Edexcel Combined 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?

Dalton's original model described atoms as indivisible, solid spheres.

2.

What discovery showed that atoms are not indivisible solid spheres as Dalton proposed?

The discovery of electrons showed that atoms contain smaller subatomic particles.

3.

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

Thomson proposed a model in which electrons were embedded in a sphere of positive charge.

4.

What discovery led to the nuclear model of the atom?

Rutherford's alpha-scattering experiment provided evidence for a small, dense, positively charged nucleus.

5.

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

The nuclear model showed that most of the atom is empty space, with electrons surrounding a tiny nucleus.

6.

How has the modern model of the atom developed from Dalton's original model as a result of discoveries of subatomic particles?

The modern model developed from Dalton's model through the discovery of electrons, the nucleus, protons and neutrons.

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

1.

What particles are found in the nucleus of an atom?

Protons and neutrons are found in the nucleus.

2.

Where are electrons found in an atom?

Electrons are found in shells surrounding the nucleus.

3.

What is meant by an electron shell?

An electron shell is a region around the nucleus occupied by electrons.

4.

What is the structure of the nucleus of an atom?

The nucleus contains protons and neutrons.

5.

Where is most of the mass of an atom located?

Most of the mass of an atom is located in the nucleus.

6.

Describe the overall structure of an atom in terms of its nucleus, protons, neutrons and electrons.

An atom consists of a tiny nucleus containing protons and neutrons, surrounded by electrons arranged in shells.

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

1.

What are the relative charge and relative mass of a proton?

A proton has a relative charge of +1 and a relative mass of 1.

2.

What are the relative charge and relative mass of a neutron?

A neutron has a relative charge of 0 and a relative mass of 1.

3.

What are the relative charge and relative mass of an electron?

An electron has a relative charge of −1 and a relative mass of approximately 1/1840.

4.

Which subatomic particle has a relative charge of −1 and a relative mass of approximately 1/1840?

The subatomic particle with a relative charge of −1 and relative mass of approximately 1/1840 is an electron.

5.

Which two subatomic particles have approximately the same relative mass?

Protons and neutrons have approximately the same relative mass.

6.

What are the relative charges of a proton, neutron and electron?

The relative charges of a proton, neutron and electron are +1, 0 and −1 respectively.

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 contains equal numbers of protons and electrons because their charges are equal in magnitude and opposite in sign.

2.

What charge does a proton have?

A proton has a relative charge of +1.

3.

What charge does an electron have?

An electron has a relative charge of −1.

4.

What would happen to the overall charge of an atom if it contained more protons than electrons?

More protons than electrons would give the atom an overall positive charge.

5.

What would happen to the overall charge of an atom if it contained more electrons than protons?

More electrons than protons would give the atom an overall negative charge.

6.

Explain why an atom with equal numbers of protons and electrons has no overall charge.

Equal numbers of +1 protons and −1 electrons cancel each other's charges, giving the atom no overall charge.

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 overall size of the atom.

2.

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

The nucleus occupies only a very small part of the overall size of an atom.

3.

Where are the electrons located relative to the tiny nucleus?

Electrons are found in shells surrounding the tiny nucleus.

4.

Why can the nucleus be described as very small compared with the overall size of an atom?

The nucleus is very small compared with the region occupied by the electrons.

5.

What does the small size of the nucleus indicate about the amount of empty space within an atom?

The small size of the nucleus indicates that most of an atom is empty space.

6.

How does the size of the nucleus compare with the distance over which the electrons are found in an atom?

The nucleus is tiny compared with the much larger region over which the electrons are found.

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?

Most of the mass of an atom is concentrated in the nucleus.

2.

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

Protons and neutrons account for almost all of an atom's mass.

3.

Why do electrons contribute very little to the mass of an atom?

Electrons contribute very little to the mass of an atom because their relative mass is approximately 1/1840.

4.

How does the mass of the nucleus compare with the mass of the electrons in an atom?

The nucleus has almost all the mass of the atom, while the electrons have very little mass.

5.

What relationship does the mass of an atom have with the number of protons and neutrons in its nucleus?

The mass of an atom is approximately determined by the total number of protons and neutrons in its nucleus.

6.

Why is the nucleus described as containing most of the mass of an atom?

The nucleus contains most of the mass because it contains the relatively massive protons and neutrons.

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

1.

What is meant by the mass number of an atom?

The mass number is the total number of protons and neutrons in an atom.

2.

How is the mass number of an atom calculated from its protons and neutrons?

Mass number = number of protons + number of neutrons.

3.

Which subatomic particles are included in the mass number?

Protons and neutrons are included in the mass number.

4.

Why are electrons not included when calculating the mass number?

Electrons are not included because their mass is negligible compared with the masses of protons and neutrons.

5.

An atom contains 11 protons and 12 neutrons. What is its mass number?

An atom containing 11 protons and 12 neutrons has a mass number of 23.

6.

An atom has a mass number of 23 and contains 11 protons. How many neutrons does it contain?

An atom with mass number 23 and 11 protons contains 12 neutrons.

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 the nucleus determines which element an atom belongs to.

2.

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

All atoms of the same element have the same number of protons.

3.

What is meant by the atomic number of an element?

The atomic number is the number of protons in the nucleus of an atom.

4.

Why is the number of protons unique to each element?

Each element has a unique number of protons.

5.

An atom contains 17 protons. Which element does it represent?

An atom containing 17 protons is chlorine.

6.

Can two different elements have atoms containing the same number of protons? Explain your answer.

No two different elements can have atoms with the same number of protons because the proton number defines the 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.

What is an isotope?

An isotope is a different atom of the same element with the same number of protons but a different number of neutrons.

2.

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

Isotopes have the same number of protons and therefore have the same chemical identity.

3.

What is the same about the nuclei of isotopes of the same element?

The nuclei of isotopes of the same element contain the same number of protons.

4.

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

The nuclei of isotopes contain different numbers of neutrons.

5.

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

Increasing the number of neutrons increases the mass number of an isotope.

6.

Explain why carbon-12 and carbon-14 are isotopes of the same element.

Carbon-12 and carbon-14 are isotopes because both have 6 protons but carbon-12 has 6 neutrons whereas carbon-14 has 8 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?

The number of protons is equal to the atomic number.

2.

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.

3.

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

In a neutral atom, the number of electrons equals the number of protons.

4.

An atom has atomic number 13 and mass number 27. How many protons, neutrons and electrons does it contain?

Atomic number 13, mass number 27: 13 protons, 14 neutrons and 13 electrons.

5.

An atom has atomic number 17 and mass number 35. Calculate the numbers of protons, neutrons and electrons in the atom.

Atomic number 17, mass number 35: 17 protons, 18 neutrons and 17 electrons.

6.

An atom has 19 protons and a mass number of 39. Calculate the numbers of neutrons and electrons in the atom.

An atom with 19 protons and mass number 39 has 20 neutrons and 19 electrons.

