.

6.

Why must water used in chemical analysis not contain dissolved salts?

Dissolved salts could interfere with chemical tests and analysis.

3.1 Recall that acids in solution are sources of hydrogen ions and alkalis in solution are sources of hydroxide ions

1.

Which ions are produced when an acid dissolves in water?

H+

2.

Which ions are produced when an alkali dissolves in water?

OH−

3.

What is the chemical formula of a hydrogen ion?

H+

4.

What is the chemical formula of a hydroxide ion?

OH−

5.

Which ions are responsible for the acidic properties of aqueous acids?

H+

6.

Which ions are responsible for the alkaline properties of aqueous alkalis?

OH−

3.2 Recall that a neutral solution has a pH of 7 and that acidic solutions have lower pH values and alkaline solutions higher pH values

1.

What pH value does a neutral solution have?

7

2.

What pH values indicate that a solution is acidic?

Below 7

3.

What pH values indicate that a solution is alkaline?

Above 7

4.

Which is more acidic: a solution with pH 2 or a solution with pH 5?

pH 2

5.

Which is more alkaline: a solution with pH 9 or a solution with pH 12?

pH 12

6.

What does a pH of 7 indicate about a solution?

It is neutral.

3.3 Recall the effect of acids and alkalis on indicators, including litmus, methyl orange and phenolphthalein

1.

What colour does blue litmus turn when placed in an acidic solution?

Red.

2.

What colour does red litmus turn when placed in an alkaline solution?

Blue.

3.

What colour is methyl orange in an acidic solution?

Red.

4.

What colour is methyl orange in an alkaline solution?

Yellow.

5.

What colour is phenolphthalein in an acidic solution?

Colourless.

6.

What colour is phenolphthalein in an alkaline solution?

Pink.

3.4 Recall that the higher the concentration of hydrogen ions in an acidic solution, the lower the pH; and the higher the concentration of hydroxide ions in an alkaline solution, the higher the pH

1.

How does increasing the concentration of hydrogen ions affect the pH of an acidic solution?

pH decreases.

2.

How does decreasing the concentration of hydrogen ions affect the pH of an acidic solution?

pH increases.

3.

How does increasing the concentration of hydroxide ions affect the pH of an alkaline solution?

pH increases.

4.

How does decreasing the concentration of hydroxide ions affect the pH of an alkaline solution?

pH decreases.

5.

Which has the higher hydrogen ion concentration: an acidic solution with pH 2 or an acidic solution with pH 4?

pH 2

6.

Which has the higher hydroxide ion concentration: an alkaline solution with pH 10 or an alkaline solution with pH 12?

pH 12

3.5 Recall that as hydrogen ion concentration in a solution increases by a factor of 10, the pH of the solution decreases by 1

1.

By how much does the pH decrease when the hydrogen ion concentration increases by a factor of 10?

By 1.

2.

By what factor does the hydrogen ion concentration increase when the pH decreases by 1?

By a factor of 10.

3.

By what factor does the hydrogen ion concentration increase when the pH decreases by 2?

By a factor of 100.

4.

An acidic solution changes from pH 5 to pH 4. By what factor has its hydrogen ion concentration increased?

A factor of 10.

5.

An acidic solution changes from pH 6 to pH 3. By what factor has its hydrogen ion concentration increased?

A factor of 1000.

6.

An acidic solution has a pH of 4. What will its pH become if its hydrogen ion concentration increases by a factor of 100?

pH 2.

3.6 Core Practical: Investigate the change in pH on adding powdered calcium hydroxide or calcium oxide to a fixed volume of dilute hydrochloric acid

1.

How could you investigate the change in pH when powdered calcium hydroxide is added to a fixed volume of dilute hydrochloric acid?

Measure a fixed volume of dilute hydrochloric acid into a container, measure the pH, add known amounts of powdered calcium hydroxide, mix and measure the pH after each addition.

2.

What should be measured after each addition of powdered calcium hydroxide to dilute hydrochloric acid when investigating changes in pH?

The pH of the mixture.

3.

Why should the volume of dilute hydrochloric acid be kept constant when investigating the effect of adding powdered calcium hydroxide on pH?

To ensure changes in pH are due to the calcium hydroxide rather than changes in acid volume.

4.

Why should known masses or measured amounts of powdered calcium hydroxide be added to the dilute hydrochloric acid during the investigation?

To know how much calcium hydroxide has been added and allow the pH change to be related to the amount added.

5.

What should be done to the hydrochloric acid mixture before measuring its pH after each addition of calcium hydroxide?

Stir the mixture thoroughly before measuring the pH.

6.

How would the pH of dilute hydrochloric acid change as increasing amounts of powdered calcium hydroxide are added?

The pH increases as more calcium hydroxide is added.

3.7 Explain the terms dilute and concentrated, with respect to amount of substances in solution

1.

What is meant by a dilute solution in terms of the amount of solute dissolved?

A solution containing a small amount of solute per unit volume.

2.

What is meant by a concentrated solution in terms of the amount of solute dissolved?

A solution containing a large amount of solute per unit volume.

3.

Which contains more solute per unit volume: a dilute solution or a concentrated solution?

A concentrated solution.

4.

How can a concentrated solution be made more dilute?

Add solvent.

5.

How can a dilute solution be made more concentrated?

Remove solvent or add more solute.

6.

What does the term concentration describe about the amount of solute in a solution?

The amount of solute present per unit volume.

3.8 Explain the terms weak and strong acids, with respect to the degree of dissociation into ions

1.

What is meant by a strong acid in terms of its dissociation into ions?

It dissociates almost completely into ions in aqueous solution.

2.

What is meant by a weak acid in terms of its dissociation into ions?

It only partially dissociates into ions.

3.

What happens to the particles of a strong acid when it dissolves in water?

Most acid molecules dissociate into ions.

4.

What happens to the particles of a weak acid when it dissolves in water?

Only some acid molecules dissociate into ions.

5.

What is meant by the dissociation of an acid into ions?

The splitting of acid molecules into ions in solution.

6.

Why does a strong acid have a greater degree of dissociation into ions than a weak acid?

A larger proportion of its molecules dissociate into ions.

3.9 Recall that a base is any substance that reacts with an acid to form a salt and water only

1.

What is a base according to the Edexcel GCSE definition?

A substance that reacts with an acid to form only a salt and water.

2.

What products are formed when a base reacts with an acid?

A salt and water.

3.

What type of reaction occurs when an acid reacts with a base to form only a salt and water?

Neutralisation.

4.

Is copper oxide classified as a base because it reacts with an acid to form a salt and water?

Yes.

5.

Is sodium hydroxide classified as a base because it reacts with an acid to form a salt and water?

Yes.

6.

What products must be formed when a substance reacts with an acid for that substance to meet the Edexcel definition of a base?

A salt and water only.

3.10 Recall that alkalis are soluble bases

1.

What is an alkali?

A soluble base.

2.

What property of a base makes it an alkali?

It is soluble in water.

3.

Is sodium hydroxide an alkali?

Yes.

4.

Is potassium hydroxide an alkali?

Yes.

5.

Is copper oxide an alkali?

No.

6.

Why is copper oxide classified as a base but not as an alkali?

Copper oxide is insoluble in water.

3.11 Explain the general reactions of aqueous solutions of acids with: a metals b metal oxides c metal hydroxides d metal carbonates to produce salts

1.

What are the general products when an aqueous acid reacts with a metal?

Salt + hydrogen.

2.

What are the general products when an aqueous acid reacts with a metal oxide?

Salt + water.

3.

What are the general products when an aqueous acid reacts with a metal hydroxide?

Salt + water.

4.

What are the general products when an aqueous acid reacts with a metal carbonate?

Salt + water + carbon dioxide.

5.

Which gas is generally produced when an aqueous acid reacts with a metal?

Hydrogen.

6.

Which gas is generally produced when an aqueous acid reacts with a metal carbonate?

Carbon dioxide.

3.12 Describe the chemical test for: a hydrogen b carbon dioxide (using limewater)

1.

How can hydrogen gas be chemically tested for?

Use a lit splint; hydrogen produces a squeaky pop.

2.

What positive result is observed when a lit splint is placed in a sample of hydrogen gas?

It burns with a squeaky pop.

3.

How can carbon dioxide gas be chemically tested for using limewater?

A squeaky pop.

4.

What positive result is observed when carbon dioxide is bubbled through limewater?

Hydrogen burns rapidly in the presence of oxygen.

5.

What gas produces a squeaky pop when a lit splint is applied to it?

A lit splint.

6.

What gas turns limewater cloudy when it is bubbled through limewater?

Carbon dioxide.

3.13 Describe a neutralisation reaction as a reaction between an acid and a base

1.

What is a neutralisation reaction?

A reaction between an acid and a base.

2.

What two types of substances react in a neutralisation reaction?

An acid and a base.

3.

What products are formed when an acid reacts with a base in a neutralisation reaction?

Salt and water.

4.

Is the reaction between hydrochloric acid and sodium hydroxide a neutralisation reaction?

Yes.

5.

Is the reaction between nitric acid and copper oxide a neutralisation reaction?

Yes.

6.

Why is the reaction between an acid and a base described as neutralisation?

The acid is neutralised by the base.

3.14 Explain an acid-alkali neutralisation as a reaction in which hydrogen ions (H+) from the acid react with hydroxide ions (OH–) from the alkali to form water

1.

Which ions from an acid react with hydroxide ions during acid-alkali neutralisation?

H+

2.

Which ions from an alkali react with hydrogen ions during acid-alkali neutralisation?

OH−

3.

What substance is formed when H+ ions react with OH− ions?

Water, H2O.

4.

What is the ionic equation for acid-alkali neutralisation?

H+ + OH− → H2O

5.

Why does the reaction between H+ ions and OH− ions result in neutralisation?

The acidic H+ ions and alkaline OH− ions are removed by forming water.

6.

What happens to the H+ ions and OH− ions during acid-alkali neutralisation?

They react together to form H2O.

3.15 Explain why, if soluble salts are prepared from an acid and an insoluble reactant: a excess of the reactant is added b the excess reactant is removed c the solution remaining is only salt and water

1.

Why is an excess of an insoluble reactant added when preparing a soluble salt from an acid?

To ensure all the acid reacts.

2.

Why must excess insoluble reactant be removed when preparing a soluble salt from an acid?

To prevent unreacted reactant contaminating the salt solution.

3.

How can excess insoluble reactant be removed from a salt solution?

By filtration.

4.

Why is filtration suitable for removing excess insoluble reactant from a salt solution?

The excess reactant is insoluble and remains as a solid.

5.

What substances should remain in the solution when an acid has completely reacted with an insoluble reactant?

Salt and water.

6.

Why does adding an excess of an insoluble reactant help ensure that all of the acid has reacted?

Any acid remaining can react with the excess solid until the acid is completely used up.

3.16 Explain why, if soluble salts are prepared from an acid and a soluble reactant: a titration must be used b the acid and the soluble reactant are then mixed in the correct proportions c the solution remaining, after reaction, is only salt and water

1.

Why must titration be used when preparing a soluble salt from an acid and a soluble reactant?

Both reactants are soluble, so excess cannot be removed by filtration.

2.

Why cannot an excess soluble reactant be removed by filtration when preparing a soluble salt?

It would remain dissolved in the solution.

3.

What is determined during a titration when preparing a soluble salt from an acid and a soluble reactant?

The exact volume needed for complete neutralisation.

4.

Why must the acid and soluble reactant be mixed in the correct proportions when preparing a pure soluble salt?

To ensure neither reactant is left in excess.

5.

What substances should remain in the solution after an acid and soluble reactant have reacted in the correct proportions?

Salt and water.

6.

Why does using the correct proportions of an acid and soluble reactant prevent unreacted reactants remaining in the final solution?

The reactants react completely with each other, so no excess reactant remains.

3.17 Core Practical: Investigate the preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath

1.

Which acid reacts with copper oxide to prepare hydrated copper sulfate crystals?

Sulfuric acid.

2.

Why is copper oxide added in excess when preparing hydrated copper sulfate crystals from copper oxide and sulfuric acid?

To ensure all sulfuric acid reacts.

3.

How is excess copper oxide removed when preparing hydrated copper sulfate crystals?

Filter the mixture.

4.

Why is a water bath used when preparing hydrated copper sulfate crystals?

To gently evaporate water without overheating the solution.

5.

How is water removed from the copper sulfate solution to allow hydrated copper sulfate crystals to form?

Gently heat the solution to evaporate some water, then allow it to cool so crystals form; filter and dry the crystals.

6.

Why should the hydrated copper sulfate crystals not be heated strongly until completely dry?

Strong heating can remove water of crystallisation.

3.18 Describe how to carry out an acid-alkali titration, using burette, pipette and a suitable indicator, to prepare a pure, dry salt

1.

What piece of apparatus is used to deliver a measured variable volume of solution during an acid-alkali titration?

Burette.

2.

What piece of apparatus is used to transfer an accurately measured fixed volume of solution during an acid-alkali titration?

It allows a variable volume to be delivered accurately.

3.

Why is a suitable indicator used during an acid-alkali titration?

To identify the end point of the reaction.

4.

What volume should be recorded from the burette during an acid-alkali titration?

Final reading − initial reading.

5.

Why is an acid-alkali titration repeated until concordant titres are obtained?

To obtain a reliable mean titre from closely agreeing results.

6.

How can an acid-alkali titration be used to prepare a pure, dry soluble salt?

Use the titration to find exact reacting volumes, repeat without indicator using the same volumes, then evaporate water and crystallise the salt.

3.19 Recall the general rules which describe the solubility of common types of substances in water: a all common sodium, potassium and ammonium salts are soluble b all nitrates are soluble c common chlorides are soluble except those of silver and lead d common sulfates are soluble except those of lead, barium and calcium e common carbonates and hydroxides are insoluble except those of sodium, potassium and ammonium

1.

Which common salts of sodium, potassium and ammonium are soluble in water?

All common sodium, potassium and ammonium salts.

2.

Are all nitrates soluble in water?

Yes.

3.

Which common chlorides are insoluble in water?

Silver chloride and lead chloride.

4.

Which common sulfates are insoluble in water?

Lead sulfate, barium sulfate and calcium sulfate.

5.

Which common carbonates are insoluble in water, apart from the stated exceptions?

Common carbonates are insoluble except sodium, potassium and ammonium carbonates.

6.

Which common hydroxides are insoluble in water, apart from the stated exceptions?

Common hydroxides are insoluble except sodium, potassium and ammonium hydroxides.

3.20 Predict, using solubility rules, whether or not a precipitate will be formed when named solutions are mixed together, naming the precipitate if any

1.

What is a precipitate in a chemical reaction between two aqueous solutions?

An insoluble solid formed when two aqueous solutions are mixed.

2.

Will a precipitate form when aqueous silver nitrate is mixed with aqueous sodium chloride?

Yes.

3.

What is the name of the precipitate formed when aqueous silver nitrate is mixed with aqueous sodium chloride?

Silver chloride, AgCl.

4.

Will a precipitate form when aqueous barium nitrate is mixed with aqueous sodium sulfate?

Yes.

5.

What is the name of the precipitate formed when aqueous barium nitrate is mixed with aqueous sodium sulfate?

Barium sulfate, BaSO4.

6.

How can solubility rules be used to predict whether mixing two named aqueous solutions will produce a precipitate?

Swap the ions, apply the solubility rules and identify any insoluble product.

3.21 Describe the method used to prepare a pure, dry sample of an insoluble salt

1.

How can an insoluble salt be prepared from two soluble solutions?

Mix two suitable soluble solutions containing the required ions.

2.

Why are two soluble solutions mixed when preparing an insoluble salt?

To provide the ions needed to form the insoluble salt.

3.

What happens when the ions of an insoluble salt combine in aqueous solution?

They form an insoluble solid precipitate.

4.

How is an insoluble salt separated from the aqueous solution in which it forms?

Filter the mixture.

5.

Why is an insoluble salt washed with distilled water after filtration?

To remove soluble impurities.

6.

How is a pure, dry sample of an insoluble salt obtained after filtration and washing?

Filter, wash the precipitate with distilled water and dry it.

3.22 Recall that electrolytes are ionic compounds in the molten state or dissolved in water

1.

What is an electrolyte?

An ionic compound that conducts electricity when molten or dissolved in water.

2.

What type of substances can act as electrolytes when molten or dissolved in water?

Ionic compounds.

3.

Why can an ionic compound conduct electricity when it is molten?

The ions are free to move.

4.

Why can an ionic compound conduct electricity when dissolved in water?

The ions are free to move through the solution.

5.

Can a solid ionic compound act as an electrolyte?

No.

6.

What must happen to an ionic compound for its ions to be able to move and conduct electricity?

Its ions must become free to move.

3.23 Describe electrolysis as a process in which electrical energy, from a direct current supply, decomposes electrolytes

1.

What is electrolysis?

The decomposition of an electrolyte using electrical energy.

2.

What type of electrical supply is used during electrolysis?

Direct current, d.c.

3.

What happens to an electrolyte during electrolysis?

It is decomposed into simpler substances.

4.

What form of energy is supplied to an electrolyte during electrolysis?

Electrical energy.

5.

What does the term decomposes mean in the context of electrolysis?

Broken down into simpler substances.

6.

Why is electrolysis considered a chemical change rather than a physical change?

New substances are formed.

3.24 Explain the movement of ions during electrolysis, in which: a positively charged cations migrate to the negatively charged cathode b negatively charged anions migrate to the positively charged anode

1.

Which electrode do positively charged cations move towards during electrolysis?

Cathode.

2.

Which electrode do negatively charged anions move towards during electrolysis?

Anode.

3.

Why do positively charged cations move towards the cathode during electrolysis?

Opposite charges attract.

4.

Why do negatively charged anions move towards the anode during electrolysis?

Opposite charges attract.

5.

What charge does the cathode have during electrolysis?

Negative.

6.

What charge does the anode have during electrolysis?

Positive.

3.25 Explain the formation of the products in the electrolysis, using inert electrodes, of some electrolytes, including: a copper chloride solution b sodium chloride solution c sodium sulfate solution d water acidified with sulfuric acid e molten lead bromide (demonstration)

1.

What products are formed at the electrodes when aqueous copper chloride is electrolysed using inert electrodes?

Cathode: copper. Anode: chlorine.

2.

What products are formed at the electrodes when aqueous sodium chloride is electrolysed using inert electrodes?

Cathode: hydrogen. Anode: chlorine.

3.

What products are formed at the electrodes when aqueous sodium sulfate is electrolysed using inert electrodes?

Cathode: hydrogen. Anode: oxygen.

4.

What products are formed at the electrodes when water acidified with sulfuric acid is electrolysed using inert electrodes?

Cathode: hydrogen. Anode: oxygen.

5.

What products are formed at the electrodes when molten lead bromide is electrolysed?

Cathode: lead. Anode: bromine.

6.

Why are inert electrodes used when electrolysing aqueous copper chloride, sodium chloride, sodium sulfate, or acidified water?

They do not react with the electrolyte and do not interfere with the products.

3.26 Predict the products of electrolysis of other binary, ionic compounds in the molten state

1.

What products are formed when a binary ionic compound is electrolysed in the molten state?

The metal forms at the cathode and the non-metal forms at the anode.

2.

What product forms from the cation at the cathode when a molten binary ionic compound is electrolysed?

The metal.

3.

What product forms from the anion at the anode when a molten binary ionic compound is electrolysed?

The non-metal.

4.

What products would form when molten sodium chloride is electrolysed?

Sodium at the cathode and chlorine at the anode.

5.

What products would form when molten magnesium bromide is electrolysed?

Magnesium at the cathode and bromine at the anode.

6.

How can the ions present in a molten binary ionic compound be used to predict the products of electrolysis?

Identify the cation and anion, then the cation forms at the cathode and the anion forms at the anode.

3.27 Write half equations for reactions occurring at the anode and cathode in electrolysis

1.

What does a half equation show during electrolysis?

It shows the electron transfer occurring at one electrode.

2.

What type of electron change is shown in a cathode half equation during electrolysis?

Gain of electrons.

3.

What type of electron change is shown in an anode half equation during electrolysis?

Loss of electrons.

4.

What is the cathode half equation for the formation of copper from Cu²⁺ ions?

Cu²+ + 2e− → Cu

5.

What is the anode half equation for the formation of chlorine from Cl⁻ ions?

2Cl− → Cl2 + 2e−

6.

What is the cathode half equation for the formation of lead from Pb²⁺ ions?

Pb²+ + 2e− → Pb

3.28 Explain oxidation and reduction in terms of loss or gain of electrons

1.

What is oxidation in terms of electrons?

Loss of electrons.

2.

What is reduction in terms of electrons?

Gain of electrons.

3.

What happens to the number of electrons when a substance is oxidised?

It loses electrons.

4.

What happens to the number of electrons when a substance is reduced?

It gains electrons.

5.

Which process involves the loss of electrons: oxidation or reduction?

Oxidation.

6.

Which process involves the gain of electrons: oxidation or reduction?

Reduction.

3.29 Recall that reduction occurs at the cathode and that oxidation occurs at the anode in electrolysis reactions

1.

At which electrode does reduction occur during electrolysis?

Cathode.

2.

At which electrode does oxidation occur during electrolysis?

Anode.

3.

What electron process occurs at the cathode during electrolysis?

Gain of electrons.

4.

What electron process occurs at the anode during electrolysis?

Loss of electrons.

5.

Why is the cathode associated with reduction during electrolysis?

The cations gain electrons there.

6.

Why is the anode associated with oxidation during electrolysis?

The anions lose electrons there.

3.30 Explain the formation of the products in the electrolysis of copper sulfate solution, using copper electrodes, and how this electrolysis can be used to purify copper

1.

What happens at the cathode when copper sulfate solution is electrolysed using copper electrodes?

Cu²+ ions gain electrons and form copper.

2.

What happens at the anode when copper sulfate solution is electrolysed using copper electrodes?

Copper atoms lose electrons and enter the solution as Cu²+ ions.

3.

What happens to the mass of the cathode during the electrolysis of copper sulfate solution using copper electrodes?

It increases.

4.

What happens to the mass of the anode during the electrolysis of copper sulfate solution using copper electrodes?

It decreases.

5.

Why does the concentration of copper ions in copper sulfate solution remain approximately constant when copper electrodes are used?

Cu²+ ions removed at the cathode are replaced by Cu²+ ions formed at the anode.

6.

How can electrolysis of copper sulfate solution using copper electrodes be used to purify impure copper?

Use impure copper as the anode and pure copper as the cathode; copper dissolves from the anode and is deposited as pure copper at the cathode.

3.31 Core Practical: Investigate the electrolysis of copper sulfate solution with inert electrodes and copper electrodes

1.

How could you investigate the electrolysis of copper sulfate solution using inert electrodes?

Connect copper sulfate solution to a d.c. supply using inert electrodes, pass current and observe the products at each electrode.

2.

How could you investigate the electrolysis of copper sulfate solution using copper electrodes?

Use copper electrodes with copper sulfate solution and pass current while observing and measuring the electrodes.

3.

What observations could be made at the cathode when copper sulfate solution is electrolysed using inert electrodes?

A copper coating forms on the cathode.

4.

What observations could be made at the anode when copper sulfate solution is electrolysed using inert electrodes?

Gas bubbles of oxygen form at the anode.

5.

How would the results differ between electrolysis of copper sulfate solution using inert electrodes and using copper electrodes?

With inert electrodes, copper is deposited at the cathode and oxygen forms at the anode; with copper electrodes, copper transfers from anode to cathode.

6.

What measurements could be made before and after electrolysis of copper sulfate solution to investigate changes at the electrodes?

Measure electrode masses before and after electrolysis and observe any gas formation or deposits.

4.1 Deduce the relative reactivity of some metals, by their reactions with water, acids and salt solutions

1.

How can the reactions of metals with water be used to compare their relative reactivities?

More reactive metals react more vigorously with water.

2.

How can the reactions of metals with dilute acids be used to establish a relative reactivity order?

More reactive metals generally react more readily with dilute acids and produce hydrogen.

3.

How can a metal displacement reaction with a salt solution be used to compare the reactivity of two metals?

A more reactive metal displaces a less reactive metal from its salt solution.

4.

If magnesium reacts with copper sulfate solution but copper does not react with magnesium sulfate solution, which metal is more reactive?

Magnesium is more reactive than copper.

5.

If metal A displaces metal B from a solution of B's salt, what does this indicate about the relative reactivities of A and B?

A is more reactive than B.

6.

Why can observations from reactions with water, acids and salt solutions be used to deduce the relative reactivity of metals?

More reactive metals react more readily and displace less reactive metals.

4.2 Explain displacement reactions as redox reactions, in terms of gain or loss of electrons

1.

Why is a metal displacement reaction classified as a redox reaction?

Electrons are transferred between the reacting species.

2.

In a displacement reaction, what happens to the electrons lost by the more reactive metal?

They are transferred to the less reactive metal ions.

3.

What happens to the metal ions when they are displaced from a compound by a more reactive metal?

They gain electrons and become metal atoms.

4.

In the reaction Mg + Cu²⁺ → Mg²⁺ + Cu, which substance is oxidised and which substance is reduced?

Mg is oxidised; Cu²+ is reduced.

5.

Why is the oxidation of one metal accompanied by the reduction of another species in a displacement reaction?

Electron loss and electron gain occur together.

6.

How can electron transfer be used to explain why magnesium displaces copper from copper sulfate?

Mg loses electrons to form Mg²+, while Cu²+ gains electrons to form Cu.

4.3 Explain the reactivity series of metals (potassium, sodium, calcium, magnesium, aluminium, (carbon), zinc, iron, (hydrogen), copper, silver, gold) in terms of the reactivity of the metals with water and dilute acids and that these reactions show the relative tendency of metal atoms to form cations

1.

What is the Edexcel GCSE Chemistry reactivity series from most reactive to least reactive?

Potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold.

2.

What does a metal's position in the reactivity series indicate about its tendency to form positive ions?

A higher position means a greater tendency to form cations.

3.

Why do metals above hydrogen in the reactivity series generally react with dilute acids to produce hydrogen?

They are above hydrogen and readily lose electrons to form ions while H+ ions are reduced to H2.

4.

Which metals in the Edexcel reactivity series are less reactive than hydrogen?

Copper, silver and gold.

5.

Why is potassium more reactive than copper according to the reactivity series?

Potassium loses electrons more readily and forms K+ ions more easily.

6.

How can reactions with water and dilute acids provide evidence for the order of metals in the reactivity series?

More reactive metals react more readily with water or acids, providing evidence for their tendency to form cations.

4.4 Recall that: a most metals are extracted from ores found in the Earth's crust b unreactive metals are found in the Earth's crust as the uncombined elements

1.

Where are most metals found naturally before they are extracted?

Ores in the Earth's crust.

2.

What is an ore?

A rock containing enough of a metal compound for the metal to be extracted economically.

3.

Why are most metals found as compounds rather than as uncombined elements in the Earth's crust?

They readily react with other substances and form compounds.

4.

Why can unreactive metals sometimes be found naturally as uncombined elements?

They are too unreactive to react readily with other elements.

5.

Which type of metals are most likely to be found in the Earth's crust as uncombined elements?

Unreactive metals.

6.

Why can gold occur naturally as an uncombined element?

Gold is very unreactive, so it can remain as the uncombined element.

4.5 Explain oxidation as the gain of oxygen and reduction as the loss of oxygen

1.

What is oxidation in terms of oxygen?

Gain of oxygen.

2.

What is reduction in terms of oxygen?

Loss of oxygen.

3.

What happens to a substance when it is oxidised in an oxygen-based definition?

It gains oxygen.

4.

What happens to a substance when it is reduced in an oxygen-based definition?

It loses oxygen.

5.

In the reaction copper oxide + carbon → copper + carbon dioxide, which substance is reduced?

Copper oxide is reduced to copper.

6.

In the reaction copper + oxygen → copper oxide, which substance is oxidised?

Copper is oxidised to copper oxide.

4.6 Recall that the extraction of metals involves reduction of ores

1.

Why does extracting a metal from its ore generally involve reduction?

The metal compound must be reduced to obtain the metal.

2.

What happens to the metal compound in an ore during the reduction stage of metal extraction?

The metal ion is converted into the metal.

3.

What does reduction mean when extracting a metal from a metal oxide?

Oxygen is removed from the metal oxide.

4.

Why must a metal ion in an ore gain electrons to become the metal?

Metal ions must gain electrons to become neutral metal atoms.

5.

What happens to the oxygen in a metal oxide when the metal oxide is reduced?

Oxygen is removed from the metal oxide.

6.

Why is the extraction of a metal from its oxide described as a reduction process?

The metal oxide loses oxygen and forms the metal.

4.7 Explain why the method used to extract a metal from its ore is related to its position in the reactivity series and the cost of the extraction process, illustrated by a heating with carbon (including iron) b electrolysis (including aluminium) (knowledge of the blast furnace is not required)

1.

Why can metals below carbon in the reactivity series be extracted by heating their ores with carbon?

They are less reactive than carbon, so their oxides can be reduced by carbon.

2.

Why cannot aluminium be extracted from its ore by heating it with carbon?

Aluminium is more reactive than carbon, so carbon cannot reduce aluminium oxide.

3.

Why is electrolysis used to extract aluminium from its ore?

Aluminium ions require electrical energy to be reduced.

4.

Why is extracting a very reactive metal by electrolysis generally more expensive than extracting a less reactive metal using carbon?

Electrolysis requires a large amount of electrical energy.

5.

Why can iron be extracted from its ore by heating the ore with carbon?

Iron is below carbon in the reactivity series.

6.

How does a metal's position relative to carbon in the reactivity series determine whether carbon reduction or electrolysis is required for extraction?

Metals below carbon can generally be extracted by carbon reduction; metals above carbon require electrolysis.

4.8 Evaluate alternative biological methods of metal extraction (bacterial and phytoextraction)

1.

What is bacterial extraction of metals?

Using bacteria to help extract metal compounds from ores.

2.

How can bacteria help to extract metals from ores?

Bacteria cause chemical reactions that convert metal compounds into soluble compounds that can be recovered.

3.

What is phytoextraction?

Using plants to absorb metal ions from soil.

4.

How can plants be used to extract metal compounds from contaminated soil?

Grow plants that absorb metal ions, harvest them and recover the metals from the plant material.

5.

What is one advantage of using biological methods to extract metals compared with traditional extraction methods?

Lower energy use and potentially less environmental damage.

6.

What is one disadvantage of using biological methods of metal extraction?

They can be slow and may produce lower yields or require large areas of land.

4.9 Explain how a metal's relative resistance to oxidation is related to its position in the reactivity series

1.

How does a metal's position in the reactivity series affect its resistance to oxidation?

Less reactive metals are more resistant to oxidation.

2.

Why are less reactive metals generally more resistant to oxidation?

They have a lower tendency to lose electrons and react with oxygen.

3.

Which is more resistant to oxidation, gold or potassium, and why?

Gold, because it is less reactive than potassium.

4.

Why does potassium oxidise more readily than gold?

Potassium readily loses electrons and is easily oxidised.

5.

How does the tendency of a metal atom to form a cation relate to its tendency to be oxidised?

A greater tendency to form a cation means a greater tendency to be oxidised.

6.

Why can the reactivity series be used to compare the relative resistance of metals to oxidation?

The reactivity series ranks metals according to how readily they react, including their tendency to be oxidised.

4.10 Evaluate the advantages of recycling metals, including economic implications and how recycling can preserve both the environment and the supply of valuable raw materials

1.

How does recycling metals help to conserve valuable raw materials?

It reduces the need to extract new ores.

2.

How can recycling metals reduce damage to the environment?

It reduces mining, habitat destruction, pollution and waste.

3.

Why can recycling metals reduce the energy needed compared with extracting metals from ores?

Recycling generally uses less energy than extracting metals from ores.

4.

How can recycling metals have economic benefits?

It reduces raw-material and energy costs and can reduce waste-disposal costs.

5.

Why can recycling metals reduce the amount of mining required?

Less new ore needs to be mined.

6.

What are the main environmental and economic advantages of recycling metals?

It conserves raw materials, reduces environmental damage and can save energy and money.

4.11 Describe that a life-cycle assessment for a product involves consideration of the effect on the environment of obtaining the raw materials, manufacturing the product, using the product and disposing of the product when it is no longer useful

1.

What is a life-cycle assessment?

An assessment of the environmental impact of a product throughout its life.

2.

Which four stages of a product's life are considered in a life-cycle assessment?

Obtaining raw materials, manufacture, use and disposal.

3.

What environmental effects of obtaining raw materials are considered in a life-cycle assessment?

Mining, extraction, transport and habitat damage.

4.

What environmental effects of manufacturing a product are considered in a life-cycle assessment?

Energy use, emissions, waste and resource use.

5.

Why does a life-cycle assessment consider how a product is used?

The product may use energy or release pollutants during its lifetime.

6.

Why does a life-cycle assessment consider how a product is disposed of when it is no longer useful?

Disposal can produce waste, require energy or release pollutants.

4.12 Evaluate data from a life cycle assessment of a product

1.

What information can be obtained from data in a life-cycle assessment?

The environmental impacts associated with each stage of the product's life.

2.

How can life-cycle assessment data be used to compare the environmental impacts of two products?

Compare the impacts across all stages and consider the overall effect.

3.

Why should all stages of a product's life be considered when evaluating its environmental impact?

A product may have low impact at one stage but high impact at another.

4.

If one product has a lower environmental impact during manufacture but a higher impact during disposal, why must both stages be considered?

The lower manufacturing impact could be outweighed by the greater disposal impact.

5.

How could numerical data from a life-cycle assessment be used to decide which product has the lower overall environmental impact?

Add or compare the relevant environmental impacts for all stages.

6.

Why might life-cycle assessment data need to be evaluated rather than simply choosing the product with the lowest impact at one stage?

Environmental impact must be considered over the whole life cycle, not one stage alone.

4.13 Recall that chemical reactions are reversible, the use of the symbol ⇌ in equations and that the direction of some reversible reactions can be altered by changing the reaction conditions

1.

What is meant by a reversible chemical reaction?

A reaction that can proceed in both directions.

2.

What symbol is used to represent a reversible chemical reaction?

⇌

3.

What does the symbol ⇌ indicate in a chemical equation?

The reaction can proceed forwards and backwards.

4.

What happens to the products and reactants in a reversible reaction?

Products can react to reform reactants.

5.

How can changing reaction conditions affect the direction of a reversible reaction?

Changing temperature, pressure or concentration can shift the equilibrium position.

6.

What is the difference between a reversible reaction and a reaction that only proceeds in one direction?

A reversible reaction can proceed in both directions, whereas an irreversible reaction proceeds effectively in one direction.

4.14 Explain what is meant by dynamic equilibrium

1.

What is meant by dynamic equilibrium in a reversible reaction?

A state in a closed system where the forward and backward reactions occur at equal rates.

2.

What happens to the forward reaction rate when a reversible reaction reaches dynamic equilibrium?

It becomes equal to the backward reaction rate.

3.

What happens to the backward reaction rate when a reversible reaction reaches dynamic equilibrium?

It becomes equal to the forward reaction rate.

4.

Why does the concentration of reactants and products remain constant at dynamic equilibrium?

Reactants are being converted into products at the same rate that products are converted back into reactants.

5.

Why does a reversible reaction continue to occur even after dynamic equilibrium has been reached?

Both reactions continue to occur.

6.

What condition must the forward and backward reaction rates satisfy for dynamic equilibrium to exist?

The forward and backward reaction rates must be equal.

4.15 Describe the formation of ammonia as a reversible reaction between nitrogen (extracted from the air) and hydrogen (obtained from natural gas) and that it can reach a dynamic equilibrium

1.

What are the reactants used to manufacture ammonia in the Haber process?

Nitrogen and hydrogen.

2.

Where is the nitrogen used in the Haber process obtained from?

From the air.

3.

Where is the hydrogen used in the Haber process obtained from?

From natural gas.

4.

What is the balanced equation for the reversible formation of ammonia from nitrogen and hydrogen?

N2 + 3H2 ⇌ 2NH3

5.

Why can the reaction between nitrogen and hydrogen reach dynamic equilibrium?

The reaction is reversible and both forward and backward reactions occur.

6.

What happens to the concentrations of nitrogen, hydrogen and ammonia when the Haber process reaches dynamic equilibrium?

Their concentrations remain constant.

4.16 Recall the conditions for the Haber process as: a temperature 450 °C b pressure 200 atmospheres c iron catalyst

1.

What temperature is used in the Haber process?

450 °C

2.

What pressure is used in the Haber process?

200 atmospheres

3.

Which catalyst is used in the Haber process?

Iron.

4.

What are the three industrial conditions used in the Haber process?

450 °C, 200 atmospheres and an iron catalyst.

5.

Why is an iron catalyst used in the Haber process?

It increases the rate of reaction.

6.

What temperature, pressure and catalyst are required industrially for the Haber process?

450 °C, 200 atmospheres and an iron catalyst.

4.17 Predict how the position of a dynamic equilibrium is affected by changes in: a temperature b pressure c concentration

1.

For the equilibrium N₂(g) + 3H₂(g) ⇌ 2NH₃(g), what happens to the position of equilibrium when the concentration of nitrogen is increased?

Shifts to the right, increasing ammonia production.

2.

For the equilibrium N₂(g) + 3H₂(g) ⇌ 2NH₃(g), what happens to the position of equilibrium when the concentration of ammonia is decreased?