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

1.

Why can the relative atomic mass of an element be a non-whole number?

Relative atomic mass can be a non-whole number because an element can exist as a mixture of isotopes with different masses.

2.

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

The different isotopes of an element contribute to its overall relative atomic mass.

3.

Why is the relative atomic mass of an element usually a weighted mean rather than the mass of one particular atom?

Relative atomic mass is a weighted mean of the masses of the naturally occurring isotopes.

4.

Why is the relative atomic mass of chlorine approximately 35.5 rather than a whole number?

Chlorine has isotopes with different masses and abundances, giving chlorine a relative atomic mass of approximately 35.5.

5.

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

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

6.

Explain why an element consisting of two isotopes with different masses can have a relative atomic mass between the two isotope masses.

The relative atomic mass lies between the masses of the isotopes when an element contains isotopes with different masses.

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 the masses and abundances of its isotopes?

Relative atomic mass = (isotope mass × abundance + isotope mass × abundance + ...) ÷ total abundance.

2.

An element has two isotopes with masses 10 and 11, present in abundances of 20% and 80%. Calculate its relative atomic mass.

(10 × 20 + 11 × 80) ÷ 100 = 10.8

3.

An element has two isotopes with masses 63 and 65, present in abundances of 70% and 30%. Calculate its relative atomic mass.

(63 × 70 + 65 × 30) ÷ 100 = 63.6

4.

An element has two isotopes with masses 24 and 26, present in abundances of 75% and 25%. Calculate its relative atomic mass.

(24 × 75 + 26 × 25) ÷ 100 = 24.5

5.

An element has isotopes of relative masses 35 and 37 with abundances of 75% and 25%. Calculate its relative atomic mass.

(35 × 75 + 37 × 25) ÷ 100 = 35.5

6.

An element has three isotopes with relative masses 20, 21 and 22 and abundances of 90%, 5% and 5%. Calculate its relative atomic mass.

(20 × 90 + 21 × 5 + 22 × 5) ÷ 100 = 20.15

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 known at the time in his periodic table?

Mendeleev arranged the known elements into a periodic table based largely on their relative atomic masses and chemical properties.

2.

What properties did Mendeleev use to help arrange the elements in his periodic table?

He used the properties of elements to identify patterns and similarities.

3.

Why did Mendeleev leave gaps in his periodic table?

Mendeleev left gaps where no known element fitted the pattern.

4.

How did Mendeleev use the properties of compounds to help identify patterns between elements?

He considered the properties of compounds when identifying similarities between elements.

5.

Why was Mendeleev's arrangement of elements an improvement on earlier attempts to classify the elements?

His arrangement improved classification by placing elements with similar properties together and revealing repeating patterns.

6.

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

Elements with similar properties were placed into the same groups.

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 to predict the existence of undiscovered elements?

Mendeleev left gaps in his periodic table for elements that he predicted had not yet been discovered.

2.

What properties of an undiscovered element could Mendeleev predict from its position in his periodic table?

He used the position of a gap to predict the existence of an undiscovered element.

3.

How could Mendeleev predict the relative atomic mass of an undiscovered element?

He could predict properties such as relative atomic mass and chemical properties from an element's position.

4.

Why was the later discovery of elements with properties predicted by Mendeleev strong evidence supporting his periodic table?

The later discovery of elements with properties matching Mendeleev's predictions supported his periodic table.

5.

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

These discoveries demonstrated that his periodic table could successfully predict previously unknown elements.

6.

Why did Mendeleev leave gaps rather than forcing all known elements into an incorrect pattern?

Mendeleev left gaps because he prioritised the pattern of chemical properties rather than forcing all known elements into incorrect positions.

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?

Mendeleev thought he had arranged elements in order of increasing relative atomic mass because he used relative atomic mass as an important basis for his arrangement.

2.

Why is increasing relative atomic mass not always the same as increasing atomic number?

Relative atomic mass and atomic number do not always increase in exactly the same order.

3.

How can the relative abundance of isotopes cause one element to have a greater relative atomic mass than the next element in atomic-number order?

Differences in the relative abundances of isotopes can give one element a greater relative atomic mass than another element with a higher atomic number.

4.

Why does the existence of isotopes explain some apparent anomalies in Mendeleev's ordering?

Isotopes with different masses and abundances can produce apparent anomalies in relative atomic mass ordering.

5.

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

The modern periodic table is arranged by increasing atomic number because atomic number directly identifies the element.

6.

How does the relative abundance of isotopes affect the relative atomic mass of an element and explain why some pairs of elements appear out of mass order?

The relative abundance of isotopes affects the weighted mean relative atomic mass and can explain why some pairs of elements do not follow increasing mass 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 is meant by the atomic number of an element?

Atomic number is the number of protons in the nucleus of an atom.

2.

How is the atomic number related to the number of protons in the nucleus of an atom?

The atomic number is equal to the number of protons in the nucleus.

3.

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

Elements are positioned in the periodic table according to increasing atomic number.

4.

Why does every element have a unique atomic number?

Every element has a unique atomic number because every element has a unique number of protons.

5.

An element has atomic number 12. How many protons are in the nucleus of each of its atoms?

An element with atomic number 12 has 12 protons.

6.

An atom contains 16 protons. What is its atomic number and what does this tell you about its position in the periodic table?

An atom with 16 protons has atomic number 16 and occupies the position of sulfur 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.

In what order are elements arranged in the modern periodic table?

Elements in the modern periodic table are arranged in order of increasing atomic number.

2.

What is a period in the periodic table?

A period is a horizontal row in the periodic table.

3.

What is a group in the periodic table?

A group is a vertical column in the periodic table.

4.

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

Elements in the same group have similar chemical properties.

5.

How does moving from left to right across a period affect the atomic number of the elements?

Moving from left to right across a period, the atomic number generally increases by one.

6.

What is the difference between a period and a group in the periodic table?

A period is a horizontal row, whereas a group is a vertical column containing elements with 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?

Metals are generally found on the left and centre of the periodic table.

2.

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

Non-metals are generally found on the right-hand side of the periodic table.

3.

What feature of atomic structure generally distinguishes metals from non-metals?

Metals generally have atoms that can lose electrons relatively easily, whereas non-metals generally gain or share electrons.

4.

Why do metals generally tend to lose electrons when they react?

Metals tend to lose electrons because they generally have fewer outer-shell electrons.

5.

Why do non-metals generally tend to gain or share electrons when they react?

Non-metals tend to gain or share electrons to achieve a more stable outer electron arrangement.

6.

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

The position of an element in the periodic table can therefore be used to predict whether it is likely to be a metal or non-metal.

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.

What is meant by the electronic configuration of an atom?

Electronic configuration describes how the electrons in an atom are arranged in electron shells.

2.

How many electrons can occupy the first electron shell?

The first electron shell can contain a maximum of 2 electrons.

3.

How many electrons can occupy the second electron shell for the first 20 elements?