Shifts to the right to replace some of the ammonia removed.

3.

For the equilibrium N₂(g) + 3H₂(g) ⇌ 2NH₃(g), what happens to the position of equilibrium when the pressure is increased?

Shifts to the right because there are fewer moles of gas on the right.

4.

For the equilibrium N₂(g) + 3H₂(g) ⇌ 2NH₃(g), what happens to the position of equilibrium when the pressure is decreased?

Shifts to the left because there are more moles of gas on the left.

5.

For the exothermic equilibrium N₂(g) + 3H₂(g) ⇌ 2NH₃(g), what happens to the position of equilibrium when the temperature is increased?

Shifts to the left because the forward reaction is exothermic.

6.

For the exothermic equilibrium N₂(g) + 3H₂(g) ⇌ 2NH₃(g), what happens to the position of equilibrium when the temperature is decreased?

Shifts to the right because the forward reaction releases heat.

5.1C Recall that most metals are transition metals and that their typical properties include: a high melting point b high density c the formation of coloured compounds d catalytic activity of the metals and their compounds as exemplified by iron

1.

What are the four typical properties of transition metals required by Edexcel GCSE Chemistry?

High melting points, high densities, coloured compounds and catalytic activity.

2.

Why do transition metals generally have high melting points?

Strong metallic bonding requires a large amount of energy to overcome.

3.

Why do transition metals generally have high densities?

Their atoms are relatively closely packed and have high relative atomic masses.

4.

What property of many transition metal compounds makes them useful in applications involving colour?

Many transition metal compounds are brightly coloured.

5.

What is catalytic activity, and why are transition metals useful as catalysts?

Catalytic activity is the ability to speed up chemical reactions without being used up; transition metals have suitable electronic structures that allow them to act as catalysts.

6.

How does iron demonstrate the catalytic activity of transition metals and their compounds?

Iron is used as a catalyst in some industrial reactions, for example the Haber process.

5.2C Recall that the oxidation of metals results in corrosion

1.

What is corrosion?

Corrosion is the gradual destruction of a metal by chemical reactions with its surroundings.

2.

What happens to a metal when it is oxidised during corrosion?

The metal loses electrons and is converted into oxidised compounds.

3.

Why does the oxidation of metals result in corrosion?

Oxidation changes the metal into compounds that weaken and damage the original metal.

4.

What is the relationship between oxidation and corrosion of metals?

Corrosion of metals is caused by their oxidation.

5.

Why does corrosion cause a metal object to deteriorate over time?

The metal is gradually converted into corrosion products, weakening the object.

6.

What chemical process is responsible for the corrosion of metals?

Oxidation of the metal.

5.3C Explain how rusting of iron can be prevented by: a exclusion of oxygen b exclusion of water c sacrificial protection

1.

What two substances are required for iron to rust?

Oxygen and water.

2.

How does excluding oxygen prevent iron from rusting?

Without oxygen, the oxidation of iron cannot occur.

3.

How does excluding water prevent iron from rusting?

Without water, the electrochemical processes needed for rusting cannot occur.

4.

How does painting an iron object help prevent rusting?

Paint forms a barrier that prevents oxygen and water reaching the iron.

5.

What is sacrificial protection?

Sacrificial protection is using a more reactive metal to protect iron by allowing the more reactive metal to oxidise instead.

6.

Why can a more reactive metal protect iron from rusting through sacrificial protection?

The more reactive metal is oxidised preferentially, preventing the iron from being oxidised.

5.4C Explain how electroplating can be used to improve the appearance and/or the resistance to corrosion of metal objects

1.

What is electroplating?

Electroplating is coating an object with a thin layer of another metal using electrolysis.

2.

How does electroplating improve the appearance of a metal object?

It gives the object a smooth, attractive surface of the plating metal.

3.

How does electroplating protect a metal object from corrosion?

The coating prevents oxygen and water reaching the underlying metal.

4.

Why can coating an object with a less reactive metal increase its resistance to corrosion?

A less reactive metal is less easily oxidised.

5.

What happens to metal ions during the electroplating of an object?

Metal ions gain electrons at the cathode and form a layer of metal on the object.

6.

Why is electroplating useful for objects that need both an attractive appearance and corrosion resistance?

It can make the object more attractive while also protecting it from corrosion.

5.5C Explain, using models, why converting pure metals into alloys often increases the strength of the product

1.

What is an alloy?

An alloy is a mixture containing a metal and one or more other elements.

2.

Why does adding different-sized atoms to a pure metal make the metal stronger?

Different-sized atoms distort the regular arrangement of metal atoms.

3.

Why can layers of atoms slide more easily in a pure metal than in an alloy?

In a pure metal, identical layers of atoms can slide over each other more easily.

4.

How does the presence of different-sized atoms disrupt the regular arrangement of atoms in an alloy?

Different-sized atoms disrupt the regular arrangement of atoms.

5.

Why does the disrupted structure of an alloy make it harder for layers of atoms to slide?

The distorted layers cannot slide as easily past one another.

6.

How does alloying a metal usually affect its strength compared with the pure metal?

Alloying generally makes a metal stronger.

5.6C Explain why iron is alloyed with other metals to produce alloy steels

1.

Why is iron alloyed with other metals to produce alloy steels?

To change and improve its properties.

2.

How does alloying change the properties of iron?

It can increase strength, hardness and other useful properties.

3.

Why are alloy steels generally stronger or more useful than pure iron?

Alloy steels can be stronger, harder and more resistant to corrosion than pure iron.

4.

Why can different metals be added to iron to produce steels with different properties?

Different metals produce different combinations of properties.

5.

How does alloying iron affect the ability of layers of atoms to slide?

Different-sized atoms disrupt the regular arrangement and make the layers harder to slide.

6.

Why are alloy steels preferred to pure iron for many applications?

They provide properties better suited to specific applications.

5.7C Explain how the uses of metals are related to their properties (and vice versa), including aluminium, copper and gold and their alloys including magnalium and brass

1.

Why is aluminium useful for aircraft despite being a metal?

Aluminium has a low density, making it useful for aircraft.

2.

Why is copper widely used for electrical wiring?

Copper has high electrical conductivity and is ductile.

3.

Why is gold useful for jewellery?

Gold is unreactive, attractive and malleable.

4.

Why is magnalium useful when a strong but relatively lightweight material is required?

Magnalium combines relatively low density with increased strength.

5.

Why is brass useful for making objects such as musical instruments and fittings?

Brass is strong, hard, corrosion-resistant and attractive.

6.

How are the uses of aluminium, copper and gold determined by their physical and chemical properties?

Their uses depend on properties such as density, conductivity, strength, malleability and resistance to corrosion.

5.8C Calculate the concentration of solutions in mol dm–3 and convert concentration in g dm–3 into mol dm–3 and vice versa

1.

What equation is used to calculate concentration in mol dm⁻³ from the number of moles and volume of solution?

Concentration = moles ÷ volume. c = n / V

2.

What concentration in mol dm⁻³ is produced by dissolving 0.50 mol of solute in 2.0 dm³ of solution?

0.50 ÷ 2.0 = 0.25 mol dm⁻³

3.

How can a concentration in g dm⁻³ be converted into mol dm⁻³?

Concentration in mol dm⁻³ = concentration in g dm⁻³ ÷ Mr.

4.

What concentration in mol dm⁻³ is produced by a solution containing 5.85 g of sodium chloride (Mr = 58.5) in 1.00 dm³?

5.85 ÷ 58.5 = 0.100 mol dm⁻³

5.

How can a concentration in mol dm⁻³ be converted into g dm⁻³?

Concentration in g dm⁻³ = concentration in mol dm⁻³ × Mr.

6.

What concentration in g dm⁻³ corresponds to a 0.250 mol dm⁻³ solution of sodium hydroxide (Mr = 40.0)?

0.250 × 40.0 = 10.0 g dm⁻³

5.9C Core Practical: Carry out an accurate acid-alkali titration, using burette, pipette and a suitable indicator

1.

What apparatus is used to measure a fixed volume of alkali accurately during an acid-alkali titration?

A pipette.

2.

Why is a burette used to add the acid accurately during an acid-alkali titration?

A burette allows an accurate variable volume of acid to be delivered.

3.

Why is an indicator used during an acid-alkali titration?

To show when the neutralisation endpoint has been reached.

4.

What is the purpose of carrying out a rough titration before obtaining accurate titres?

To estimate the approximate titre and identify a suitable range for accurate titres.

5.

Why should concordant titres be obtained in an accurate acid-alkali titration?

To show that the results are consistent and reliable.

6.

How is the titre calculated from the initial and final burette readings?

Titre = final burette reading − initial burette reading.

5.10C Carry out simple calculations using the results of titrations to calculate an unknown concentration of a solution or an unknown volume of solution required

1.

How can titration results be used to calculate the unknown concentration of an acid?

Use the titre to calculate moles of the known solution, use the balanced equation to find the mole ratio, then calculate the unknown concentration. concentration = moles ÷ volume

2.

How can titration results be used to calculate the unknown concentration of an alkali?

Use the same method, using the known acid or alkali concentration and the balanced equation.

3.

What information from a titration is needed to calculate an unknown concentration?

The balanced equation, known concentration, known volume and titre of the unknown solution.

4.

If 25.0 cm³ of 0.100 mol dm⁻³ HCl exactly neutralises 20.0 cm³ of NaOH, what is the concentration of the NaOH?

HCl + NaOH → NaCl + H₂O; moles HCl = 0.100 × 25.0/1000 = 0.00250 mol; 1:1 ratio, so moles NaOH = 0.00250 mol; concentration NaOH = 0.00250 ÷ 20.0/1000 = 0.125 mol dm⁻³

5.

How can a balanced chemical equation be used when calculating an unknown concentration from titration results?

It gives the mole ratio between the reacting substances.

6.

How can titration results be used to calculate the volume of solution required for a reaction?

Use the mole ratio from the balanced equation to calculate the required moles, then use: volume = moles ÷ concentration

5.11C Calculate the percentage yield of a reaction from the actual yield and the theoretical yield

1.

What equation is used to calculate percentage yield?

Percentage yield = (actual yield ÷ theoretical yield) × 100

2.

What is meant by the theoretical yield of a reaction?

The maximum amount of product that could be produced.

3.

What is meant by the actual yield of a reaction?

The amount of product actually obtained.

4.

What is the percentage yield when the theoretical yield is 50 g and the actual yield is 40 g?

(40 ÷ 50) × 100 = 80%

5.

What does a percentage yield of 100% mean?

The actual yield equals the theoretical yield.

6.

Why is the actual yield of a reaction usually less than the theoretical yield?

Reactions may be incomplete, products may be lost, or side reactions may occur.

5.12C Describe that the actual yield of a reaction is usually less than the theoretical yield and that the causes of this include: a incomplete reactions b practical losses during the experiment c competing, unwanted reactions (side reactions)

1.

Why is the actual yield of a reaction usually less than the theoretical yield?

Reactions may be incomplete, product may be lost, or unwanted side reactions may occur.

2.

How can an incomplete reaction reduce the actual yield of a product?

Some reactants remain unreacted, so less product is formed.

3.

How can practical losses during an experiment reduce the actual yield?

Product can be lost during filtration, transfer, washing or purification.

4.

How can side reactions reduce the actual yield of the desired product?

Some reactants form unwanted products instead of the desired product.

5.

Why can transferring a product between containers cause a practical loss?

Some product remains in the original container or equipment.

6.

What are the three main causes of an actual yield being lower than the theoretical yield?

Incomplete reactions, practical losses and side reactions.

5.13C Recall the atom economy of a reaction forming a desired product

1.

What is atom economy?

Atom economy is the proportion of reactant atoms incorporated into the desired product.

2.

What does a high atom economy indicate about a chemical reaction?

A high proportion of the reactants becomes the desired product.

3.

What does a low atom economy indicate about a chemical reaction?

A larger proportion of reactant atoms becomes waste or unwanted products.

4.

Why is atom economy important when choosing a reaction pathway for manufacturing a product?

It reduces waste and makes better use of raw materials.

5.

What happens to the atoms that are not incorporated into the desired product?

They become part of by-products or waste.

6.

Why does a reaction producing fewer unwanted products generally have a higher atom economy?

Fewer atoms are wasted in unwanted products.

5.14C Calculate the atom economy of a reaction forming a desired product

1.

What equation is used to calculate the percentage atom economy of a reaction?

Percentage atom economy = (Mr of desired product ÷ total Mr of reactants) × 100

2.

What is the atom economy of a reaction with a desired product Mr of 80 and total Mr of reactants of 100?

(80 ÷ 100) × 100 = 80%

3.

What is the atom economy of the reaction CaCO₃ → CaO + CO₂ if CaO is the desired product? (Ar: Ca = 40, C = 12, O = 16)

Mr(CaCO₃) = 40 + 12 + 48 = 100; Mr(CaO) = 40 + 16 = 56; Atom economy = (56 ÷ 100) × 100 = 56%

4.

How can the relative formula masses of the reactants and desired product be used to calculate atom economy?

Calculate the Mr of the desired product and the total Mr of the reactants, then apply the atom economy equation.

5.

Why must the desired product be identified before calculating atom economy?

The desired product determines which reactant atoms are counted as being used usefully.

6.

What does an atom economy of 75% mean?

75% of the atoms in the reactants are incorporated into the desired product.

5.15C Explain why a particular reaction pathway is chosen to produce a specified product, given appropriate data such as atom economy, yield, rate, equilibrium position and usefulness of by-products

1.

Why might a reaction with a high atom economy be preferred in industry?

Less waste is produced and raw materials are used more efficiently.

2.

Why might a reaction with a high percentage yield be preferred?

More of the desired product is obtained from a given amount of reactants.

3.

Why might the rate of a reaction affect the choice of industrial reaction pathway?

A faster reaction can reduce production time and costs.

4.

Why might the position of equilibrium affect the choice of reaction pathway?

It affects how much product can be obtained at equilibrium.

5.

Why can useful by-products make a reaction pathway more attractive?

They can provide an additional useful product and reduce waste.

6.

How should atom economy, yield, rate, equilibrium position and useful by-products be considered when choosing an industrial reaction pathway?

The pathway should balance environmental, economic and practical factors, including high atom economy, high yield, suitable rate, favourable equilibrium and useful by-products.

5.16C Describe the molar volume, of any gas at room temperature and pressure, as the volume occupied by one mole of molecules of any gas at room temperature and pressure (The molar volume will be provided as 24 dm3 or 24000 cm3 in calculations where it is required)

1.

What is the molar volume of any gas at room temperature and pressure?

The volume occupied by one mole of gas at room temperature and pressure.

2.

What volume does one mole of gas occupy at room temperature and pressure in dm³?

24 dm³

3.

What volume does one mole of gas occupy at room temperature and pressure in cm³?

24 000 cm³

4.

What volume would 2.5 mol of a gas occupy at room temperature and pressure?

2.5 × 24 = 60 dm³

5.

How can the number of moles of a gas be calculated from its volume at room temperature and pressure?

Moles = volume ÷ molar volume.

6.

Why can the molar volume be used for different gases at room temperature and pressure?

Equal numbers of gas particles occupy the same volume at the same temperature and pressure.

5.17C Use the molar volume and balanced equations in calculations involving the masses of solids and volumes of gases

1.

What volume of gas at room temperature and pressure is produced when 0.50 mol of gas is formed?

0.50 × 24 = 12 dm³

2.

How can a balanced equation be used to calculate the volume of gas produced from a known mass of a solid?

Convert the mass of the solid to moles, use the equation to find the moles of gas, then multiply by 24 dm³.

3.

What volume of CO₂ at room temperature and pressure is produced when 10.0 g of CaCO₃ decomposes according to CaCO₃ → CaO + CO₂? (Mr CaCO₃ = 100)

Moles CaCO₃ = 10.0 ÷ 100 = 0.100 mol; 1:1 ratio, so CO₂ = 0.100 mol; Volume = 0.100 × 24 = 2.40 dm³

4.

What mass of calcium carbonate is required to produce 4.80 dm³ of CO₂ at room temperature and pressure according to CaCO₃ → CaO + CO₂?

Moles CO₂ = 4.80 ÷ 24 = 0.200 mol; 1:1 ratio, so CaCO₃ = 0.200 mol; Mass = 0.200 × 100 = 20.0 g

5.

Why must the equation be balanced before using it to calculate gas volumes?

The coefficients give the mole ratio between substances.

6.

How can the molar volume of a gas be combined with relative formula mass to calculate the mass of a solid involved in a reaction?

Convert gas volume to moles using the molar volume, use the mole ratio, then convert moles of solid to mass using Mr.

5.18C Use Avogadro's law to calculate volumes of gases involved in a gaseous reaction, given the relevant equation

1.

What does Avogadro's law state about the volumes of gases involved in a reaction at the same temperature and pressure?

Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

2.

According to N₂ + 3H₂ → 2NH₃, what volume of hydrogen reacts with 10 dm³ of nitrogen when all gases are measured at the same temperature and pressure?

N₂ : H₂ = 1 : 3, so 30 dm³ H₂

3.

According to N₂ + 3H₂ → 2NH₃, what volume of ammonia is produced from 15 dm³ of nitrogen when all gases are measured at the same temperature and pressure?

N₂ : NH₃ = 1 : 2, so 30 dm³ NH₃

4.

According to 2H₂ + O₂ → 2H₂O, what volume of oxygen reacts with 40 dm³ of hydrogen when all gases are measured at the same temperature and pressure?

H₂ : O₂ = 2 : 1, so 20 dm³ O₂

5.

How can the coefficients in a balanced gaseous equation be used to calculate reacting gas volumes?

The coefficients give the ratio of reacting gas volumes.

6.

Why can gas volumes be compared directly using the mole ratios in a balanced equation when temperature and pressure are constant?

At the same temperature and pressure, gas volume is proportional to the number of moles.

5.19C Describe the Haber process as a reversible reaction between nitrogen and hydrogen to form ammonia

1.

What reversible reaction occurs in the Haber process?

Nitrogen reacts reversibly with hydrogen to form ammonia.

2.

What are the two reactants used to produce ammonia in the Haber process?

Nitrogen and hydrogen.

3.

What is the balanced symbol equation for the Haber process?

N₂ + 3H₂ ⇌ 2NH₃

4.

Where is the nitrogen used in the Haber process obtained from?

From the air.

5.

Where is the hydrogen used in the Haber process obtained from?

Mainly from natural gas.

6.

Why does the Haber process reach a dynamic equilibrium?

The forward and reverse reactions can both occur, leading to equal forward and backward rates.

5.20C Predict how the rate of attainment of equilibrium is affected by: a changes in temperature b changes in pressure c changes in concentration d use of a catalyst

1.

How does increasing the temperature affect the rate at which equilibrium is reached in the Haber process?

Increasing temperature increases the rate of reaching equilibrium.

2.

How does increasing the pressure affect the rate at which equilibrium is reached in the Haber process?

Increasing pressure generally increases the rate of reaching equilibrium for reacting gases.

3.

How does increasing the concentration of a reactant affect the rate at which equilibrium is reached?

Increasing reactant concentration increases the rate of reaching equilibrium.

4.

How does increasing the concentration of a product affect the rate at which equilibrium is reached?

Increasing product concentration increases the rate of the reverse reaction and allows equilibrium to be reached sooner.

5.

How does adding a catalyst affect the rate at which equilibrium is reached?

It increases the rate of reaching equilibrium.

6.

Why does a catalyst allow equilibrium to be reached more quickly without changing the position of equilibrium?

It lowers the activation energy for both forward and reverse reactions equally, without changing the equilibrium position.

5.21C Explain how, in industrial reactions, including the Haber process, conditions used are related to: a the availability and cost of raw materials and energy supplies b the control of temperature, pressure and catalyst used produce an acceptable yield in an acceptable time

1.

Why must the cost and availability of raw materials be considered when choosing industrial reaction conditions?

Expensive or limited raw materials increase production costs.

2.

Why must the energy requirements of an industrial reaction be considered when selecting its operating conditions?

High energy use increases operating costs.

3.

Why is a compromise temperature used in the Haber process rather than simply using the temperature that gives the highest equilibrium yield?

A lower temperature would give a higher equilibrium yield but a slower rate, so a compromise temperature is used.

4.

Why is a compromise pressure used in the Haber process rather than simply using an extremely high pressure?

Higher pressure favours ammonia formation but very high pressures are expensive and hazardous.

5.

Why is a catalyst used in the Haber process?

It increases the rate without changing the equilibrium position.

6.

Why must industrial reaction conditions produce an acceptable yield in an acceptable time?

Industry must balance yield, rate, cost, safety and energy use.

5.22C Recall that fertilisers may contain nitrogen, phosphorus and potassium compounds to promote plant growth

1.

Which three elements are commonly supplied by fertilisers to promote plant growth?

Nitrogen, phosphorus and potassium.

2.

Why are nitrogen compounds included in fertilisers?

Nitrogen is needed to make proteins and other important compounds.

3.

Why are phosphorus compounds included in fertilisers?

Phosphorus is needed for healthy growth and processes such as energy transfer.

4.

Why are potassium compounds included in fertilisers?

Potassium supports healthy plant growth and enzyme activity.

5.

What are the three main types of nutrients supplied by fertilisers?

Nitrogen, phosphorus and potassium compounds.

6.

Why are fertilisers added to soil?

To replace essential mineral ions and increase crop growth.

5.23C Describe how ammonia reacts with nitric acid to produce a salt that is used as a fertiliser

1.

What salt is produced when ammonia reacts with nitric acid?

Ammonium nitrate.

2.

What is the word equation for the reaction between ammonia and nitric acid?

Ammonia + nitric acid → ammonium nitrate

3.

What is the balanced symbol equation for the reaction between ammonia and nitric acid?

NH₃ + HNO₃ → NH₄NO₃

4.

Why is ammonium nitrate used as a fertiliser?

It provides nitrogen needed for plant growth.

5.

What type of reaction occurs when ammonia reacts with nitric acid?

Neutralisation.

6.

Which ions are present in ammonium nitrate?

NH₄⁺ and NO₃⁻.

5.24C Describe and compare: a the laboratory preparation of ammonium sulfate from ammonia solution and dilute sulfuric acid on a small scale b the industrial production of ammonium sulfate, used as a fertiliser, in which several stages are required to produce ammonia and sulfuric acid from their raw materials and the production is carried out on a much larger scale (details of the industrial production of sulfuric acid are not required)

1.

How is ammonium sulfate prepared in the laboratory from ammonia solution and dilute sulfuric acid?

Ammonia solution is reacted with dilute sulfuric acid, with the correct proportions found by titration, then the solution is concentrated and crystallised.

2.

What are the reactants used to prepare ammonium sulfate in the laboratory?

Ammonia solution and dilute sulfuric acid.

3.

What is the balanced symbol equation for the reaction between ammonia and sulfuric acid?

2NH₃ + H₂SO₄ → (NH₄)₂SO₄

4.

How does the scale of laboratory production of ammonium sulfate differ from industrial production?

Laboratory production is small-scale; industrial production is large-scale.

5.

Why does industrial production of ammonium sulfate require several stages before the final salt is produced?

Raw materials must first be processed to obtain ammonia and sulfuric acid.

6.

How do the laboratory and industrial methods of producing ammonium sulfate differ in scale and process?

The laboratory method directly reacts prepared chemicals on a small scale, whereas industrial production involves multiple large-scale stages.

5.25C Recall that a chemical cell produces a voltage until one of the reactants is used up

1.

What does a chemical cell produce while it is operating?

A voltage.

2.

Why does a chemical cell eventually stop producing a voltage?

One of the reactants is eventually used up.

3.

What happens when one of the reactants in a chemical cell is used up?

The chemical reaction can no longer continue effectively and the voltage falls to zero.

4.

What type of energy conversion occurs in a chemical cell?

Chemical energy → electrical energy.

5.

What determines how long a chemical cell can continue producing a voltage?

The amount of reactants available.

6.

Why cannot a chemical cell produce a voltage indefinitely?

The reactants are consumed during the chemical reaction.

5.26C Recall that in a hydrogen–oxygen fuel cell hydrogen and oxygen are used to produce a voltage and water is the only product

1.

Which two reactants are used in a hydrogen–oxygen fuel cell?

Hydrogen and oxygen.

2.

What product is formed in a hydrogen–oxygen fuel cell?

Water.

3.

What is the overall word equation for a hydrogen–oxygen fuel cell?

Hydrogen + oxygen → water

4.

What is the overall balanced symbol equation for a hydrogen–oxygen fuel cell?

2H₂ + O₂ → 2H₂O

5.

Why does a hydrogen–oxygen fuel cell produce a voltage?

Chemical reactions in the cell transfer energy to an electrical circuit.

6.

Why is water described as the only product of a hydrogen–oxygen fuel cell?

The overall reaction produces water without forming carbon dioxide.

5.27C Evaluate the strengths and weaknesses of fuel cells for given uses

1.

What is one advantage of using a hydrogen–oxygen fuel cell compared with a fossil-fuel-powered system?

They produce electricity with water as the direct product and no carbon dioxide at the point of use.

2.

Why can hydrogen–oxygen fuel cells be considered environmentally beneficial when the hydrogen is produced using renewable energy?

Renewable electricity can be used to produce hydrogen without directly using fossil fuels.

3.

What is one disadvantage of storing and transporting hydrogen for use in fuel cells?

Hydrogen is difficult to store and transport because it has a low density and can require high pressure or low temperatures.

4.

Why can the production of hydrogen affect the environmental benefits of fuel cells?

Producing hydrogen from fossil fuels can generate carbon dioxide.

5.

Why are fuel cells useful for applications that require a continuous supply of electrical energy?

They can continuously generate electricity while reactants are supplied.

6.

What factors should be considered when evaluating whether fuel cells are suitable for a particular use?

Cost, fuel availability, storage, infrastructure, efficiency, emissions and intended use.

6.1 Explain why some elements can be classified as alkali metals (group 1), halogens (group 7) or noble gases (group 0), based on their position in the periodic table

1.

Why are the elements in group 1 classified as alkali metals?

They are the elements in Group 1.

2.

Why are the elements in group 7 classified as halogens?

They are the elements in Group 7.

3.

Why are the elements in group 0 classified as noble gases?

They are the elements in Group 0.

4.

Which group of the periodic table contains the alkali metals?

Group 1.

5.

Which group of the periodic table contains the halogens?

Group 7.

6.

Which group of the periodic table contains the noble gases?

Group 0.

6.2 Recall that alkali metals: a are soft b have relatively low melting points

1.

What is the physical texture of alkali metals?

They are soft.

2.

How do the melting points of alkali metals compare with those of many other metals?

They have relatively low melting points.

3.

Why can alkali metals be described as soft metals?

Their atoms are held together less strongly than in many other metals.

4.

Which two physical properties of alkali metals are characteristic of this group?

Soft and relatively low melting points.

5.

What happens to the melting point of an alkali metal compared with many transition metals?

They are generally lower.

6.

Why can alkali metals be cut relatively easily?

Their metallic bonding is relatively weak.

6.3 Describe the reactions of lithium, sodium and potassium with water

1.

What happens when lithium is added to water?

It floats, moves slowly, fizzes and gradually disappears, forming lithium hydroxide and hydrogen.

2.

What happens when sodium is added to water?

It floats, moves rapidly and melts into a ball, producing sodium hydroxide and hydrogen.

3.

What happens when potassium is added to water?

It reacts very vigorously, usually with a lilac flame, producing potassium hydroxide and hydrogen.

4.

What products are formed when an alkali metal reacts with water?

A metal hydroxide and hydrogen.

5.

What gas is produced when lithium, sodium or potassium reacts with water?

Hydrogen.

6.

What is the general word equation for the reaction between an alkali metal and water?

Metal + water → metal hydroxide + hydrogen

6.4 Describe the pattern in reactivity of the alkali metals, lithium, sodium and potassium, with water; and use this pattern to predict the reactivity of other alkali metals

1.

How does the reactivity of lithium, sodium and potassium with water change down group 1?

Reactivity increases down Group 1.

2.

Which is more reactive with water, lithium or sodium?

Sodium.

3.

Which is more reactive with water, sodium or potassium?

Potassium.

4.

How can the reactivity of an unfamiliar group 1 metal be predicted from its position in the group?

Metals lower down the group are predicted to be more reactive.

5.

Which would be expected to react more vigorously with water, potassium or rubidium?

Rubidium.

6.

What happens to the vigour of the reaction with water as you move down group 1?

The reaction becomes more vigorous.

6.5 Explain this pattern in reactivity in terms of electronic configurations

1.

Why does the reactivity of group 1 metals increase down the group?

The outer electron becomes easier to remove down the group.

2.

How does the distance of the outer electron from the nucleus change down group 1?

It becomes further from the nucleus.

3.

How does increased electron shielding affect the attraction between the nucleus and the outer electron in group 1 metals?

Shielding increases, reducing the attraction between the nucleus and the outer electron.

4.

Why is the outer electron easier to remove from potassium than from lithium?

Potassium has a greater distance and more shielding.

5.

What happens to the first ionisation energy of group 1 metals down the group?

It decreases.

6.

How does the electronic configuration of group 1 metals explain their tendency to form +1 ions?

They each have one electron in their outer shell and tend to lose it to form +1 ions.

6.6 Recall the colours and physical states of chlorine, bromine and iodine at room temperature

1.

What colour and physical state is chlorine at room temperature?

Yellow-green gas.

2.

What colour and physical state is bromine at room temperature?

Red-brown liquid.

3.

What colour and physical state is iodine at room temperature?

Grey-black solid.

4.

Which of chlorine, bromine and iodine is a yellow-green gas at room temperature?

Chlorine.

5.

Which of chlorine, bromine and iodine is a red-brown liquid at room temperature?

Bromine.

6.

Which of chlorine, bromine and iodine is a grey-black solid at room temperature?

Iodine.

6.7 Describe the pattern in the physical properties of the halogens, chlorine, bromine and iodine, and use this pattern to predict the physical properties of other halogens

1.

How does the physical state of the halogens change down group 7?

Gas → liquid → solid.

2.

How does the colour of the halogens change down group 7?

Colours become darker.

3.

How does the melting point of the halogens change down group 7?

Melting point increases.

4.

How does the boiling point of the halogens change down group 7?

Boiling point increases.

5.

What physical state would be predicted for a halogen below iodine in group 7 at room temperature?

A solid.

6.

How can the physical properties of an unfamiliar halogen be predicted from the trends in chlorine, bromine and iodine?

Use the trends in colour, melting point, boiling point and physical state down the group.

6.8 Describe the chemical test for chlorine

1.

How can chlorine gas be chemically tested?

Expose damp blue litmus paper to the gas.

2.

What happens to damp blue litmus paper when it is exposed to chlorine?

It is bleached.

3.

What colour change occurs when chlorine is tested using damp blue litmus paper?

Blue → red → white.

4.

Why does chlorine eventually bleach damp blue litmus paper?

Chlorine reacts with water to form substances that bleach the dye.

5.

What observation confirms the presence of chlorine gas using damp blue litmus paper?

Damp blue litmus paper is bleached.

6.

Why must the litmus paper be damp when testing for chlorine?

Water is needed for chlorine to produce the reactive species responsible for bleaching.

6.9 Describe the reactions of the halogens, chlorine, bromine and iodine, with metals to form metal halides, and use this pattern to predict the reactions of other halogens

1.

What type of compound is formed when a halogen reacts with a metal?

Metal halides.

2.

What compound is formed when chlorine reacts with sodium?

Sodium chloride, NaCl.

3.

What compound is formed when bromine reacts with potassium?

Potassium bromide, KBr.

4.

What compound is formed when iodine reacts with aluminium?

Aluminium iodide, AlI₃.

5.

What is the general word equation for the reaction between a metal and a halogen?

Metal + halogen → metal halide

6.

How can the reactions of chlorine, bromine and iodine with metals be used to predict the reactions of other halogens?

Use the pattern that halogens react with metals to form metal halides.

6.10 Recall that the halogens, chlorine, bromine and iodine, form hydrogen halides which dissolve in water to form acidic solutions, and use this pattern to predict the reactions of other halogens

1.

What type of compounds are formed when chlorine, bromine or iodine reacts with hydrogen?

Hydrogen halides.

2.

What is formed when hydrogen chloride dissolves in water?

Hydrochloric acid solution.

3.

Why do hydrogen halides form acidic solutions when they dissolve in water?

They release hydrogen ions when dissolved in water.

4.

What ions are produced when hydrogen chloride dissolves in water?

H⁺ and Cl⁻ ions.

5.

What acidic solution is formed when hydrogen bromide dissolves in water?

Hydrobromic acid solution.

6.

How can the reactions of chlorine, bromine and iodine with hydrogen be used to predict the reactions of other halogens?

Other halogens are predicted to react with hydrogen to form hydrogen halides that produce acidic solutions in water.

6.11 Describe the relative reactivity of the halogens chlorine, bromine and iodine, as shown by their displacement reactions with halide ions in aqueous solution, and use this pattern to predict the reactions of astatine

1.

How does the reactivity of chlorine, bromine and iodine change down group 7?

Reactivity decreases down the group.

2.

What happens when chlorine is added to potassium bromide solution?

Chlorine displaces bromine: Cl₂ + 2KBr → 2KCl + Br₂

3.

What happens when chlorine is added to potassium iodide solution?

Chlorine displaces iodine: Cl₂ + 2KI → 2KCl + I₂

4.

What happens when bromine is added to potassium iodide solution?

Bromine displaces iodine: Br₂ + 2KI → 2KBr + I₂

5.

Why can chlorine displace bromine and iodine from their halide solutions?

Chlorine is more reactive and gains electrons more readily.

6.

Using the group 7 reactivity trend, which halogens would astatine be expected to be less reactive than?

Astatine would be expected to be less reactive than chlorine, bromine and iodine.

6.12 Explain why these displacement reactions are redox reactions in terms of gain and loss of electrons, identifying which of the substances are oxidised and which are reduced

1.

Why is a halogen displacement reaction classified as a redox reaction?

Electrons are transferred between the halogen and halide ion.

2.

What happens to a halide ion when it is oxidised during a displacement reaction?

It loses electrons.

3.

What happens to a halogen molecule when it is reduced during a displacement reaction?

It gains electrons.

4.

In the reaction Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, which species is oxidised?

Br⁻ is oxidised.

5.

In the reaction Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, which species is reduced?

Cl₂ is reduced.

6.

In the reaction Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, which species gains electrons and which species loses electrons?

Cl₂ gains electrons and Br⁻ loses electrons.

6.13 Explain the relative reactivity of the halogens in terms of electronic configurations

1.

Why does the reactivity of halogens decrease down group 7?

The outer shell is further from the nucleus and shielding increases down the group.

2.

Why does a halogen need to gain one electron to form a stable outer electron shell?

To achieve a full outer electron shell.

3.

How does electron shielding change down group 7?

It increases.

4.

How does the distance between the nucleus and the outer shell change down group 7?

It increases.

5.

Why is chlorine more reactive than bromine?

Chlorine has less shielding and its outer shell is closer to the nucleus.

6.

Why does the ability of a halogen atom to attract an additional electron decrease down group 7?

The attraction between the nucleus and an incoming electron becomes weaker.

6.14 Explain why the noble gases are chemically inert, compared with the other elements, in terms of their electronic configurations

1.

Why are noble gases chemically inert?

They have full outer electron shells.

2.

What is special about the outer electron shell of noble gases?

The outer shell is complete.

3.

Why does a full outer electron shell make noble gases unreactive?

They do not need to gain, lose or share electrons readily.

4.

Why do noble gases have little tendency to gain electrons?

Gaining an electron would disrupt a stable electron arrangement.

5.

Why do noble gases have little tendency to lose electrons?

Losing an electron would disrupt a stable electron arrangement.

6.

How does the electronic configuration of a noble gas explain its lack of chemical reactivity?

Their full outer shells make them chemically unreactive.

6.15 Explain how the uses of noble gases depend on their inertness, low density and/or non-flammability

1.

Why is helium used to fill balloons?

Helium has low density, so it provides lift.

2.

Why is helium suitable for use where a non-flammable gas is required?

Helium is non-flammable.

3.

Why is argon used inside filament lamps?

Argon is inert and prevents the hot filament reacting with oxygen.

4.

Why is argon suitable for use in welding?

Argon is inert, so it does not readily react during welding.

5.

How does the low density of helium make it useful in balloons and airships?

Its density is lower than air, so it provides lift.

6.

How does the inertness of noble gases determine their uses?

Their lack of reactivity, low density and non-flammability make them useful in specific applications.

6.16 Describe the pattern in the physical properties of some noble gases and use this pattern to predict the physical properties of other noble gases

1.

How does the boiling point of noble gases change down group 0?

It increases down Group 0.

2.

How does the melting point of noble gases change down group 0?

It increases down Group 0.

3.

How does the density of noble gases change down group 0?

It increases down Group 0.

4.

How does the relative atomic mass of noble gases change down group 0?

It increases down Group 0.

5.

What physical properties would be predicted for a noble gas below xenon in group 0?

It would be expected to have a higher relative atomic mass, density, melting point and boiling point than xenon.

6.

How can the physical properties of an unfamiliar noble gas be predicted from the trends in group 0?

Extend the trends observed in known noble gases down the group.

7.1 Core Practical: Investigate the effects of changing the conditions of a reaction on the rates of chemical reactions by: a measuring the production of a gas (in the reaction between hydrochloric acid and marble chips) b observing a colour change (in the reaction between sodium thiosulfate and hydrochloric acid)

1.