For the first 20 elements, the second shell can contain a maximum of 8 electrons.

4.

What is the electronic configuration of sodium?

Sodium has the electronic configuration 2.8.1.

5.

What is the electronic configuration of chlorine?

Chlorine has the electronic configuration 2.8.7.

6.

What is the electronic configuration of calcium?

Calcium has the electronic configuration 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 the period number of an element?

The number of occupied electron shells corresponds to the period number.

2.

How is the number of electrons in the outer shell related to the group number for the main groups?

For the main groups, the number of electrons in the outer shell corresponds to the group number, apart from helium.

3.

What can the electronic configuration of an element tell you about its position in the periodic table?

Electronic configuration can therefore be used to identify an element's position in the periodic table.

4.

An element has an electronic configuration of 2.8.1. What period and group is the element in?

An element with electronic configuration 2.8.1 is in Group 1 and Period 3.

5.

An element has an electronic configuration of 2.8.7. What period and group is the element in?

An element with electronic configuration 2.8.7 is in Group 7 and Period 3.

6.

An element is in Group 2 and Period 3. What electronic configuration would you predict for this element?

An element in Group 2 and Period 3 has the electronic configuration 2.8.2.

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.

What happens to electrons when an ionic bond forms between a metal atom and a non-metal atom?

Ionic bonding involves the transfer of electrons from one atom to another.

2.

How does an atom become a cation during ionic bonding?

A metal atom loses electrons to form a positively charged cation.

3.

How does an atom become an anion during ionic bonding?

A non-metal atom gains electrons to form a negatively charged anion.

4.

Why does electron transfer allow the atoms involved in ionic bonding to achieve more stable electronic configurations?

Electron transfer allows the resulting ions to achieve more stable electron configurations.

5.

How can a dot and cross diagram show the transfer of electrons during ionic bonding?

Dot and cross diagrams show the outer electrons from each atom and the transfer of electrons between them.

6.

Using a dot and cross diagram, how would you show the formation of sodium chloride from sodium and chlorine atoms?

In sodium chloride, sodium transfers one electron to chlorine, forming Na⁺ and Cl⁻ ions with stable outer electron arrangements.

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

2.

How does an atom become a positive ion?

An atom becomes a positive ion when it loses one or more electrons.

3.

How does an atom become a negative ion?

An atom becomes a negative ion when it gains one or more electrons.

4.

What is a cation?

A cation is a positively charged ion.

5.

What is an anion?

An anion is a negatively charged ion.

6.

Can an ion consist of a group of atoms rather than a single atom?

An ion can consist of a group of atoms, such as sulfate, SO₄²⁻.

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 an ion be determined from its atomic number?

The number of protons in an ion is equal to its atomic number.

2.

How can the number of neutrons in an 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?

The number of electrons in a positive ion is found by subtracting the positive charge from the atomic number.

4.

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

The number of electrons in a negative ion is found by adding the magnitude of the negative charge to the atomic number.

5.

A sodium ion, Na⁺, has atomic number 11 and mass number 23. Calculate its numbers of protons, neutrons and electrons.

Na⁺: 11 protons, 12 neutrons and 10 electrons.

6.

A chloride ion, Cl⁻, has atomic number 17 and mass number 35. Calculate its numbers of protons, neutrons and electrons.

Cl⁻: 17 protons, 18 neutrons and 18 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.

What ion does a Group 1 metal atom form when it loses one electron?

Group 1 atoms lose one electron to form +1 ions.

2.

What ion does a Group 2 metal atom form when it loses two electrons?

Group 2 atoms lose two electrons to form +2 ions.

3.

What ion does a Group 6 non-metal atom form when it gains two electrons?

Group 6 atoms gain two electrons to form −2 ions.

4.

What ion does a Group 7 non-metal atom form when it gains one electron?

Group 7 atoms gain one electron to form −1 ions.

5.

Explain how magnesium and oxygen atoms form Mg²⁺ and O²⁻ ions.

Magnesium loses two electrons to form Mg²⁺, while oxygen gains two electrons to form O²⁻.

6.

Explain how calcium and chlorine atoms form Ca²⁺ and Cl⁻ ions in calcium chloride.

Calcium forms Ca²⁺ and chlorine forms Cl⁻, so two chloride ions are required for each calcium ion, giving CaCl₂.

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 an ionic compound?

The ending –ide generally indicates a compound containing a simple negative ion.

2.

What does the ending –ate generally indicate in the name of an ionic compound?

The ending –ate generally indicates a compound containing a negative ion containing oxygen.

3.

What is the difference between a chloride ion and a chlorate ion in terms of the naming convention?

Chloride is a simple negative ion, whereas chlorate contains oxygen.

4.

What does the name sodium chloride tell you about the ions present in the compound?

Sodium chloride contains sodium ions and chloride ions.

5.

What does the name sodium sulfate tell you about the type of negative ion present?

Sodium sulfate contains sulfate ions, SO₄²⁻.

6.

How can the ending of the name of an ionic compound help you identify the type of anion present?

The ending –ide or –ate can therefore provide information about the type of negative ion present.

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 you deduce the formula of an ionic compound from the charges of its constituent ions?

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

2.

What is the formula of the ionic compound formed from Na⁺ and O²⁻ ions?

Na⁺ and O²⁻ form Na₂O.

3.

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

Mg²⁺ and Cl⁻ form MgCl₂.

4.

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

Al³⁺ and SO₄²⁻ form Al₂(SO₄)₃.

5.

What is the formula of the ionic compound formed from Ca²⁺ and NO₃⁻ ions?

Ca²⁺ and NO₃⁻ form Ca(NO₃)₂.

6.

What is the formula of the ionic compound formed from Al³⁺ and CO₃²⁻ ions?

Al³⁺ and CO₃²⁻ form Al₂(CO₃)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 meant by a lattice structure in an ionic compound?

An ionic lattice is a regular, repeating arrangement of ions.

2.

How are the ions arranged within an ionic lattice?

Positive and negative ions are arranged in a regular pattern throughout the lattice.

3.

What holds the oppositely charged ions together in an ionic lattice?

Oppositely charged ions are held together by strong electrostatic forces.

4.

What are electrostatic forces in an ionic compound?

Electrostatic forces are forces of attraction between oppositely charged particles.

5.

Why does an ionic compound contain a regular arrangement of both positive and negative ions?

The regular arrangement contains positive and negative ions because each ion is attracted to ions of opposite charge.

6.

Explain how strong electrostatic forces between oppositely charged ions hold an ionic lattice together.

Strong electrostatic attractions between oppositely charged ions hold the ionic lattice together.

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

1.

What is a covalent bond?

A covalent bond is a strong bond formed when two atoms share a pair of electrons.

2.

How is a covalent bond formed between two atoms?

A covalent bond forms when two atoms share electrons.

3.

How many electrons are involved in a single covalent bond?

A single covalent bond contains one shared pair of electrons.