How can the rate of reaction between hydrochloric acid and marble chips be investigated by measuring the production of a gas?

Place marble chips in hydrochloric acid and collect the carbon dioxide produced using a gas syringe; record gas volume against time.

2.

What measurements should be taken during the reaction between hydrochloric acid and marble chips to determine the rate at which carbon dioxide is produced?

Record the volume of carbon dioxide produced at regular time intervals.

3.

How can a graph of gas volume against time be used to determine the rate of the reaction between hydrochloric acid and marble chips?

Calculate the gradient of the graph; a steeper gradient means a faster reaction.

4.

How can the rate of reaction between sodium thiosulfate and hydrochloric acid be investigated by observing a colour change?

React sodium thiosulfate with hydrochloric acid and time how long it takes for the cross beneath the flask to disappear.

5.

Why does the reaction between sodium thiosulfate and hydrochloric acid allow the rate of reaction to be measured by timing how long it takes for a cross beneath the reaction mixture to disappear?

Sulfur precipitate forms and makes the mixture opaque, so the cross disappears as the reaction proceeds.

6.

How could the effect of changing the concentration of hydrochloric acid on the rate of reaction with marble chips be investigated experimentally?

Keep marble-chip mass and size constant, vary acid concentration, measure gas volume at regular time intervals and compare reaction rates.

7.2 Suggest practical methods for determining the rate of a given reaction

1.

What practical method could be used to determine the rate of a reaction that produces a gas?

Collect the gas in a gas syringe or over water and measure its volume against time.

2.

What practical method could be used to determine the rate of a reaction that produces a precipitate?

Measure the time taken for a precipitate to obscure a mark or use a colorimeter.

3.

How could the rate of a reaction involving a colour change be measured experimentally?

Measure the time taken for the colour to change.

4.

How could the rate of a reaction be determined by measuring the change in mass of the reaction mixture over time?

Measure the mass at regular time intervals and calculate change in mass ÷ time.

5.

How could the rate of a reaction be determined from measurements of the concentration of a reactant over time?

Measure concentration at different times and calculate change in concentration ÷ time.

6.

What measurements are needed to calculate the rate of a chemical reaction from experimental data?

A quantity that changes with time, such as gas volume, mass, concentration or colour intensity.

7.3 Explain how reactions occur when particles collide and that rates of reaction are increased when the frequency and/or energy of collisions is increased

1.

Why must reacting particles collide for a chemical reaction to occur?

Reacting particles must come into contact for bonds to break and form.

2.

What is meant by a successful collision between reacting particles?

A collision with enough energy and the correct orientation to cause reaction.

3.

How does increasing the frequency of collisions between reacting particles affect the rate of a chemical reaction?

More collisions per second increase the number of successful collisions per second.

4.

How does increasing the energy of collisions between reacting particles affect the rate of a chemical reaction?

More particles have enough energy to overcome the activation energy.

5.

Why do not all collisions between reacting particles result in a reaction?

They may have insufficient energy or an unsuitable orientation.

6.

What two factors involving collisions can increase the rate of a chemical reaction?

Collision frequency and collision energy.

7.4 Explain the effects on rates of reaction of changes in temperature, concentration, surface area to volume ratio of a solid and pressure (on reactions involving gases) in terms of frequency and/or energy of collisions between particles

1.

Why does increasing the temperature generally increase the rate of a chemical reaction?

Particles move faster, collide more frequently and more collisions have enough energy.

2.

How does increasing the concentration of a reactant affect the frequency of collisions between reacting particles?

More particles are present in a given volume, increasing collision frequency.

3.

Why does increasing the surface area to volume ratio of a solid increase the rate of reaction?

More particles are exposed at the surface, increasing collision frequency.

4.

Why does increasing the pressure of reacting gases increase the rate of reaction?

Gas particles are closer together, increasing collision frequency.

5.

How does increasing temperature affect both the frequency and energy of collisions between reacting particles?

Both collision frequency and the proportion of particles with sufficient energy increase.

6.

Explain why powdered calcium carbonate reacts faster with hydrochloric acid than the same mass of large calcium carbonate chips.

Powdered calcium carbonate has a greater surface area to volume ratio, so more particles can collide with hydrochloric acid each second.

7.5 Interpret graphs of mass, volume or concentration of reactant or product against time

1.

What does the gradient of a mass-against-time graph represent during a chemical reaction?

Rate of change of mass.

2.

What does a steeper gradient on a volume-of-gas-against-time graph indicate about the rate of reaction?

A faster rate of reaction.

3.

How can the rate of a reaction be determined from a concentration-against-time graph?

Calculate the gradient of the graph.

4.

What does a horizontal section on a graph of product volume against time indicate about a reaction?

The reaction has stopped or no more measurable product is being formed.

5.

How can two mass-against-time graphs be compared to determine which reaction is faster?

The graph with the steeper gradient represents the faster reaction.

6.

A reaction produces 60 cm³ of gas in 30 seconds; what is the average rate of gas production?

60 ÷ 30 = 2.0 cm³ s⁻¹

7.6 Describe a catalyst as a substance that speeds up the rate of a reaction without altering the products of the reaction, being itself unchanged chemically and in mass at the end of the reaction

1.

What is a catalyst?

A substance that increases reaction rate without being chemically changed overall.

2.

How does a catalyst affect the rate of a chemical reaction?

It increases the rate.

3.

Does a catalyst change the products formed in a chemical reaction?

No.

4.

What happens to the chemical composition of a catalyst during a reaction?

It remains chemically unchanged overall.

5.

What happens to the mass of a catalyst at the end of a reaction?

Its mass remains unchanged overall.

6.

Why can a catalyst be described as being chemically unchanged at the end of a reaction?

It participates in the reaction pathway but is regenerated by the end.

7.7 Explain how the addition of a catalyst increases the rate of a reaction in terms of activation energy

1.

What is activation energy?

The minimum energy required for a successful reaction.

2.

How does a catalyst affect the activation energy of a chemical reaction?

It lowers the activation energy.

3.

Why does lowering the activation energy increase the rate of a chemical reaction?

More particles have enough energy to react successfully.

4.

How does a catalyst provide an alternative reaction pathway?

It provides a pathway with a lower activation energy.

5.

Why does a catalyst increase the number of successful collisions between reacting particles?

More collisions have enough energy to overcome the lower activation energy.

6.

How would the activation energy shown on a reaction profile change when a catalyst is added?

The peak is lower, so the activation energy is smaller.

7.8 Recall that enzymes are biological catalysts and that enzymes are used in the production of alcoholic drinks

1.

What are enzymes?

Biological catalysts.

2.

Why are enzymes described as biological catalysts?

They speed up biochemical reactions without being used up overall.

3.

How do enzymes affect the rate of biochemical reactions?

They increase reaction rates.

4.

How are enzymes used in the production of alcoholic drinks?

Yeast enzymes catalyse fermentation of sugars to produce ethanol.

5.

Which biological process involving enzymes produces ethanol during alcoholic fermentation?

Fermentation.

6.

Why are enzymes important in the production of alcoholic drinks?

They catalyse the reactions that convert carbohydrates into ethanol.

7.9 Recall that changes in heat energy accompany the following changes: a salts dissolving in water b neutralisation reactions c displacement reactions d precipitation reactions and that, when these reactions take place in solution, temperature changes can be measured to reflect the heat changes

1.

What happens to heat energy when a salt dissolves in water?

Depending on the salt, heat energy may be absorbed or released.

2.

Why can a temperature change be measured during an acid-alkali neutralisation reaction?

Energy is released or absorbed, causing a measurable temperature change.

3.

Why can a temperature change be measured during a displacement reaction carried out in solution?

A displacement reaction can release or absorb heat energy.

4.

Why can a temperature change be measured when a precipitate forms in a solution?

Formation of the precipitate can involve a heat energy change.

5.

How can a temperature change be used as evidence for a heat energy change during a reaction in solution?

A temperature increase indicates heat released; a temperature decrease indicates heat absorbed.

6.

Which four types of chemical change specified in GCSE Chemistry can involve measurable heat energy changes?

Dissolving salts, neutralisation, displacement and precipitation.

7.10 Describe an exothermic change or reaction as one in which heat energy is given out

1.

What is an exothermic reaction?

A reaction that gives out heat energy.

2.

What happens to the surroundings during an exothermic reaction?

They warm up.

3.

Why does the temperature of the surroundings usually increase during an exothermic reaction?

Heat is transferred to the surroundings.

4.

What happens to heat energy during an exothermic change?

It is released to the surroundings.

5.

Is the energy change associated with an exothermic reaction positive or negative if energy released to the surroundings is represented by a negative energy change?

Negative.

6.

Give one example of an exothermic chemical reaction.

Combustion or neutralisation.

7.11 Describe an endothermic change or reaction as one in which heat energy is taken in

1.

What is an endothermic reaction?

A reaction that takes in heat energy.

2.

What happens to the surroundings during an endothermic reaction?

They cool down.

3.

Why does the temperature of the surroundings usually decrease during an endothermic reaction?

Heat is transferred from the surroundings into the reaction.

4.

What happens to heat energy during an endothermic change?

It is absorbed from the surroundings.

5.

Is the energy change associated with an endothermic reaction positive or negative when energy absorbed by the reaction is represented as a positive energy change?

Positive.

6.

Give one example of an endothermic chemical reaction.

Thermal decomposition or some dissolving reactions.

7.12 Recall that the breaking of bonds is endothermic and the making of bonds is exothermic

1.

Why is energy required when a chemical bond is broken?

Energy is required to overcome the attractive forces holding the atoms together.

2.

Is bond breaking an endothermic or exothermic process?

Endothermic.

3.

Is bond making an endothermic or exothermic process?

Exothermic.

4.

Why is energy released when a chemical bond is formed?

The forming atoms move into a lower-energy, more stable arrangement.

5.

What happens to energy when bonds in reactants are broken?

Energy is absorbed.

6.

What happens to energy when new bonds in products are formed?

Energy is released.

7.13 Recall that the overall heat energy change for a reaction is: a exothermic if more heat energy is released in forming bonds in the products than is required in breaking bonds in the reactants b endothermic if less heat energy is released in forming bonds in the products than is required in breaking bonds in the reactants

1.

When is an overall chemical reaction exothermic in terms of the energy required to break bonds and the energy released when bonds form?

More energy is released making bonds than is required to break bonds.

2.

When is an overall chemical reaction endothermic in terms of the energy required to break bonds and the energy released when bonds form?

More energy is required to break bonds than is released making bonds.

3.

A reaction requires 500 kJ mol⁻¹ to break bonds and releases 650 kJ mol⁻¹ when new bonds form; is the reaction exothermic or endothermic?

Exothermic.

4.

A reaction requires 750 kJ mol⁻¹ to break bonds and releases 600 kJ mol⁻¹ when new bonds form; is the reaction exothermic or endothermic?

Endothermic.

5.

Why is a reaction exothermic when more energy is released making bonds than is required to break bonds?

More energy leaves the system than enters it.

6.

Why is a reaction endothermic when more energy is required to break bonds than is released when new bonds form?

More energy enters the system than leaves it during bond formation.

7.14 Calculate the energy change in a reaction given the energies of bonds (in kJ mol–1)

1.

How is the overall energy change of a reaction calculated using bond energies?

Energy change = energy required to break bonds − energy released forming bonds.

2.

Calculate the energy change for a reaction in which 500 kJ mol⁻¹ is required to break bonds and 700 kJ mol⁻¹ is released when bonds form.

500 − 700 = −200 kJ mol⁻¹

3.

Calculate the energy change for a reaction in which 900 kJ mol⁻¹ is required to break bonds and 650 kJ mol⁻¹ is released when bonds form.

900 − 650 = +250 kJ mol⁻¹

4.

The bond energies required to break all bonds in the reactants total 1200 kJ mol⁻¹, while the bond energies released when bonds form in the products total 1500 kJ mol⁻¹. Calculate the overall energy change.

1200 − 1500 = −300 kJ mol⁻¹

5.

The bond energies required to break all bonds in the reactants total 1800 kJ mol⁻¹, while the bond energies released when bonds form in the products total 1400 kJ mol⁻¹. Calculate the overall energy change.

1800 − 1400 = +400 kJ mol⁻¹

6.

A reaction has an energy change of −250 kJ mol⁻¹ and requires 1000 kJ mol⁻¹ to break all bonds in the reactants. How much energy is released when the bonds in the products form?

Energy released = 1250 kJ mol⁻¹

7.15 Explain the term activation energy

1.

What is activation energy?

The minimum energy required for a reaction to occur.

2.

Why must reacting particles have at least the activation energy for a successful collision to occur?

The particles need enough energy for bonds to break and a reaction to take place.

3.

What happens if colliding particles have less energy than the activation energy?

The collision is unsuccessful.

4.

Where is the activation energy shown on a reaction profile?

From the energy level of the reactants to the top of the reaction profile.

5.

How does activation energy affect the rate of a chemical reaction?

A higher activation energy generally means a slower reaction.

6.

Why does increasing the number of particles with energy greater than the activation energy increase the rate of reaction?

More particles can undergo successful collisions.

7.16 Draw and label reaction profiles for endothermic and exothermic reactions, identifying activation energy

1.

What features should be labelled on a reaction profile for an exothermic reaction?

Reactants, products, activation energy and energy change.

2.

What features should be labelled on a reaction profile for an endothermic reaction?

Reactants, products, activation energy and energy change.

3.

Where should the activation energy be shown on an exothermic reaction profile?

From the reactant energy level to the peak.

4.

Where should the activation energy be shown on an endothermic reaction profile?

From the reactant energy level to the peak.

5.

How should the relative energy levels of the reactants and products differ on an exothermic reaction profile?

Products are at a lower energy than reactants.

6.

How should the relative energy levels of the reactants and products differ on an endothermic reaction profile?

Products are at a higher energy than reactants.

8.1 Recall that hydrocarbons are compounds that contain carbon and hydrogen only

1.

What is a hydrocarbon?

A compound containing carbon and hydrogen only.

2.

Which two elements are present in every hydrocarbon?

Carbon and hydrogen.

3.

Is methane, CH₄, a hydrocarbon?

Yes.

4.

Is ethanol, C₂H₅OH, a hydrocarbon?

No; ethanol also contains oxygen.

5.

Is carbon dioxide, CO₂, a hydrocarbon?

No; carbon dioxide contains oxygen.

6.

How can the formula of a compound be used to determine whether it is a hydrocarbon?

Check whether the formula contains only carbon and hydrogen.

8.2 Describe crude oil as: a a complex mixture of hydrocarbons b containing molecules in which carbon atoms are in chains or rings (names, formulae and structures of specific ring molecules not required) c an important source of useful substances (fuels and feedstock for the petrochemical industry) d a finite resource

1.

What is crude oil?

A complex mixture of hydrocarbons.

2.

Why is crude oil described as a complex mixture of hydrocarbons?

It contains many different hydrocarbon molecules.

3.

What types of structures can the carbon atoms in crude oil molecules form?

Carbon atoms can form chains or rings.

4.

Why is crude oil an important source of useful substances?

It provides fuels and feedstock for the petrochemical industry.

5.

How is crude oil used as a feedstock for the petrochemical industry?

It provides raw materials for making useful chemicals and materials.

6.

Why is crude oil described as a finite resource?

It is finite and cannot be replaced quickly enough to match consumption.

8.3 Describe and explain the separation of crude oil into simpler, more useful mixtures by the process of fractional distillation

1.

What process is used to separate crude oil into simpler, more useful mixtures?

Fractional distillation.

2.

Why can the hydrocarbons in crude oil be separated by fractional distillation?

They have different boiling points.

3.

How does fractional distillation separate the different fractions of crude oil?

Crude oil is heated and the hydrocarbons vaporise and condense at different temperatures.

4.

What happens to hydrocarbons with lower boiling points during fractional distillation?

They rise higher in the column before condensing.

5.

What happens to hydrocarbons with higher boiling points during fractional distillation?

They condense lower down in the column.

6.

Why does fractional distillation produce fractions containing hydrocarbons with similar boiling points rather than individual pure hydrocarbons?

Each fraction contains hydrocarbons with a similar range of boiling points.

8.4 Recall the names and uses of the following fractions: a gases, used in domestic heating and cooking b petrol, used as fuel for cars c kerosene, used as fuel for aircraft d diesel oil, used as fuel for some cars and trains e fuel oil, used as fuel for large ships and in some power stations f bitumen, used to surface roads and roofs

1.

What is the main use of the gases fraction obtained from crude oil?

Domestic heating and cooking.

2.

What is petrol used for?

Fuel for cars.

3.

What is kerosene used for?

Fuel for aircraft.

4.

What is diesel oil used for?

Fuel for some cars and trains.

5.

What is fuel oil used for?

Fuel for large ships and some power stations.

6.

What is bitumen used for?

Surfacing roads and roofing.

8.5 Explain how hydrocarbons in different fractions differ from each other in: a the number of carbon and hydrogen atoms their molecules contain b boiling points c ease of ignition d viscosity and are mostly members of the alkane homologous series

1.

How does the number of carbon atoms in hydrocarbon molecules generally change between different crude oil fractions?

Generally increases from lighter to heavier fractions.

2.

How does boiling point generally change as the size of hydrocarbon molecules increases?

It increases.

3.

How does the ease of ignition generally change as hydrocarbon molecules become larger?

They become harder to ignite.

4.

How does viscosity generally change as hydrocarbon molecules become larger?

It increases.

5.

Which homologous series are most hydrocarbons in crude oil fractions members of?

Alkanes.

6.

Why do different crude oil fractions have different physical properties?

Their molecules have different sizes and therefore different intermolecular forces.

8.6 Explain an homologous series as a series of compounds which: a have the same general formula b differ by CH2 in molecular formulae from neighbouring compounds c show a gradual variation in physical properties, as exemplified by their boiling points d have similar chemical properties

1.

What is a homologous series?

A series of compounds with the same general formula, similar chemical properties and a gradual change in physical properties.

2.

What do members of the same homologous series have in common in terms of their general formula?

The same general formula.

3.

How do the molecular formulae of neighbouring members of a homologous series differ?

They differ by CH₂.

4.

How do physical properties change across a homologous series?

They change gradually.

5.

How do the chemical properties of members of the same homologous series compare?

They are similar.

6.

What molecular group is added between neighbouring members of a homologous series?

CH₂.

8.7 Describe the complete combustion of hydrocarbon fuels as a reaction in which: a carbon dioxide and water are produced b energy is given out

1.

What products are formed when a hydrocarbon undergoes complete combustion?

Carbon dioxide and water.

2.

What type of energy change occurs during the complete combustion of a hydrocarbon?

Energy is released.

3.

What two substances must be present for the complete combustion of a hydrocarbon fuel?

Hydrocarbon fuel and oxygen.

4.

Write the word equation for the complete combustion of a hydrocarbon.

Hydrocarbon + oxygen → carbon dioxide + water

5.

What happens to the carbon atoms in a hydrocarbon during complete combustion?

They are oxidised to carbon dioxide.

6.

What happens to the hydrogen atoms in a hydrocarbon during complete combustion?

They are oxidised to water.

8.8 Explain why the incomplete combustion of hydrocarbons can produce carbon and carbon monoxide

1.

Why can incomplete combustion of hydrocarbons occur?

There is insufficient oxygen.

2.

What solid product can be formed when a hydrocarbon undergoes incomplete combustion?

Carbon (soot).

3.

What poisonous gas can be produced during incomplete combustion of hydrocarbons?

Carbon monoxide, CO.

4.

Why does incomplete combustion produce carbon monoxide instead of only carbon dioxide?

There is not enough oxygen for all carbon to be fully oxidised to CO₂.

5.

What happens to the amount of oxygen available when incomplete combustion occurs?

It is limited.

6.

What are the two carbon-containing products that can be formed during incomplete combustion of hydrocarbons?

Carbon particles and carbon monoxide.

8.9 Explain how carbon monoxide behaves as a toxic gas

1.

Why is carbon monoxide a toxic gas?

It prevents blood from transporting enough oxygen.

2.

How does carbon monoxide affect the blood?

It binds strongly to haemoglobin.

3.

Which molecule in red blood cells does carbon monoxide bind to?

Haemoglobin.

4.

Why does carbon monoxide reduce the ability of the blood to transport oxygen?

It reduces the amount of oxygen haemoglobin can carry.

5.

Why can carbon monoxide poisoning prevent body cells from receiving enough oxygen?

Cells receive less oxygen for aerobic respiration.

6.

Why is carbon monoxide particularly dangerous because it is difficult to detect without a detector?

It is colourless and odourless.

8.10 Describe the problems caused by incomplete combustion producing carbon monoxide and soot in appliances that use carbon compounds as fuels

1.

What health problem can carbon monoxide produced by incomplete combustion cause?

Carbon monoxide poisoning.

2.

Why is carbon monoxide from faulty fuel-burning appliances dangerous?

Carbon monoxide reduces the blood's ability to transport oxygen.

3.

What is soot?

Soot is fine particles of carbon.

4.

How can soot produced by incomplete combustion affect human health?

It can damage the respiratory system and contribute to breathing problems.

5.

Why can incomplete combustion be particularly dangerous in poorly ventilated rooms?

There may be less oxygen available, increasing incomplete combustion.

6.

What two harmful products of incomplete combustion are particularly important in fuel-burning appliances?

Carbon monoxide and soot.

8.11 Explain how impurities in some hydrocarbon fuels result in the production of sulfur dioxide

1.

Why can burning some hydrocarbon fuels produce sulfur dioxide?

The fuels contain sulfur impurities.

2.

Where does the sulfur that produces sulfur dioxide come from?

From sulfur-containing compounds in the fuel.

3.

What happens to sulfur-containing impurities when a hydrocarbon fuel is burned?

They are oxidised during combustion.

4.

What pollutant is produced when sulfur impurities in fuels are oxidised during combustion?

Sulfur dioxide, SO₂.

5.

Why do sulfur-containing impurities in fuels contribute to air pollution?

It contributes to air pollution and acid rain.

6.

What must happen to sulfur in a fuel for sulfur dioxide to be produced during combustion?

It must be oxidised by oxygen during combustion.

8.12 Explain some problems associated with acid rain caused when sulfur dioxide dissolves in rain water

1.

How does sulfur dioxide contribute to acid rain?

Sulfur dioxide dissolves in rainwater and forms acidic solutions.

2.

What happens when sulfur dioxide dissolves in rainwater?

Acidic substances are formed.

3.

How can acid rain damage aquatic ecosystems?

It can lower the pH of lakes and rivers, harming aquatic organisms.

4.

How can acid rain damage buildings and statues made from limestone?

It reacts with limestone and damages buildings and statues.

5.

How can acid rain affect plants and forests?

It can damage leaves, soil and forests.

6.

Why is sulfur dioxide from burning fuels an environmental problem?

It causes acid rain and environmental damage.

8.13 Explain why, when fuels are burned in engines, oxygen and nitrogen can react together at high temperatures to produce oxides of nitrogen, which are pollutants

1.

Why can nitrogen and oxygen react together inside an engine?

High temperatures provide enough energy for nitrogen and oxygen molecules to react.

2.

What conditions inside an engine allow nitrogen and oxygen to react?

Very high temperatures, such as those inside engines.

3.

What products are formed when nitrogen and oxygen react at high temperatures in engines?

Oxides of nitrogen, NOₓ.

4.

Why are oxides of nitrogen considered pollutants?

They contribute to air pollution and can cause acid rain and photochemical pollution.

5.

Why does the high temperature inside an engine increase the likelihood of nitrogen and oxygen reacting?

More particles have enough energy to overcome the activation energy.

6.

Which two gases from the air react together at high temperatures in engines to form oxides of nitrogen?

Nitrogen and oxygen.

8.14 Evaluate the advantages and disadvantages of using hydrogen, rather than petrol, as a fuel in cars

1.

What is one environmental advantage of using hydrogen rather than petrol as a fuel in cars?

Hydrogen produces no carbon dioxide at the point of use in a fuel cell.

2.

Why does using hydrogen in a fuel cell avoid producing carbon dioxide directly at the point of use?

The fuel cell's overall product is water rather than carbon dioxide.

3.

What is one disadvantage of using hydrogen as a fuel compared with petrol?

Hydrogen is difficult to store and transport and requires suitable infrastructure.

4.

Why can the method used to produce hydrogen affect its overall environmental impact?

Hydrogen made from fossil fuels can produce carbon dioxide.

5.

What storage problem is associated with using hydrogen as a fuel in cars?

It must be stored at high pressure or low temperature because of its low density.

6.

What factors should be considered when evaluating hydrogen against petrol as a car fuel?

Emissions, production method, storage, cost, infrastructure, availability, efficiency and safety.

8.15 Recall that petrol, kerosene and diesel oil are non-renewable fossil fuels obtained from crude oil and methane is a nonrenewable fossil fuel found in natural gas

1.

Why are petrol, kerosene and diesel oil classified as non-renewable fossil fuels?

They are obtained from finite crude oil reserves.

2.

What raw material is petrol obtained from?

Crude oil.

3.

What raw material is kerosene obtained from?

Crude oil.

4.

What raw material is diesel oil obtained from?

Crude oil.

5.

Where is methane found as a non-renewable fossil fuel?

Natural gas.

6.

Why are fossil fuels described as non-renewable?

They are used much faster than they can be naturally replaced.

8.16 Explain how cracking involves the breaking down of larger, saturated hydrocarbon molecules (alkanes) into smaller, more useful ones, some of which are unsaturated (alkenes)

1.

What is cracking?

Breaking large hydrocarbon molecules into smaller, more useful molecules.

2.

What type of hydrocarbon molecules are broken down during cracking?

Large, saturated hydrocarbons called alkanes.

3.

What happens to large alkane molecules during cracking?

They are broken into smaller hydrocarbon molecules.

4.

What types of smaller hydrocarbons can be produced by cracking?

Smaller alkanes and alkenes.

5.

Why are some of the products of cracking described as unsaturated hydrocarbons?

They contain carbon-carbon double bonds.

6.

What is the difference between the saturated reactants and unsaturated products that can be formed during cracking?

Alkanes have only single C–C bonds; alkenes contain a C=C bond.

8.17 Explain why cracking is necessary

1.

Why is cracking necessary in the petroleum industry?

It produces more useful smaller hydrocarbons from larger ones.

2.

Why is there a greater demand for some smaller hydrocarbons than the larger hydrocarbons found in crude oil?

Smaller hydrocarbons are in greater demand.

3.

How does cracking help meet the demand for petrol?

It produces more hydrocarbons suitable for petrol and other fuels.

4.

Why are alkenes produced by cracking useful to the chemical industry?

They are used to make polymers and other chemicals.

5.

How does cracking make crude oil more useful to society?

It matches supply more closely to demand.

6.

Why would simply separating crude oil into fractions not be enough to meet the demand for all petroleum products?

Crude oil contains too much of some heavy fractions and too little of some more useful smaller hydrocarbons.

8.18 Recall that the gases produced by volcanic activity formed the Earth's early atmosphere

1.

What formed the gases that made up the Earth's early atmosphere?

Volcanic activity.

2.

How did volcanic activity contribute to the formation of the Earth's early atmosphere?

Volcanoes released gases into the atmosphere.

3.

What happened to gases released by volcanic activity to form the early atmosphere?

The gases accumulated around the Earth.

4.

Why was volcanic activity important in the formation of the Earth's early atmosphere?

It released large amounts of gases from Earth's interior.

5.

Which major gases were released by volcanic activity that contributed to the Earth's early atmosphere?

Mainly carbon dioxide and water vapour, with smaller amounts of other gases.

6.

How did the Earth's early atmosphere differ from the atmosphere present today?

It contained little or no oxygen and much more carbon dioxide and water vapour.

8.19 Describe that the Earth's early atmosphere was thought to contain: a little or no oxygen b a large amount of carbon dioxide c water vapour d small amounts of other gases and interpret evidence relating to this

1.

What was the estimated oxygen content of the Earth's early atmosphere?

Little or no oxygen.

2.

Which gas was present in large amounts in the Earth's early atmosphere?

Carbon dioxide.

3.

Why was water vapour present in the Earth's early atmosphere?

Water vapour was released by volcanic activity.

4.

What other gases were thought to be present in small amounts in the Earth's early atmosphere?

Small amounts of other gases.

5.

How does evidence from the composition of ancient rocks help scientists understand the Earth's early atmosphere?

Ancient rocks can provide evidence of the atmospheric conditions when they formed.

6.

What evidence can be used to interpret the likely composition of the Earth's early atmosphere?

Evidence from ancient rocks, fossils and geological records.

8.20 Explain how condensation of water vapour formed oceans

1.

How did water vapour in the early atmosphere lead to the formation of oceans?

As the Earth cooled, water vapour condensed into liquid water.

2.

What process changed water vapour into liquid water as the Earth cooled?

Condensation.

3.

Why did cooling of the Earth cause water vapour to condense?

Cooling lowers the energy of water molecules, allowing liquid water to form.

4.

How did condensation contribute to the formation of the first oceans?

Condensed water collected on the Earth's surface and formed oceans.

5.

What happened to atmospheric water vapour as the Earth's surface became cooler?

It condensed into liquid water.

6.

Why could liquid water accumulate on the Earth's surface after condensation occurred?

Gravity allowed liquid water to collect on the surface.

8.21 Explain how the amount of carbon dioxide in the atmosphere was decreased when carbon dioxide dissolved as the oceans formed

1.

How did the formation of the oceans decrease the amount of carbon dioxide in the atmosphere?

Carbon dioxide dissolved in the newly formed oceans.

2.

Why did atmospheric carbon dioxide dissolve into the oceans?

Carbon dioxide is soluble in water.

3.

What happened to atmospheric carbon dioxide when liquid oceans formed?

It moved from the atmosphere into the oceans.

4.

How did the formation of oceans contribute to the reduction of atmospheric carbon dioxide?

More carbon dioxide was removed from the atmosphere as more ocean water formed.

5.

Why was carbon dioxide able to dissolve in the oceans as they formed?

The newly formed liquid water dissolved carbon dioxide.

6.

What happened to the concentration of carbon dioxide in the atmosphere as increasing amounts dissolved in the oceans?

The atmospheric concentration decreased.

8.22 Explain how the growth of primitive plants used carbon dioxide and released oxygen by photosynthesis and consequently the amount of oxygen in the atmosphere gradually increased

1.

How did primitive plants affect the amount of carbon dioxide in the atmosphere?

They removed carbon dioxide from the atmosphere.

2.

How did photosynthesis by primitive plants affect the amount of oxygen in the atmosphere?

They released oxygen.

3.

Which gas did primitive plants remove from the atmosphere during photosynthesis?

Carbon dioxide.

4.

Which gas did primitive plants release during photosynthesis?

Oxygen.

5.

Why did the amount of oxygen in the atmosphere gradually increase as primitive plants became more widespread?

More photosynthesis occurred, releasing more oxygen.

6.

How did photosynthesis by primitive plants contribute to the change from an oxygen-poor early atmosphere to an oxygen-rich atmosphere?

Photosynthesis gradually removed CO₂ and added O₂ to the atmosphere.

8.23 Describe the chemical test for oxygen

1.

How is oxygen tested for in a laboratory?

Insert a glowing splint into the gas.

2.

What happens to a glowing splint when it is placed in oxygen?

The splint relights.

3.

What positive result indicates the presence of oxygen using a glowing splint?

A glowing splint relights.

4.

Why does a glowing splint relight when placed in oxygen?

Oxygen supports combustion.

5.

What apparatus or test is used to distinguish oxygen from gases that do not support combustion?

The glowing-splint test.

6.

A gas relights a glowing splint; which gas is indicated by this test?

Oxygen.

8.24 Describe how various gases in the atmosphere, including carbon dioxide, methane and water vapour, absorb heat radiated from the Earth, subsequently releasing energy which keeps the Earth warm: this is known as the greenhouse effect

1.

What is the greenhouse effect?

The warming effect caused when atmospheric gases absorb and re-emit infrared radiation.

2.

Which type of radiation is emitted by the Earth and absorbed by greenhouse gases?

Infrared radiation.

3.

Which atmospheric gases specified in the GCSE Chemistry specification contribute to the greenhouse effect?

Carbon dioxide, methane and water vapour.

4.

How do carbon dioxide, methane and water vapour help keep the Earth warm?

They absorb infrared radiation emitted by the Earth and re-release energy in all directions.

5.

What happens after greenhouse gases absorb infrared radiation emitted by the Earth?

They emit infrared radiation again.

6.

Why would the Earth be colder without the natural greenhouse effect?

Less heat energy would be retained in the atmosphere.

8.25 Evaluate the evidence for human activity causing climate change, considering: a the correlation between the change in atmospheric carbon dioxide concentration, the consumption of fossil fuels and temperature change b the uncertainties caused by the location where these measurements are taken and historical accuracy

1.

What correlation is observed between atmospheric carbon dioxide concentration and global temperature change?

Atmospheric CO₂ concentration and global temperature have increased over time.

2.

How is the consumption of fossil fuels related to changes in atmospheric carbon dioxide concentration?

Burning fossil fuels increases atmospheric CO₂.

3.

Why does a correlation between fossil fuel consumption, atmospheric carbon dioxide concentration and temperature provide evidence for human-caused climate change?

The three quantities show correlated long-term changes consistent with human activity contributing to warming.

4.

Why does correlation between atmospheric carbon dioxide concentration and temperature not, by itself, prove causation?

Correlation alone does not establish cause and effect.

5.

How can the locations where atmospheric carbon dioxide and temperature measurements are taken introduce uncertainty into climate data?

Different locations can have different local conditions and measurements may not represent the whole planet equally.

6.

Why can the historical accuracy of measurements introduce uncertainty when evaluating evidence for human-caused climate change?

Older measurements may be less accurate or less complete.

8.26 Describe: a the composition of today's atmosphere b the potential effects on the climate of increased levels of carbon dioxide and methane generated by human activity, including burning fossil fuels and livestock farming c that these effects may be mitigated: consider scale, risk and environmental implications

1.

What are the main gases that make up today's atmosphere?

Mainly nitrogen and oxygen, with smaller amounts of argon, carbon dioxide and other gases.

2.

How can burning fossil fuels increase the concentration of carbon dioxide in the atmosphere?

Burning fossil fuels transfers carbon from long-term stores into the atmosphere as CO₂.

3.

How can livestock farming increase the concentration of methane in the atmosphere?

Livestock farming produces methane, particularly from digestion and waste.

4.

How can increased atmospheric concentrations of carbon dioxide and methane affect the climate?

They increase the greenhouse effect and can contribute to global warming and climate change.

5.

What methods could be used to mitigate climate change caused by increased carbon dioxide and methane levels?

Renewable energy, energy efficiency, carbon capture, reforestation and changes in agriculture and transport.

6.

Why should the scale, risks and environmental implications of proposed climate-change mitigation methods be considered?

Different methods have different costs, scales, risks and environmental impacts.

9.1C Explain why the test for any ion must be unique

1.

Why must a chemical test for an ion produce a unique result?

It must distinguish that ion from other possible ions.

2.

What is meant by a unique test for an ion?

A test that produces a characteristic result for one specific ion.

3.

Why would a test that gives the same positive result for several different ions be unsuitable for identifying a specific ion?

It could not reliably identify the specific ion.

4.

How does a unique test allow an unknown ion to be identified?

The characteristic result can be matched to the known behaviour of the ion.

5.

Why is it important to distinguish between different ions when analysing an unknown substance?

Different ions have different chemical properties.

6.

What would be the problem if two different ions produced identical results in the same chemical test?

The identity of the ion would be ambiguous.

9.2C Describe flame tests to identify the following ions in solids: a lithium ion, Li+ (red) b sodium ion, Na+ (yellow) c potassium ion, K+ (lilac) d calcium ion, Ca2+ (orange-red) e copper ion, Cu2+ (blue-green)

1.

What flame-test colour is produced by lithium ions, Li⁺?

Red.

2.

What flame-test colour is produced by sodium ions, Na⁺?

Yellow.

3.

What flame-test colour is produced by potassium ions, K⁺?

Lilac.

4.

What flame-test colour is produced by calcium ions, Ca²⁺?

Orange-red.

5.

What flame-test colour is produced by copper(II) ions, Cu²⁺?

Blue-green.

6.

How can a flame test be used to identify a metal ion present in a solid compound?

Heat the solid in a flame and compare the observed flame colour with known colours.

9.3C Describe tests to identify the following ions in solids or solutions as appropriate: a aluminium ion, Al3+ b calcium ion, Ca2+ c copper ion, Cu2+ d iron(II) ion, Fe2+ e iron(III) ion, Fe3+ f ammonium ion, NH4 + using sodium hydroxide solution

1.

How can sodium hydroxide solution be used to test for aluminium ions, Al³⁺?

Add sodium hydroxide solution: a white precipitate forms, which dissolves in excess sodium hydroxide.

2.

How can sodium hydroxide solution be used to distinguish calcium ions, Ca²⁺, from other specified metal ions?

A white precipitate is formed.

3.

What observation is produced when sodium hydroxide solution is added to a solution containing copper(II) ions, Cu²⁺?

A blue precipitate.

4.

What observation is produced when sodium hydroxide solution is added to a solution containing iron(II) ions, Fe²⁺?

A green precipitate.

5.

What observation is produced when sodium hydroxide solution is added to a solution containing iron(III) ions, Fe³⁺?

A red-brown precipitate.

6.

How can sodium hydroxide solution be used to test for ammonium ions, NH₄⁺?