4.

Why does sharing a pair of electrons allow atoms to achieve a more stable electron configuration?

Sharing electrons allows atoms to obtain a more stable electron configuration.

5.

What happens to the shared pair of electrons in a covalent bond?

The shared electrons are attracted to the nuclei of both bonded atoms.

6.

Explain how two chlorine atoms form a covalent bond by sharing a pair of electrons.

Two chlorine atoms each contribute one electron to form one shared pair, producing a covalent bond in Cl₂.

1.29 Recall that covalent bonding results in the formation of molecules

1.

What type of structure is formed when atoms are joined by covalent bonds in simple molecular substances?

Covalent bonding between atoms can form simple molecular substances.

2.

What is a molecule?

A molecule is a group of atoms held together by covalent bonds.

3.

How are the atoms within a molecule held together?

The atoms within a molecule are held together by covalent bonds.

4.

What type of bonding holds the atoms together within a molecule?

Covalent bonds hold the atoms together within a molecule.

5.

What is the difference between a molecule and an individual atom?

A molecule contains two or more atoms chemically bonded together, whereas an individual atom consists of a single atom.

6.

Give two examples of substances that exist as simple covalent molecules.

Examples include hydrogen, H₂, and water, H₂O.

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?

The typical order of magnitude of the size of an atom is 10⁻¹⁰ m.

2.

What is the typical order of magnitude of the size of a small molecule?

The typical order of magnitude of a small molecule is also around 10⁻¹⁰ m.

3.

Which is smaller, an atom or a small molecule?

A small molecule is generally larger than an individual atom because it contains multiple atoms.

4.

Approximately how many metres is the diameter of a typical atom?

A typical atom has a diameter of approximately 10⁻¹⁰ m.

5.

Why is it difficult to observe individual atoms using the naked eye?

Individual atoms are far too small to be observed with the naked eye.

6.

How does the size of an atom compare with the size of an everyday object such as a football?

A football is approximately 10⁸ times larger in diameter than a typical atom.

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 does a dot and cross diagram show the formation of a hydrogen molecule, H₂?

In H₂, each hydrogen atom contributes one electron to form one shared pair.

2.

How does a dot and cross diagram show the formation of a hydrogen chloride molecule, HCl?

In HCl, hydrogen and chlorine share one pair of electrons, giving both atoms a stable outer arrangement.

3.

How does a dot and cross diagram show the formation of a water molecule, H₂O?

In H₂O, oxygen forms two covalent bonds with two hydrogen atoms and has two lone pairs of electrons.

4.

How does a dot and cross diagram show the formation of a methane molecule, CH₄?

In CH₄, carbon forms four covalent bonds with four hydrogen atoms.

5.

How does a dot and cross diagram show the formation of an oxygen molecule, O₂?

In O₂, the two oxygen atoms share two pairs of electrons, forming a double covalent bond.

6.

How does a dot and cross diagram show the formation of a carbon dioxide molecule, CO₂?

In CO₂, carbon forms two double covalent bonds, one with each oxygen atom.

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 classified according to their structure and bonding?

The four main types are ionic, simple molecular covalent, giant covalent and metallic.

2.

How does the structure of an ionic substance differ from that of a simple molecular covalent substance?

Ionic substances contain a lattice of oppositely charged ions, whereas simple molecular covalent substances contain separate molecules.

3.

How does the structure of a giant covalent substance differ from that of a simple molecular covalent substance?

Giant covalent substances contain a giant network of atoms joined by strong covalent bonds.

4.

How does metallic bonding result in metals being able to conduct electricity as solids?

Metals conduct electricity as solids because they contain delocalised electrons that can move through the structure.

5.

Why do substances with different structures and bonding have different melting points and boiling points?

Different structures and bonding produce different strengths of attraction between particles, resulting in different melting and boiling points.

6.

How can the structure and bonding of a substance be used to predict its solubility in water and ability to conduct electricity?

The type of structure and bonding affects whether a substance dissolves in water and whether charged particles or electrons are available to conduct electricity.

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 points?

Ionic compounds have high melting points because strong electrostatic attractions between oppositely charged ions require a lot of energy to overcome.

2.

Why do ionic compounds generally have high boiling points?

Ionic compounds have high boiling points because the strong electrostatic attractions between ions require a lot of energy to overcome.

3.

Why do solid ionic compounds not conduct electricity?

Solid ionic compounds do not conduct electricity because their ions are fixed in position and cannot move.

4.

Why can molten ionic compounds conduct electricity?

Molten ionic compounds conduct electricity because their ions are free to move.

5.

Why can ionic compounds dissolved in water conduct electricity?

Ionic compounds in aqueous solution conduct electricity because their ions are free to move through the solution.

6.

Explain how the movement of ions determines whether an ionic compound can conduct electricity as a solid, molten liquid or aqueous solution.

Ionic compounds conduct electricity only when their ions are able to move, such as when molten or dissolved in water.

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 points?

Simple molecular covalent substances generally have low melting points because the intermolecular forces between molecules are weak.

2.

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

They generally have low boiling points because only weak intermolecular forces need to be overcome.

3.

What are intermolecular forces?

Intermolecular forces are forces of attraction between molecules.

4.

Why does melting a simple molecular substance not require breaking the covalent bonds within its molecules?

Melting a simple molecular substance involves overcoming intermolecular forces rather than breaking the strong covalent bonds within the molecules.

5.

Why do simple molecular covalent substances generally conduct electricity poorly?

Simple molecular covalent substances generally conduct electricity poorly because they do not contain freely moving charged particles.

6.

Explain how weak intermolecular forces account for the relatively low melting and boiling points of simple molecular covalent substances.

Weak intermolecular forces account for the relatively low melting and boiling points of simple molecular covalent substances.

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

1.

What element are graphite and diamond both made from?

Graphite and diamond are both made from carbon.

2.

What is meant by different forms of the same element?

They are different forms, or allotropes, of the element carbon.

3.

Why are graphite and diamond classified as giant covalent substances?

Graphite and diamond are giant covalent substances because their carbon atoms form large structures joined by covalent bonds.

4.

What type of bonding holds the carbon atoms together in graphite?

Carbon atoms in graphite are joined by strong covalent bonds.

5.

What type of bonding holds the carbon atoms together in diamond?

Carbon atoms in diamond are joined by strong covalent bonds.

6.

Why can graphite and diamond have different properties despite both being made only from carbon?

Their different structures and bonding arrangements give graphite and diamond different physical properties.

1.36 Describe the structures of graphite and diamond

1.

How are carbon atoms arranged in the structure of diamond?

Each carbon atom in diamond is covalently bonded to four other carbon atoms.

2.

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

Diamond has a giant three-dimensional structure of carbon atoms.

3.

How are carbon atoms arranged in the structure of graphite?

Each carbon atom in graphite is covalently bonded to three other carbon atoms within a layer.

4.

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

Graphite consists of layers of carbon atoms arranged in hexagonal structures.

5.

What is the structure of graphite made up of between its layers?

Weak forces act between the layers of graphite.

6.

What structural difference between graphite and diamond explains their different physical properties?

Diamond has a rigid three-dimensional network, whereas graphite has layers that can slide over one another.

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?

Graphite contains delocalised electrons that can move through its layers, allowing it to conduct electricity.

2.

Why can graphite be used as an electrode?

Graphite can conduct electricity, so it can be used as an electrode.

3.

Why can graphite be used as a lubricant?

Graphite can act as a lubricant because its layers can slide over one another.

4.

Why is diamond suitable for use in cutting tools?

Diamond is suitable for cutting tools because it is extremely hard.

5.

How does the structure of graphite allow its layers to slide over one another?

The layers of graphite are held together by weak forces, allowing them to move relative to one another.

6.

Explain why the strong covalent bonding throughout diamond makes it suitable for cutting tools, whereas graphite is suitable as a lubricant and electrode.

Diamond has strong covalent bonds throughout its rigid three-dimensional structure, making it very hard, whereas graphite's layered structure allows it to conduct electricity and act as a lubricant.

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

1.

What is a fullerene?

A fullerene is a form of carbon in which carbon atoms form hollow structures such as cages or tubes.

2.

What is the structure of C₆₀?

C₆₀ consists of 60 carbon atoms arranged in a hollow spherical structure.

3.

Why are some fullerenes useful as lubricants?

Some fullerenes can act as lubricants because their structures allow layers or molecules to slide over one another.

4.

What is graphene?

Graphene is a single layer of carbon atoms arranged in a hexagonal lattice.

5.

Why can graphene conduct electricity?

Graphene can conduct electricity because it contains delocalised electrons that can move through the structure.

6.

Explain how the structure and bonding of graphene and fullerenes give them useful properties.

The structures and bonding of fullerenes and graphene give them properties such as low density, strength, conductivity and useful behaviour as lubricants.

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 polymer is a substance made from very large molecules formed from many repeating units.

2.

What is the structure of a poly(ethene) molecule?

Poly(ethene) consists of long chains of carbon atoms with hydrogen atoms attached.

3.

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

The main chain of poly(ethene) contains carbon atoms.

4.

Why is poly(ethene) described as having large molecules?

Poly(ethene) is described as having large molecules because each molecule contains a very large number of atoms.

5.

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

The carbon atoms form a long continuous chain in a poly(ethene) molecule.

6.

What is the relationship between ethene and poly(ethene)?

Poly(ethene) is formed from ethene molecules joining together to form long polymer chains.

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?

Malleability is the ability of a metal to be hammered or bent into different shapes without breaking.

2.

Why can metals be bent and shaped without breaking?

Metals can be bent and shaped because layers of metal atoms can move relative to one another while the metallic bonding remains.

3.

Why can metals conduct electricity?

Metals conduct electricity because they contain delocalised electrons that can move through the structure.

4.

What particles carry electrical charge through a metal?

Delocalised electrons carry electrical charge through a metal.

5.

How does metallic bonding allow the layers of metal atoms to move while keeping the structure together?

Metallic bonding allows layers of atoms to move while the attraction between positive metal ions and delocalised electrons maintains the structure.

6.

Explain how the structure and bonding of metals account for their malleability and ability to conduct electricity.

The arrangement of metal ions and delocalised electrons accounts for the malleability and electrical conductivity of metals.

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 using a dot and cross diagram to represent bonding?

Dot and cross diagrams do not show the actual sizes or three-dimensional shapes of atoms and molecules.

2.

What is one limitation of using a ball and stick model to represent a molecule?

Ball and stick models do not accurately represent the relative sizes of atoms and the distances between them.

3.

Why can a ball and stick model give a misleading impression of the relative sizes of atoms?

Ball and stick models can make atoms appear larger or smaller relative to their actual sizes.

4.

Why can a two-dimensional representation fail to show the true three-dimensional arrangement of atoms?

A two-dimensional representation cannot fully show the three-dimensional arrangement of atoms.

5.

Why can a three-dimensional model still be an imperfect representation of a real chemical structure?

A three-dimensional model is still a simplified representation and cannot show every feature of a real chemical structure accurately.

6.

Why is it useful to use different representations and models when describing chemical structures?

Different representations are useful because each model highlights some features while having limitations in what it can show.

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?

Most metals are shiny solids with high melting points, high density and good electrical conductivity.

2.

Why are most metals good conductors of electricity?

Most metals conduct electricity well because they contain delocalised electrons.

3.

What is the typical appearance of most metals?

Metals generally have a shiny appearance.

4.

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

Most non-metals have relatively low boiling points and are poor conductors of electricity.

5.

How do the typical melting points and densities of metals compare with those of non-metals?

Metals generally have higher melting points and densities than non-metals.

6.

How do the electrical conductivities of most metals compare with those of most non-metals?

Most metals are good electrical conductors, whereas most non-metals are poor electrical 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?

Relative formula mass is calculated by adding the relative atomic masses of all the atoms shown in the formula.

2.

Calculate the relative formula mass of CaCO₃ using Ar values Ca = 40, C = 12 and O = 16.

Mr(CaCO₃) = 40 + 12 + (3 × 16) = 100.

3.

Calculate the relative formula mass of Al₂(SO₄)₃ using Ar values Al = 27, S = 32 and O = 16.

Mr(Al₂(SO₄)₃) = (2 × 27) + (3 × 32) + (12 × 16) = 342.

4.

What calculation is used to determine the percentage by mass of an element in a compound?

Percentage by mass = (mass of element in the formula ÷ relative formula mass) × 100.

5.

Calculate the percentage by mass of oxygen in CO₂ using Ar values C = 12 and O = 16.

Percentage oxygen in CO₂ = (32 ÷ 44) × 100 = 72.7%.

6.

Calculate the percentage by mass of calcium in CaCO₃ using Ar values Ca = 40, C = 12 and O = 16.

Percentage calcium in CaCO₃ = (40 ÷ 100) × 100 = 40.0%.

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?

An empirical formula shows 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 the reacting masses into moles and divide the mole values by the smallest value to obtain the simplest whole-number ratio.

3.

A compound contains 24 g of carbon and 4 g of hydrogen. Calculate its empirical formula.

C: 24 ÷ 12 = 2; H: 4 ÷ 1 = 4 → ratio 1:2 → CH₂.

4.

A compound contains 14 g of nitrogen and 16 g of oxygen. Calculate its empirical formula.

N: 14 ÷ 14 = 1; O: 16 ÷ 16 = 1 → NO.

5.

A compound contains 40.0% carbon, 6.7% hydrogen and 53.3% oxygen by mass. Calculate its empirical formula.

Assume 100 g: C = 40/12 = 3.33; H = 6.7/1 = 6.7; O = 53.3/16 = 3.33 → ratio 1:2:1 → CH₂O.

6.