Add sodium hydroxide and warm gently; ammonia gas is released and turns damp red litmus paper blue.

9.4C Describe the chemical test for ammonia

1.

How can you test a gas to determine whether it is ammonia?

Expose damp red litmus paper to the gas.

2.

What happens when ammonia gas comes into contact with damp red litmus paper?

It turns blue.

3.

What colour does damp red litmus paper turn when exposed to ammonia?

Blue.

4.

Why must litmus paper be damp when testing for ammonia?

Ammonia must dissolve in water to produce alkaline conditions.

5.

What does the colour change of damp red litmus paper show when testing for ammonia?

It indicates an alkaline gas, consistent with ammonia.

6.

What observation would confirm that an unknown gas is ammonia?

Damp red litmus paper turns blue.

9.5C Describe tests to identify the following ions in solids or solutions as appropriate: a carbonate ion, CO3 2–, using dilute acid and identifying the carbon dioxide evolved b sulfate ion, SO4 2–, using dilute hydrochloric acid and barium chloride solution c chloride ion, Cl–, bromide ion, Br–, iodide ion, I–, using dilute nitric acid and silver nitrate solution

1.

How can carbonate ions, CO₃²⁻, be identified using dilute acid?

Add dilute acid; effervescence occurs as CO₂ is produced.

2.

How can the gas produced from a carbonate ion test be confirmed as carbon dioxide?

Bubble the gas through limewater; it turns cloudy.

3.

How can sulfate ions, SO₄²⁻, be tested for using dilute hydrochloric acid and barium chloride solution?

Add dilute hydrochloric acid, then barium chloride solution; a white precipitate indicates sulfate.

4.

What observation indicates the presence of sulfate ions when barium chloride solution is added?

A white precipitate of barium sulfate.

5.

How can chloride, bromide and iodide ions be tested for using dilute nitric acid and silver nitrate solution?

Acidify with dilute nitric acid, then add silver nitrate solution.

6.

How can the results of the silver nitrate test be used to distinguish between chloride, bromide and iodide ions?

Chloride → white precipitate; bromide → cream precipitate; iodide → yellow precipitate.

9.6C Core Practical: Identify the ions in unknown salts, using the tests for the specified cations and anions in 9.2C, 9.3C, 9.4C, 9.5C

1.

How can you systematically identify the cation and anion present in an unknown salt?

Test separate portions for the cation and anion using the appropriate tests, then combine the results.

2.

Which flame tests can be used to identify lithium, sodium, potassium, calcium and copper(II) ions in an unknown solid?

Lithium → red; sodium → yellow; potassium → lilac; calcium → orange-red; copper(II) → blue-green.

3.

Which chemical tests can be used to identify aluminium, calcium, copper(II), iron(II), iron(III) and ammonium ions?

Add sodium hydroxide solution and observe precipitate colours or ammonia release for ammonium ions.

4.

Which chemical tests can be used to identify carbonate, sulfate, chloride, bromide and iodide ions?

Carbonate → acid and CO₂ test; sulfate → hydrochloric acid and barium chloride; halides → nitric acid and silver nitrate.

5.

Why should separate portions of an unknown salt be used for different ion tests?

To avoid one test reagent contaminating or interfering with another test.

6.

How can a combination of positive test results be used to identify an unknown salt?

Match the full set of positive results to the known ions.

9.7C Identify the ions in unknown salts, using results of the tests above

1.

How can the results of chemical tests be used to identify the ions present in an unknown salt?

Compare each test result with the known characteristic test results.

2.

An unknown salt gives a lilac flame and produces a cream precipitate with silver nitrate. Which ions are present?

K⁺ and Br⁻.

3.

An unknown salt gives a blue-green flame and produces a white precipitate with barium chloride after acidification. Which ions are present?

Cu²⁺ and SO₄²⁻.

4.

An unknown salt produces a brick-red flame and a white precipitate with silver nitrate after adding dilute nitric acid. Which ions are present?

Ca²⁺ and Cl⁻.

5.

An unknown salt produces a red-brown precipitate with sodium hydroxide solution. Which metal ion is present?

Fe³⁺.

6.

Why should more than one test result be considered when identifying an unknown salt?

One result may not uniquely identify the salt, so multiple results confirm the identification.

9.8C Describe that instrumental methods of analysis are available and that these may improve sensitivity, accuracy and speed of tests

1.

What are instrumental methods of chemical analysis?

Methods using instruments to detect or measure substances.

2.

Why can instrumental methods be more sensitive than simple chemical tests?

They can detect very small concentrations.

3.

How can instrumental methods improve the accuracy of chemical analysis?

They provide precise numerical measurements.

4.

How can instrumental methods make chemical analysis faster?

They can automate measurements and analyse samples rapidly.

5.

What is meant by the sensitivity of an analytical method?

The ability to detect small quantities of a substance.

6.

Why might an instrumental method be preferred over a traditional chemical test?

They may provide greater sensitivity, accuracy and speed.

9.9C Evaluate data from a flame photometer: a to determine the concentration of ions in dilute solution using a calibration curve b to identify metal ions by comparing the data with reference data (no knowledge of the instrument or how it works is required)

1.

How can a calibration curve from a flame photometer be used to determine the concentration of an ion in a dilute solution?

Read the unknown's signal on the calibration curve and determine the corresponding concentration.

2.

What information is plotted on a calibration curve used with a flame photometer?

Ion concentration against measured signal intensity.

3.

How can the concentration of an unknown solution be determined from a flame-photometer calibration curve?

Locate the unknown reading on the graph and read across to the concentration axis.

4.

How can flame-photometer data be used to identify an unknown metal ion?

Compare the measured data with reference values for known ions.

5.

How can the data for an unknown metal ion be compared with reference data to identify the ion?

Match the unknown value with the closest reference value.

6.

Why is a calibration curve needed when determining an unknown ion concentration using flame-photometer data?

It provides the relationship needed to convert instrument readings into concentration.

9.10C Recall the formulae of molecules of the alkanes, methane, ethane, propane and butane, and draw the structures of these molecules, showing all covalent bonds

1.

What is the molecular formula of methane?

CH₄

2.

What is the molecular formula of ethane?

C₂H₆

3.

What is the molecular formula of propane?

C₃H₈

4.

What is the molecular formula of butane?

C₄H₁₀

5.

What is the general formula for the alkane homologous series?

CnH2n+2

6.

How should the structures of methane, ethane, propane and butane be represented to show all covalent bonds?

Methane: CH₄; ethane: CH₃–CH₃; propane: CH₃–CH₂–CH₃; butane: CH₃–CH₂–CH₂–CH₃, showing all covalent bonds.

9.11C Explain why the alkanes are saturated hydrocarbons

1.

What is meant by the term saturated hydrocarbon?

A hydrocarbon containing only single bonds between carbon atoms and the maximum possible number of hydrogen atoms.

2.

Why are alkanes described as saturated hydrocarbons?

They contain only single C–C bonds and have the maximum number of hydrogen atoms.

3.

What type of carbon-carbon bond is present in alkanes?

Single covalent C–C bonds.

4.

Why do alkanes contain the maximum possible number of hydrogen atoms?

There are no multiple C–C bonds, so each carbon has the maximum number of hydrogen atoms.

5.

What feature of an alkane molecule distinguishes it from an unsaturated hydrocarbon?

It contains only single C–C bonds.

6.

Why can alkanes not contain a carbon-carbon double bond?

A C=C bond would mean fewer hydrogen atoms and therefore an unsaturated hydrocarbon.

9.12C Recall the formulae of molecules of the alkenes, ethene, propene, butene, and draw the structures of these molecules, showing all covalent bonds (but-1-ene and but-2-ene only)

1.

What is the molecular formula of ethene?

C₂H₄

2.

What is the molecular formula of propene?

C₃H₆

3.

What is the molecular formula of butene?

C₄H₈

4.

What is the molecular formula of but-1-ene?

C₄H₈

5.

What is the molecular formula of but-2-ene?

C₄H₈

6.

How should the structures of ethene, propene, but-1-ene and but-2-ene be represented to show all covalent bonds?

Ethene: CH₂=CH₂; propene: CH₂=CH–CH₃; but-1-ene: CH₂=CH–CH₂–CH₃; but-2-ene: CH₃–CH=CH–CH₃.

9.13C Explain why the alkenes are unsaturated hydrocarbons, describing that their molecules contain the functional group C=C

1.

What is meant by an unsaturated hydrocarbon?

A hydrocarbon containing one or more carbon-carbon double or triple bonds.

2.

Why are alkenes described as unsaturated hydrocarbons?

They contain a carbon-carbon double bond and therefore do not contain the maximum number of hydrogen atoms.

3.

What functional group is present in all alkenes?

C=C

4.

What type of bond is represented by C=C?

A double covalent bond.

5.

Why does the C=C bond make an alkene unsaturated?

It means the molecule does not have the maximum possible number of hydrogen atoms.

6.

How can an alkene be distinguished structurally from an alkane?

Alkenes contain C=C; alkanes contain only single C–C bonds.

9.14C Recall the addition reaction of ethene with bromine, showing the structures of reactants and products, and extend this to other alkenes

1.

What type of reaction occurs when ethene reacts with bromine?

Addition reaction.

2.

What are the products when ethene reacts with bromine?

1,2-dibromoethane.

3.

What happens to the C=C bond when ethene reacts with bromine?

The C=C double bond becomes a C–C single bond.

4.

What is the structural formula of the product formed when ethene reacts with bromine?

CH₂Br–CH₂Br

5.

What is meant by an addition reaction of an alkene?

A reaction in which atoms are added across a carbon-carbon double bond.

6.

How can the reaction of ethene with bromine be extended to other alkenes?

Bromine adds across the C=C bond of other alkenes in the same way.

9.15C Explain how bromine water is used to distinguish between alkanes and alkenes

1.

How can bromine water be used to distinguish between an alkane and an alkene?

Add bromine water to the hydrocarbon.

2.

What happens to the colour of bromine water when it is added to an alkene?

It decolourises.

3.

What happens to the colour of bromine water when it is added to an alkane under normal test conditions?

It remains orange.

4.

Why does an alkene decolourise bromine water?

Bromine reacts by addition across the C=C bond.

5.

What does a positive bromine-water test indicate about a hydrocarbon?

The hydrocarbon contains a C=C double bond.

6.

Which structural feature of an alkene allows it to react with bromine water?

The C=C bond.

9.16C Describe how the complete combustion of alkanes and alkenes involves the oxidation of the hydrocarbons to produce carbon dioxide and water

1.

What are the products of the complete combustion of an alkane?

Carbon dioxide and water.

2.

What are the products of the complete combustion of an alkene?

Carbon dioxide and water.

3.

Why is the complete combustion of hydrocarbons described as an oxidation reaction?

Oxygen is added to the carbon and hydrogen during oxidation.

4.

What happens to the carbon atoms in a hydrocarbon during complete combustion?

They are oxidised to carbon dioxide.

5.

What happens to the hydrogen atoms in a hydrocarbon during complete combustion?

They are oxidised to water.

6.

What two products are formed when a hydrocarbon undergoes complete combustion in excess oxygen?

Carbon dioxide and water.

9.17C Recall that a polymer is a substance of high average relative molecular mass made up of small repeating units

1.

What is a polymer?

A substance of high average relative molecular mass made from repeating units.

2.

What does it mean for a polymer to have a high average relative molecular mass?

It contains a very large number of atoms joined in long chains.

3.

What are the small repeating units that make up a polymer called?

Monomers.

4.

Why is a polymer described as being made up of repeating units?

The same structural unit is repeated many times.

5.

How does the relative molecular mass of a polymer compare with that of its repeating units?

It is much higher.

6.

What structural feature distinguishes a polymer from a small molecule?

A polymer contains many repeating units joined into a long chain.

9.18C Describe: a how ethene molecules can combine together in a polymerisation reaction b that the addition polymer formed is called poly(ethene) (conditions and mechanisms not required)

1.

What type of reaction occurs when ethene molecules combine to form poly(ethene)?

Addition polymerisation.

2.

What polymer is formed when ethene molecules undergo addition polymerisation?

Poly(ethene).

3.

What happens to the C=C bond in an ethene molecule during addition polymerisation?

The C=C opens and forms single C–C bonds to neighbouring monomers.

4.

What is the repeating unit in poly(ethene)?

–CH₂–CH₂–

5.

What is the monomer used to produce poly(ethene)?

Ethene.

6.

Why is poly(ethene) described as an addition polymer?

No small molecule is eliminated; the monomers add together.

9.19C Describe how other addition polymers can be made by combining together other monomer molecules containing C=C, to include poly(propene), poly(chloroethene) (PVC) and poly(tetrafluoroethene) (PTFE) (conditions and mechanisms not required)

1.

What type of monomer is required to make an addition polymer?

A monomer containing a C=C bond.

2.

Which monomer is used to make poly(propene)?

Propene.

3.

Which monomer is used to make poly(chloroethene), PVC?

Chloroethene.

4.

Which monomer is used to make poly(tetrafluoroethene), PTFE?

Tetrafluoroethene.

5.

What happens to the C=C bond when a monomer containing C=C forms an addition polymer?

The double bond opens to form single bonds in the polymer chain.

6.

Why can propene, chloroethene and tetrafluoroethene form addition polymers?

Each contains a C=C double bond that can undergo addition polymerisation.

9.20C Deduce the structure of a monomer from the structure of an addition polymer and vice versa

1.

How can the monomer of an addition polymer be identified from the polymer's repeating unit?

Identify the repeating unit and restore the C=C bond.

2.

What bond in an addition polymer indicates where the original C=C bond of the monomer has reacted?

The C–C single bond in the polymer backbone between the appropriate atoms.

3.

How can the C=C bond be reconstructed when deducing a monomer from an addition polymer?

Replace the relevant C–C single bond with C=C and add the appropriate hydrogen or substituent arrangement.

4.

How can the repeating unit of an addition polymer be deduced from its monomer?

Open the C=C bond and join the monomers into a chain.

5.

What structural change occurs when a C=C-containing monomer becomes part of an addition polymer?

The C=C bond becomes C–C bonds linking neighbouring monomer units.

6.

How can you determine the monomer used to make a polymer from a displayed structure of the polymer?

Find the repeating unit, then restore the C=C bond to obtain the monomer.

9.21C Explain how the uses of polymers are related to their properties and vice versa: including poly(ethene), poly(propene), poly(chloroethene) (PVC) and poly(tetrafluoroethene) (PTFE)

1.

How are the uses of polymers determined by their physical properties?

Properties such as strength, flexibility, chemical resistance and electrical insulation determine suitable uses.

2.

Why is poly(ethene) suitable for making plastic bags and containers?

It is lightweight, flexible and waterproof.

3.

Why is poly(propene) suitable for making items such as ropes and packaging?

It is strong and durable.

4.

Why is PVC suitable for making products such as pipes and cable insulation?

It is durable, chemically resistant and a good electrical insulator.

5.

Why is PTFE suitable for non-stick coatings?

It is chemically unreactive and has a low-friction surface.

6.

How can the use of a polymer give information about the properties it must have?

The required properties of the use indicate which polymer is suitable.

9.22C Explain: a why polyesters are condensation polymers b how a polyester is formed when a monomer molecule containing two carboxylic acid groups is reacted with a monomer molecule containing two alcohol groups c how a molecule of water is formed each time an ester link is formed

1.

Why are polyesters classified as condensation polymers?

Small molecules, including water, are eliminated during polymer formation.

2.

What two types of functional groups react to form a polyester?

Carboxylic acid groups and alcohol groups.

3.

What type of monomer contains two carboxylic acid groups in polyester formation?

A di-carboxylic acid monomer.

4.

What type of monomer contains two alcohol groups in polyester formation?

A diol monomer.

5.

What is formed when a carboxylic acid group reacts with an alcohol group during polyester formation?

An ester link.

6.

Why is a molecule of water formed each time an ester link is formed?

An –OH group and an –H from the reacting groups combine to form H₂O.

9.23C Describe some problems associated with polymers including the: a availability of starting materials b persistence in landfill sites, due to non-biodegradability c gases produced during disposal by combustion d requirement to sort polymers so that they can be melted and reformed into a new product

1.

Why can the availability of starting materials be a problem when producing polymers?

Many polymers depend on finite fossil-fuel feedstocks.

2.

Why do many polymers persist in landfill sites?

They are often non-biodegradable.

3.

Why is the non-biodegradability of many polymers an environmental problem?

They remain in landfill for long periods.

4.

What problem can be caused by gases released when polymers are disposed of by combustion?

Toxic or harmful gases may be released.

5.

Why must polymers be sorted before they can be melted and reformed into new products?

Different polymers have different properties and must be separated before processing.

6.

Why can the disposal of polymers create environmental problems?

They can consume finite resources, persist in landfill and cause pollution during disposal.

9.24C Evaluate the advantages and disadvantages of recycling polymers, including economic implications, availability of starting materials and environmental impact

1.

What are the environmental advantages of recycling polymers?

Less polymer waste enters landfill and less new material is needed.

2.

How can recycling polymers reduce the need for new starting materials?

Recycled polymers replace some new raw materials.

3.

How can recycling polymers affect the economic cost of producing plastic products?

It can reduce raw-material costs but collection, sorting and processing can be expensive.

4.

What are the disadvantages or limitations of recycling polymers?

Sorting, contamination and repeated recycling can limit their usefulness.

5.

Why does sorting polymers affect the practicality of recycling?

Mixed polymers may not process together successfully.

6.

How should economic, environmental and resource factors be considered when evaluating polymer recycling?

Consider environmental benefits, processing costs, starting-material availability and the quality of recycled products.

9.25C Recall that: a DNA is a polymer made from four different monomers called nucleotides (names of nucleotides not required) b starch is a polymer based on sugars c proteins are polymers based on amino acids

1.

What type of polymer is DNA?

DNA is a polymer.

2.

What are the monomers that make up DNA called?

Nucleotides.

3.

How many different types of nucleotide monomer are used to make DNA?

Four.

4.

What type of molecules are the repeating units in starch based on?

Sugars.

5.

What are proteins polymers of?

Amino acids.

6.

Which monomers form the polymer chains of DNA, starch and proteins respectively?

DNA → nucleotides; starch → sugars; proteins → amino acids.

9.26C Recall the formulae of molecules of the alcohols, methanol, ethanol, propanol (propan-1-ol only) and butanol (butan-1-ol only), and draw the structures of these molecules, showing all covalent bonds

1.

What is the molecular formula of methanol?

CH₃OH

2.

What is the molecular formula of ethanol?

C₂H₅OH

3.

What is the molecular formula of propan-1-ol?

C₃H₇OH

4.

What is the molecular formula of butan-1-ol?

C₄H₉OH

5.

What is the general molecular formula of the alcohols methanol, ethanol, propan-1-ol and butan-1-ol?

CnH2n+1OH

6.