Why must the masses or percentage compositions of elements be converted into moles before determining an empirical formula?

Masses or percentages must be converted to moles because chemical formulae represent ratios of numbers of atoms, and moles give the number of particles.

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 the subscripts in the molecular formula by their greatest common factor to obtain the simplest whole-number ratio.

2.

What is the empirical formula of C₆H₁₂O₆?

The empirical formula of C₆H₁₂O₆ is CH₂O.

3.

What is the empirical formula of C₂H₄?

The empirical formula of C₂H₄ is CH₂.

4.

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

First calculate the empirical formula mass, then divide the molecular relative mass by the empirical formula mass to find the multiplier.

5.

A compound has empirical formula CH₂O and relative molecular mass 180. Calculate its molecular formula.

CH₂O has empirical formula mass 30; 180 ÷ 30 = 6 → molecular formula C₆H₁₂O₆.

6.

A compound has empirical formula NO₂ and relative molecular mass 92. Calculate its molecular formula.

NO₂ has empirical formula mass 46; 92 ÷ 46 = 2 → molecular formula N₂O₄.

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

1.

How could you experimentally determine the empirical formula of magnesium oxide?

Weigh a clean, dry crucible and lid, add magnesium, then weigh the crucible, lid and magnesium before heating strongly in air.

2.

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

Magnesium is heated strongly in oxygen/air so that it reacts completely to form magnesium oxide.

3.

Why must the crucible and lid be weighed before heating magnesium?

The crucible and lid must be weighed first so that the mass of the magnesium can be determined accurately.

4.

Why is the magnesium oxide heated, cooled and reweighed repeatedly until a constant mass is obtained?

Heating, cooling and reweighing until a constant mass ensures that the magnesium has reacted completely and no further mass increase occurs.

5.

How are the masses of magnesium and oxygen used to calculate the empirical formula of magnesium oxide?

Calculate the mass of magnesium and the mass of oxygen gained, convert each to moles and determine the simplest whole-number ratio.

6.

What safety precaution should be taken when heating magnesium strongly during an experiment to determine the empirical formula of magnesium oxide?

Wear safety goggles and use appropriate heat-resistant equipment; do not look directly at burning magnesium.

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?

The law of conservation of mass states that mass is neither created nor destroyed during a chemical reaction.

2.

Why does the total mass remain constant when a reaction takes place in a closed system?

In a closed system, no substances can enter or leave, so the total mass remains constant.

3.

A precipitation reaction takes place in a sealed flask. Why does the total mass remain unchanged?

In a sealed flask, all substances involved in the precipitation reaction remain inside the system, so the total mass is unchanged.

4.

Why can the measured mass change when a reaction takes place in an open flask and produces a gas?

If a gas is produced and escapes from an open flask, the measured mass of the apparatus and contents decreases.

5.

Why can the measured mass increase when an open system takes in a gas from the surroundings?

If an open system takes in a gas from the surroundings, the measured mass of the apparatus and contents increases.

6.

Explain why mass is conserved in an open system even when the measured mass of the apparatus and contents changes.

Mass is still conserved in an open system because the gas has simply moved into or out of the system rather than being created or destroyed.

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

1.

What information from a balanced chemical equation is needed to calculate the mass of an unknown reactant or product?

Use the balanced equation to determine the mole ratio, convert the known mass to moles, use the ratio to find the unknown moles, then convert to mass.

2.

In the reaction 2Mg + O₂ → 2MgO, what mass of magnesium oxide is produced when 24 g of magnesium reacts completely? Use Ar values Mg = 24 and O = 16.

24 g Mg = 1 mol Mg; ratio Mg:MgO = 1:1; Mr(MgO) = 40 → 40 g MgO.

3.

In the reaction 2H₂ + O₂ → 2H₂O, what mass of water is produced when 4 g of hydrogen reacts completely? Use Ar values H = 1 and O = 16.

4 g H₂ = 2 mol H₂; ratio H₂:H₂O = 1:1 → 2 mol H₂O; 2 × 18 = 36 g H₂O.

4.

In the reaction CaCO₃ → CaO + CO₂, what mass of carbon dioxide is produced when 100 g of calcium carbonate decomposes completely? Use Ar values Ca = 40, C = 12 and O = 16.

100 g CaCO₃ = 1 mol; ratio CaCO₃:CO₂ = 1:1; Mr(CO₂) = 44 → 44 g CO₂.

5.

In the reaction 2Al + 3Cl₂ → 2AlCl₃, what mass of aluminium chloride is produced when 54 g of aluminium reacts completely? Use Ar values Al = 27 and Cl = 35.5.

54 g Al = 2 mol; ratio Al:AlCl₃ = 1:1; Mr(AlCl₃) = 133.5 → 267 g AlCl₃.

6.

In the reaction Fe₂O₃ + 3CO → 2Fe + 3CO₂, what mass of iron is produced from 160 g of iron(III) oxide? Use Ar values Fe = 56, O = 16 and C = 12.

160 g Fe₂O₃ = 1 mol; ratio Fe₂O₃:Fe = 1:2 → 2 mol Fe; mass = 112 g Fe.

1.49 Calculate the concentration of solutions in g dm–3

1.

How is the concentration of a solution in g dm⁻³ calculated?

Concentration (g dm⁻³) = mass of solute (g) ÷ volume of solution (dm³).

2.

A solution contains 10 g of sodium chloride dissolved to make 2 dm³ of solution. Calculate its concentration in g dm⁻³.

10 ÷ 2 = 5 g dm⁻³.

3.

A solution contains 15 g of potassium nitrate in 500 cm³ of solution. Calculate its concentration in g dm⁻³.

500 cm³ = 0.500 dm³; 15 ÷ 0.500 = 30 g dm⁻³.

4.

A solution has a concentration of 20 g dm⁻³. What mass of solute is present in 0.5 dm³ of the solution?

Mass = concentration × volume = 20 × 0.5 = 10 g.

5.

A solution contains 12 g of solute and has a concentration of 8 g dm⁻³. Calculate the volume of the solution in dm³.

Volume = mass ÷ concentration = 12 ÷ 8 = 1.5 dm³.

6.

A solution contains 25 g of solute in 250 cm³ of solution. Calculate its concentration in g dm⁻³.

250 cm³ = 0.250 dm³; 25 ÷ 0.250 = 100 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.

What is meant by one mole of a substance?

One mole is an amount of substance containing the Avogadro constant number of particles.

2.

What is the Avogadro constant?

The Avogadro constant is 6.02 × 10²³ mol⁻¹.

3.

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

One mole contains 6.02 × 10²³ particles.

4.

What types of particles can be counted using the mole?

The mole can count atoms, molecules, formula units and ions.

5.

What is meant by the relative particle mass of a substance?

Relative particle mass is the relative mass of a particle compared with the chosen standard.

6.

What mass in grams is equal to one mole of a substance in terms of its relative particle mass?

One mole of a substance has a mass in grams numerically equal to its 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.

What equation is used to calculate the number of moles from the mass and relative particle mass of a substance?

Number of moles = mass ÷ relative particle mass.

2.

Calculate the number of moles in 18 g of water, H₂O, given Mr(H₂O) = 18.

Moles of H₂O = 18 ÷ 18 = 1.00 mol.

3.

Calculate the mass of 0.50 mol of carbon dioxide, CO₂, given Mr(CO₂) = 44.

Mass = moles × relative particle mass = 0.50 × 44 = 22 g CO₂.

4.

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

Number of particles = number of moles × Avogadro constant.

5.

Calculate the number of molecules in 0.25 mol of oxygen, O₂, using the Avogadro constant 6.02 × 10²³ mol⁻¹.

0.25 × 6.02 × 10²³ = 1.505 × 10²³ molecules of O₂.

6.

Calculate the number of molecules in 9 g of water, H₂O, given Mr(H₂O) = 18 and the Avogadro constant = 6.02 × 10²³ mol⁻¹.

9 g H₂O = 9 ÷ 18 = 0.5 mol; particles = 0.5 × 6.02 × 10²³ = 3.01 × 10²³ molecules.

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 a reactant being in excess?

A reactant in excess is present in a greater amount than is required to react completely with the other reactants.

2.

What is meant by a limiting reactant?

A limiting reactant is the reactant that is completely used up first.

3.

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

The limiting reactant determines the maximum amount of product because the reaction cannot continue once it has been used up.

4.

If 10 g of reactant A is in excess and 5 g of reactant B is completely used up, which reactant controls the amount of product formed?

Reactant B controls the amount of product because it is completely used up.

5.

Why does adding more of a reactant that is already in excess not increase the amount of product formed?

Adding more of a reactant already in excess does not increase the product because the limiting reactant still becomes completely used up first.

6.

Explain why the mass of product formed in a reaction is controlled by the reactant that is not in excess.

The reactant that is not in excess is the limiting reactant, so its amount determines the maximum mass of product formed.

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

1.

What is meant by the stoichiometry of a chemical reaction?

Stoichiometry is the quantitative relationship between the amounts of reactants and products in a chemical reaction.

2.

How can the masses of reactants and products be used to deduce the stoichiometry of a reaction?

Convert the masses of the reactants and products to moles, then compare the mole amounts to obtain the simplest whole-number ratio.

3.

Aluminium reacts with oxygen according to 4Al + 3O₂ → 2Al₂O₃. Calculate the mass of aluminium oxide formed when 135 g of aluminium reacts completely with excess oxygen. Use Ar values Al = 27 and O = 16.

135 g Al = 5 mol; ratio Al:Al₂O₃ = 2:1 → 2.5 mol Al₂O₃; Mr = 102 → 255 g Al₂O₃.

4.

Magnesium reacts with oxygen according to 2Mg + O₂ → 2MgO. Calculate the mass of magnesium oxide formed when 48 g of magnesium reacts completely with excess oxygen. Use Ar values Mg = 24 and O = 16.

48 g Mg = 2 mol; ratio Mg:MgO = 1:1 → 2 mol MgO; Mr = 40 → 80 g MgO.

5.

Calcium carbonate decomposes according to CaCO₃ → CaO + CO₂. Calculate the mass of calcium oxide produced when 250 g of calcium carbonate decomposes completely. Use Ar values Ca = 40, C = 12 and O = 16.

250 g CaCO₃ = 2.5 mol; ratio CaCO₃:CaO = 1:1 → 2.5 mol CaO; Mr = 56 → 140 g CaO.

6.

Iron reacts with sulfur according to Fe + S → FeS. A reaction uses 11.2 g of iron and produces 17.6 g of iron sulfide. Use these masses and the relative atomic masses Fe = 56 and S = 32 to deduce the stoichiometric ratio of iron to sulfur in the reaction.

11.2 g Fe = 0.2 mol; 17.6 g FeS = 0.2 mol. Therefore Fe:S = 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, liquid and gas?

In a solid, particles are closely packed in a fixed, regular arrangement.

2.

How does the movement of particles differ between solids, liquids and gases?

In a liquid, particles are close together but can move past one another.

3.

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

In a gas, particles are far apart and move freely in random directions.

4.

Why do particles in a solid remain in fixed positions rather than moving freely?

Particles in a solid vibrate about fixed positions.

5.

Why can particles in a liquid move past one another while remaining close together?

Particles in a liquid have more movement than particles in a solid while remaining relatively close together.

6.

Why can particles in a gas move freely and spread throughout a container?

Particles in a gas have the greatest movement and generally the greatest kinetic energy of the three states.

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?

Solid → liquid: melting.

2.

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

Liquid → gas: boiling/evaporation.

3.

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

Gas → liquid: condensation.

4.

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

Liquid → solid: freezing.

5.

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

Solid → gas: sublimation.

6.

What is the name of the change from a gas directly to a solid, and how do physical changes differ from chemical changes?

Gas → solid: deposition. These changes of state are physical changes because no new substance is formed, whereas chemical reactions produce chemical changes.

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

1.

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

During melting, particles become less ordered and can move past one another; their energy increases.

2.

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

During evaporation, particles gain energy and move further apart.

3.

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

During condensation, particles lose energy, move closer together and become less free to move.

4.

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

During freezing, particles lose energy and become arranged more closely in fixed positions.

5.

What happens to the particles when a solid sublimes directly into a gas?

During sublimation, particles gain energy, become much further apart and move freely.

6.

What happens to the particles when a gas undergoes deposition to form a solid?

During deposition, particles lose energy and form a more closely packed, ordered solid arrangement.

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

1.

How can melting point and boiling point data be used to predict whether a substance is a solid, liquid or gas at a specified temperature?

Compare the temperature with the melting point and boiling point: below the melting point is solid, between the melting and boiling points is liquid, and above the boiling point is gas.

2.

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

At 10 °C, the substance is a solid.

3.

A substance has a melting point of 20 °C and a boiling point of 80 °C. What state is it in at 50 °C?

At 50 °C, the substance is a liquid.

4.

A substance has a melting point of 20 °C and a boiling point of 80 °C. What state is it in at 100 °C?

At 100 °C, the substance is a gas.

5.

What physical state would a substance be in at a temperature below its melting point?

Below its melting point, a substance is a solid.

6.

What physical state would a substance be in at a temperature above its boiling point?

Above its boiling point, a substance is a gas.

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?

In chemistry, a pure substance contains only one substance and has no other substances mixed with it.

2.

How does the scientific meaning of pure differ from its everyday meaning?

The everyday meaning of pure can mean clean, uncontaminated or not adulterated, rather than chemically containing only one substance.

3.

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

A pure substance contains only one element or compound, whereas a mixture contains two or more substances.

4.

Why is air classified as a mixture rather than a pure substance?

Air contains several gases, so it is a mixture rather than a pure substance.