How should the structures of methanol, ethanol, propan-1-ol and butan-1-ol be represented to show all covalent bonds?

Methanol: CH₃OH; ethanol: CH₃CH₂OH; propan-1-ol: CH₃CH₂CH₂OH; butan-1-ol: CH₃CH₂CH₂CH₂OH.

9.27C Recall that the functional group in alcohols is –OH and that alcohols can be dehydrated to form alkenes

1.

What is the functional group present in alcohols?

–OH

2.

What is the name of the functional group –OH in an alcohol?

Hydroxyl group.

3.

What happens when an alcohol is dehydrated?

Water is removed.

4.

What type of hydrocarbon is formed when an alcohol is dehydrated?

An alkene.

5.

What structural feature is introduced when an alcohol is dehydrated to form an alkene?

A C=C double bond.

6.

Why does dehydration of an alcohol produce an unsaturated hydrocarbon?

Removing H₂O leaves a carbon-carbon double bond.

9.28C Core Practical: Investigate the temperature rise produced in a known mass of water by the combustion of the alcohols ethanol, propanol, butanol and pentanol

1.

How can the temperature rise produced by burning an alcohol be investigated experimentally?

Burn a measured mass or volume of alcohol beneath a container of a known mass of water and measure the temperature rise.

2.

Why should a known mass of water be used when comparing the temperature rises produced by different alcohols?

It allows a fair comparison because the same amount of water is heated each time.

3.

What measurements are needed to determine the temperature rise of water when an alcohol is burned?

Initial temperature, final temperature and mass of water.

4.

How can the temperature rise produced by ethanol, propanol, butanol and pentanol be compared fairly?

Use the same mass of water, same apparatus, same starting temperature and same burner position.

5.

What variable should be kept constant when comparing the energy released by different alcohols using this experiment?

The mass of water and experimental setup should remain constant.

6.

Why should the alcohol burner be positioned at a fixed distance from the water container during this investigation?

To keep the transfer of heat to the water as similar as possible between experiments.

9.29C Recall the formulae of molecules of the carboxylic acids, methanoic, ethanoic, propanoic and butanoic acids, and draw the structures of these molecules, showing all covalent bonds

1.

What is the molecular formula of methanoic acid?

HCOOH

2.

What is the molecular formula of ethanoic acid?

CH₃COOH

3.

What is the molecular formula of propanoic acid?

CH₃CH₂COOH

4.

What is the molecular formula of butanoic acid?

CH₃CH₂CH₂COOH

5.

What is the general molecular formula of the carboxylic acids methanoic, ethanoic, propanoic and butanoic acid?

CnH2n+1COOH

6.

How should the structures of methanoic, ethanoic, propanoic and butanoic acid be represented to show all covalent bonds?

Methanoic acid: HCOOH; ethanoic acid: CH₃COOH; propanoic acid: CH₃CH₂COOH; butanoic acid: CH₃CH₂CH₂COOH.

9.30C Recall that the functional group in carboxylic acids is –COOH and that solutions of carboxylic acids have typical acidic properties

1.

What is the functional group present in carboxylic acids?

–COOH

2.

What is the functional group –COOH called?

Carboxyl group.

3.

Why do solutions of carboxylic acids have acidic properties?

They partially dissociate in water to produce H⁺ ions.

4.

What type of ions are produced when carboxylic acids dissolve in water?

H⁺ ions.

5.

How can the acidic nature of a carboxylic acid solution be demonstrated using an indicator?

They turn blue litmus red or give other acidic indicator results.

6.

Which structural feature is responsible for the acidic properties of carboxylic acids?

The –COOH group.

9.31C Recall that ethanol can be oxidised to produce ethanoic acid and extend this to other alcohols (reagents not required)

1.

What product is formed when ethanol is oxidised?

Ethanoic acid.

2.

What type of compound is ethanoic acid?

A carboxylic acid.

3.

What happens to the functional group of ethanol when ethanol is oxidised?

It is oxidised to the –COOH group.

4.

What product is formed when a suitable primary alcohol is oxidised?

A carboxylic acid.

5.

What type of alcohol can be oxidised to form a carboxylic acid?

A primary alcohol.

6.

How can the oxidation of ethanol be extended to other suitable alcohols?

Other primary alcohols can similarly be oxidised to the corresponding carboxylic acids.

9.32C Recall members of a given homologous series have similar reactions because their molecules contain the same functional group and use this to predict the products of other members of these series

1.

Why do members of the same homologous series have similar chemical reactions?

They contain the same functional group.

2.

What structural feature do members of the same homologous series share?

The same functional group.

3.

How does the functional group determine the chemical reactions of members of a homologous series?

It is responsible for the main chemical reactions.

4.

How can the reaction of one member of a homologous series be used to predict the reaction of another member?

Identify the functional group and apply the known reaction to the new member.

5.

What information about a homologous series is needed to predict the products of reactions of its members?

The homologous series and its functional group.

6.

Why can members of the same homologous series have similar chemical properties despite having different molecular masses?

Their functional groups are the same.

9.33C Describe the production of ethanol by fermentation of carbohydrates in aqueous solution, using yeast to provide enzymes

1.

How is ethanol produced by fermentation?

Ferment carbohydrates using yeast in aqueous solution.

2.

What type of substances are fermented to produce ethanol?

Carbohydrates such as sugars.

3.

What role does yeast play in the production of ethanol by fermentation?

Yeast provides enzymes that catalyse fermentation.

4.

Why must the carbohydrates be present in aqueous solution during fermentation?

Water provides the medium in which the yeast enzymes and substrates can react.

5.

What products are formed when carbohydrates undergo fermentation?

Ethanol and carbon dioxide.

6.

What biological catalyst in yeast carries out the reactions involved in fermentation?

Enzymes in yeast.

9.34C Explain how to obtain a concentrated solution of ethanol by fractional distillation of the fermentation mixture

1.

Why is fractional distillation used to obtain a concentrated solution of ethanol from a fermentation mixture?

Ethanol and water have different boiling points.

2.

What property of ethanol and water allows them to be separated by fractional distillation?

Their different boiling points.

3.

What happens to the ethanol-water mixture during fractional distillation?

The mixture is heated and vapours condense at different temperatures.

4.

Why can fractional distillation produce a more concentrated ethanol solution than the original fermentation mixture?

Ethanol has a lower boiling point than water, so the distillate becomes richer in ethanol.

5.

Why can fractional distillation not normally produce completely pure ethanol from an ethanol-water mixture?

Ethanol and water cannot normally be completely separated by simple fractional distillation.

6.

How does the boiling point of ethanol compare with that of water during fractional distillation?

Ethanol has a lower boiling point than water.

9.35C Compare the size of nanoparticles with the sizes of atoms and molecules

1.

What is the typical size range of nanoparticles?

About 1 nm to 100 nm.

2.

How does the size of a nanoparticle compare with the size of an atom?

Nanoparticles are much larger than individual atoms.

3.

How does the size of a nanoparticle compare with the size of a small molecule?

Nanoparticles are generally larger than small molecules.

4.

Why are nanoparticles described as being on the nanoscale?

Their dimensions are measured in nanometres.

5.

What unit is commonly used to describe the size of nanoparticles?

Nanometres (nm).

6.

How does the size scale of nanoparticles differ from that of ordinary bulk materials?

Nanoparticles are much smaller than ordinary bulk materials.

9.36C Describe how the properties of nanoparticulate materials are related to their uses including surface area to volume ratio of the particles they contain, including sunscreens

1.

Why do nanoparticles have a high surface area to volume ratio?

Their particles are very small, giving a large surface area relative to volume.

2.

How can the high surface area to volume ratio of nanoparticles affect their properties?

It can increase reactivity and alter optical or other physical properties.

3.

Why are nanoparticles useful in some applications where larger particles of the same material would not be as effective?

They can have properties different from the same material in bulk form.

4.

Why are nanoparticles used in some sunscreens?

They can provide effective UV protection while remaining relatively transparent.

5.

How does particle size affect the surface area to volume ratio of a nanoparticulate material?

Smaller particles give a larger surface area to volume ratio.

6.

How can the properties of nanoparticulate materials determine their uses?

Their properties determine suitable applications.

9.37C Explain the possible risks associated with some nanoparticulate materials

1.

Why can nanoparticulate materials pose potential health risks?

Their very small size may allow them to enter cells or tissues.

2.

Why might nanoparticles behave differently from larger particles of the same substance?

Their high surface area and small size can produce different chemical and biological behaviour.

3.

How could nanoparticles enter the human body?

Through inhalation, ingestion or through damaged skin.

4.

Why can the high surface area to volume ratio of nanoparticles contribute to potential risks?

More surface area is available for chemical interactions.

5.

Why is the long-term effect of some nanoparticles uncertain?

Some nanoparticles have not been studied over sufficiently long timescales.

6.

Why should the risks of nanoparticulate materials be considered when developing new applications?

Potential health and environmental effects must be assessed before widespread use.

9.38C Compare, using data, the physical properties of glass and clay ceramics, polymers, composites and metals

1.

Which physical properties could be compared when evaluating glass, clay ceramics, polymers, composites and metals?

Strength, density, hardness, melting point, electrical conductivity, thermal conductivity and flexibility.

2.

How can data be used to compare the physical properties of different materials?

Compare numerical or measured values for the relevant properties.

3.

Why might a metal have different physical properties from a polymer?

Their different structures and bonding give different physical properties.

4.

How can the physical properties of glass affect its suitability for particular uses?

Its hardness, brittleness, transparency and chemical resistance affect its uses.

5.

How can the physical properties of clay ceramics affect their suitability for particular uses?

Their hardness, heat resistance and brittleness affect their uses.

6.

Why are composites often used when the properties of a single material are insufficient?

They combine useful properties from different materials.

9.39C Explain why the properties of a material make it suitable for a given use and use data to select materials appropriate for specific uses

1.

Why must the properties of a material be considered when selecting it for a particular use?

The material must have properties that meet the requirements of the application.

2.

How can data be used to select the most suitable material for a specific application?

Compare the supplied data with the properties required for the application.

3.

Which properties would be important when selecting a material for an electrical conductor?

High electrical conductivity, suitable strength and low resistance.

4.

Which properties would be important when selecting a material for a lightweight structural component?

Low density combined with high strength.

5.

Why might a material with a particular combination of properties be preferred over a material with one superior property?

The overall combination may better satisfy the application's requirements.

6.

How can experimental or supplied data be used to justify the choice of a material for a specific use?

Identify the required properties, compare the data and justify the material that best meets the requirements.

Chemistry

Edexcel GCSE Triple Science: Chemistry

Topic 1 – Key concepts in chemistry

1.1 Describe how the Dalton model of an atom has changed over time because of the discovery of subatomic particles

1.2 Describe the structure of an atom as a nucleus containing protons and neutrons, surrounded by electrons in shells

1.3 Recall the relative charge and relative mass of: a a proton b a neutron c an electron

1.4 Explain why atoms contain equal numbers of protons and electrons

1.5 Describe the nucleus of an atom as very small compared to the overall size of the atom

1.6 Recall that most of the mass of an atom is concentrated in the nucleus

1.7 Recall the meaning of the term mass number of an atom

1.8 Describe atoms of a given element as having the same number of protons in the nucleus and that this number is unique to that element

1.9 Describe isotopes as different atoms of the same element containing the same number of protons but different numbers of neutrons in their nuclei

1.10 Calculate the numbers of protons, neutrons and electrons in atoms given the atomic number and mass number

1.11 Explain how the existence of isotopes results in relative atomic masses of some elements not being whole numbers

1.12 Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes

1.13 Describe how Dmitri Mendeleev arranged the elements, known at that time, in a periodic table by using properties of these elements and their compounds

1.14 Describe how Dmitri Mendeleev used his table to predict the existence and properties of some elements not then discovered

1.15 Explain that Dmitri Mendeleev thought he had arranged elements in order of increasing relative atomic mass but this was not always true because of the relative abundance of isotopes of some pairs of elements in the periodic table

1.16 Explain the meaning of atomic number of an element in terms of position in the periodic table and number of protons in the nucleus

1.17 Describe that in the periodic table a elements are arranged in order of increasing atomic number, in rows called periods b elements with similar properties are placed in the same vertical columns called groups

1.18 Identify elements as metals or non-metals according to their position in the periodic table, explaining this division in terms of the atomic structures of the elements

1.19 Predict the electronic configurations of the first 20 elements in the periodic table as diagrams and in the form, for example 2.8.1

1.20 Explain how the electronic configuration of an element is related to its position in the periodic table

1.21 Explain how ionic bonds are formed by the transfer of electrons between atoms to produce cations and anions, including the use of dot and cross diagrams

1.22 Recall that an ion is an atom or group of atoms with a positive or negative charge

1.23 Calculate the numbers of protons, neutrons and electrons in simple ions given the atomic number and mass number

1.24 Explain the formation of ions in ionic compounds from their atoms, limited to compounds of elements in groups 1, 2, 6 and 7

1.25 Explain the use of the endings –ide and –ate in the names of compounds

1.26 Deduce the formulae of ionic compounds (including oxides, hydroxides, halides, nitrates, carbonates and sulfates) given the formulae of the constituent ions

1.27 Explain the structure of an ionic compound as a lattice structure: a consisting of a regular arrangement of ions b held together by strong electrostatic forces (ionic bonds) between oppositely-charged ions

1.28 Explain how a covalent bond is formed when a pair of electrons is shared between two atoms

1.29 Recall that covalent bonding results in the formation of molecules

1.30 Recall the typical size (order of magnitude) of atoms and small molecules

1.31 Explain the formation of simple molecular, covalent substances, using dot and cross diagrams, including: a hydrogen b hydrogen chloride c water d methane e oxygen f carbon dioxide

1.33 Explain the properties of ionic compounds limited to: a high melting points and boiling points, in terms of forces between ions b whether or not they conduct electricity as solids, when molten and in aqueous solution

1.34 Explain the properties of typical covalent, simple molecular compounds limited to: a low melting points and boiling points, in terms of forces between molecules (intermolecular forces) b poor conduction of electricity

1.35 Recall that graphite and diamond are different forms of carbon and that they are examples of giant covalent substances

1.36 Describe the structures of graphite and diamond

1.37 Explain, in terms of structure and bonding, why graphite is used to make electrodes and as a lubricant, whereas diamond is used in cutting tools

1.38 Explain the properties of fullerenes including C60 and graphene in terms of their structures and bonding

1.39 Describe, using poly(ethene) as the example, that simple polymers consist of large molecules containing chains of carbon atoms

1.40 Explain the properties of metals, including malleability and the ability to conduct electricity

1.41 Describe the limitations of particular representations and models, to include dot and cross, ball and stick models and two- and three-dimensional representations

1.42 Describe most metals as shiny solids which have high melting points, high density and are good conductors of electricity whereas most non-metals have low boiling points and are poor conductors of electricity

1.43 Calculate: a relative formula mass given relative atomic masses b percentage by mass of an element in a compound given relative atomic masses

1.44 Calculate the formulae of simple compounds from reacting masses or percentage composition and understand that these are empirical formulae

1.45 Deduce: a the empirical formula of a compound from the formula of its molecule b the molecular formula of a compound from its empirical formula and its relative molecular mass

1.46 Describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide

1.47 Explain the law of conservation of mass applied to: a a closed system including a precipitation reaction in a closed flask b a non-enclosed system including a reaction in an open flask that takes in or gives out a gas

1.48 Calculate masses of reactants and products from balanced equations, given the mass of one substance

1.49 Calculate the concentration of solutions in g dm⁻³

1.50 Recall that one mole of particles of a substance is defined as: a the Avogadro constant number of particles (6.02 × 10²³ atoms, molecules, formulae or ions) of that substance b a mass of 'relative particle mass'

1.51 Calculate the number of: a moles of particles of a substance in a given mass of that substance and vice versa b particles of a substance in a given number of moles of that substance and vice versa c particles of a substance in a given mass of that substance and vice versa

1.52 Explain why, in a reaction, the mass of product formed is controlled by the mass of the reactant which is not in excess

1.53 Deduce the stoichiometry of a reaction from the masses of the reactants and products

Topic 2 – States of matter and mixtures

2.1 Describe the arrangement, movement and the relative energy of particles in each of the three states of matter: solid, liquid and gas

2.2 Recall the names used for the interconversions between the three states of matter, recognising that these are physical changes: contrasted with chemical reactions that result in chemical changes

2.3 Explain the changes in arrangement, movement and energy of particles during these interconversions

2.4 Predict the physical state of a substance under specified conditions, given suitable data

2.5 Explain the difference between the use of 'pure' in chemistry compared with its everyday use and the differences in chemistry between a pure substance and a mixture

2.6 Interpret melting point data to distinguish between pure substances which have a sharp melting point and mixtures which melt over a range of temperatures

2.7 Explain the types of mixtures that can be separated by using the following experimental techniques: a simple distillation b fractional distillation c filtration d crystallisation e paper chromatography

2.8 Describe an appropriate experimental technique to separate a mixture, knowing the properties of the components of the mixture

2.9 Describe paper chromatography as the separation of mixtures of soluble substances by running a solvent (mobile phase) through the mixture on the paper (the paper contains the stationary phase), which causes the substances to move at different rates over the paper

2.10 Interpret a paper chromatogram: a to distinguish between pure and impure substances b to identify substances by comparison with known substances c to identify substances by calculation and use of Rf values

2.11 Core Practical: Investigate the composition of inks using simple distillation and paper chromatography

2.12 Describe how: a waste and ground water can be made potable, including the need for sedimentation, filtration and chlorination b sea water can be made potable by using distillation c water used in analysis must not contain any dissolved salts

Topic 3 – Chemical changes

3.1 Recall that acids in solution are sources of hydrogen ions and alkalis in solution are sources of hydroxide ions

3.2 Recall that a neutral solution has a pH of 7 and that acidic solutions have lower pH values and alkaline solutions higher pH values

3.3 Recall the effect of acids and alkalis on indicators, including litmus, methyl orange and phenolphthalein

3.4 Recall that the higher the concentration of hydrogen ions in an acidic solution, the lower the pH; and the higher the concentration of hydroxide ions in an alkaline solution, the higher the pH

3.5 Recall that as hydrogen ion concentration in a solution increases by a factor of 10, the pH of the solution decreases by 1

3.6 Core Practical: Investigate the change in pH on adding powdered calcium hydroxide or calcium oxide to a fixed volume of dilute hydrochloric acid

3.7 Explain the terms dilute and concentrated, with respect to amount of substances in solution

3.8 Explain the terms weak and strong acids, with respect to the degree of dissociation into ions

3.9 Recall that a base is any substance that reacts with an acid to form a salt and water only

3.10 Recall that alkalis are soluble bases

3.11 Explain the general reactions of aqueous solutions of acids with: a metals b metal oxides c metal hydroxides d metal carbonates to produce salts