5.

Why can a mixture contain more than one substance while a pure substance contains only one substance?

A mixture contains more than one substance, whereas a pure substance contains only one substance.

6.

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

The components of a mixture retain their individual chemical identities and 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 behaviour of a pure substance differ from that of a mixture?

A pure substance generally melts sharply at a specific temperature, whereas a mixture melts over a range.

2.

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

A sharp melting point suggests that a substance is pure.

3.

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

Melting over a range of temperatures suggests that a substance is a mixture or impure.

4.

A substance melts sharply at 80 °C. What does this suggest about its purity?

A substance that melts sharply at 80 °C is likely to be pure.

5.

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

A substance that melts from 70 °C to 76 °C 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?

Melting point data can distinguish purity because pure substances have sharp melting points while mixtures generally melt 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?

Simple distillation separates a soluble solid from a solution and can be used to obtain the solvent.

2.

What type of mixture can be separated using fractional distillation?

Fractional distillation separates a mixture of miscible liquids with different boiling points.

3.

What type of mixture can be separated using filtration?

Filtration separates an insoluble solid from a liquid.

4.

What type of mixture can be separated using crystallisation?

Crystallisation separates a soluble solid from a solution.

5.

What type of mixture can be separated using paper chromatography?

Paper chromatography separates a mixture of soluble substances, such as dyes.

6.

What property of the components of a mixture determines which separation technique is appropriate?

The relevant property is the difference in physical properties such as solubility, particle size, boiling point or attraction to the stationary phase.

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

1.

Which separation technique would be appropriate for separating an insoluble solid from a liquid, and why?

Use filtration to separate an insoluble solid from a liquid because the solid particles are too large to pass through the filter.

2.

Which separation technique would be appropriate for obtaining a solvent from a solution containing a dissolved solid, and why?

Use simple distillation to obtain the solvent from a solution because the solvent has a lower boiling point and can be vaporised and condensed.

3.

Which separation technique would be appropriate for separating two miscible liquids with different boiling points, and why?

Use fractional distillation to separate miscible liquids with different boiling points.

4.

Which separation technique would be appropriate for separating several soluble substances with different affinities for a stationary phase, and why?

Use paper chromatography to separate soluble substances with different attractions to the stationary and mobile phases.

5.

How would you choose an appropriate technique to separate the components of an unknown mixture?

Identify the physical properties of each component, then select a technique that exploits a difference in those properties.

6.

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

The components' physical properties determine whether they can be separated by filtration, distillation, crystallisation or chromatography.

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?

The solvent dissolves the substances and carries them through the chromatography paper.

2.

What is meant by the mobile phase in paper chromatography?

The mobile phase is the solvent that moves through the paper.

3.

What is meant by the stationary phase in paper chromatography?

The stationary phase is the chromatography paper.

4.

Why do different soluble substances move different distances during paper chromatography?

Different substances move at different rates because they have different attractions to the stationary and mobile phases.

5.

Why must the substances being separated be soluble in the solvent used for paper chromatography?

The substances must dissolve in the solvent so that they can move with the mobile phase.

6.

How does the movement of the solvent through the paper cause a mixture to separate?

The solvent travels through the paper and carries the dissolved substances different distances, separating the mixture.

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?

A pure substance produces one spot, whereas an impure substance generally produces more than one spot.

2.

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

Compare the position of an unknown spot with the position of a known substance under the same conditions to help identify it.

3.

What does a single spot on a chromatogram indicate about a substance?

A single spot indicates that the sample may contain one substance.

4.

What does more than one spot on a chromatogram indicate about a substance?

More than one spot indicates that the sample contains more than one soluble substance.

5.

How is the Rf value of a substance calculated from a paper chromatogram?

Rf = distance travelled by substance ÷ distance travelled by solvent front.

6.

How can an Rf value be used to identify an unknown substance?

An unknown substance can be identified by comparing its Rf value with the Rf value of a known substance under the same conditions.

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

1.

How can simple distillation be used to separate the solvent from an ink sample?

Simple distillation can separate and collect the solvent from the ink because the solvent is vaporised and then condensed.

2.

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

Paper chromatography separates the different soluble dyes in the ink so their composition can be investigated.

3.

Why must the ink sample be placed near the bottom of the chromatography paper?

The ink is placed near the bottom of the paper so the solvent can carry the dyes upwards through the paper.

4.

Why should the solvent level be below the ink spot when carrying out paper chromatography?

The solvent level must be below the ink spot so the ink does not dissolve directly into the solvent reservoir.

5.

How can the chromatogram be used to determine whether an ink contains more than one dye?

Multiple spots on the chromatogram show that the ink contains more than one dye.

6.

How can the dyes in an unknown ink be identified using known substances and their chromatogram results?

Compare the positions or Rf values of the unknown ink's dyes with known dyes to identify the components.

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 and groundwater potable?

Waste water and groundwater can be treated using sedimentation, filtration and chlorination to make them potable.

2.

What is the purpose of sedimentation when treating waste water or groundwater?

Sedimentation allows suspended particles to settle out of the water.

3.

What is the purpose of filtration when treating waste water or groundwater?

Filtration removes remaining insoluble particles and suspended solids.

4.

Why is chlorination used when making water potable?

Chlorination kills microorganisms and makes the water safer to drink.

5.

How can sea water be made potable using distillation?

Sea water can be made potable by distillation, which separates water from dissolved salts.

6.

Why must water used in chemical analysis contain no dissolved salts?

Water used in chemical analysis must contain no dissolved salts because dissolved ions could interfere with or alter the results of the analysis.

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

1.

What ion do acids produce when they dissolve in water?

Hydrogen ions, H⁺.

2.

What ion do alkalis produce when they dissolve in water?

Hydroxide ions, OH⁻.

3.

Which ion is responsible for the acidic properties of an aqueous acid?

Hydrogen ions, H⁺.

4.

Which ion is responsible for the alkaline properties of an aqueous alkali?

Hydroxide ions, OH⁻.

5.

What is the formula of the hydrogen ion produced by acids in aqueous solution?

H⁺.

6.

What is the formula of the hydroxide ion produced by alkalis in aqueous solution?

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?

pH 7.

2.

What pH values do acidic solutions have?

pH values below 7.

3.

What pH values do alkaline solutions have?

pH values above 7.

4.

Is a solution with a pH of 4 acidic, neutral or alkaline?

Acidic.

5.

Is a solution with a pH of 11 acidic, neutral or alkaline?

Alkaline.

6.

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

The solution with pH 2.

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

1.

What colour does blue litmus turn in an acidic solution?

Red.

2.

What colour does red litmus turn 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 hydrogen ion concentration affect the pH of an acidic solution?

The pH decreases.

2.

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

The pH increases.

3.

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

The pH increases.

4.

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

The pH decreases.

5.

Which solution has the lower pH: one with a hydrogen ion concentration of 0.1 mol dm⁻³ or one with 0.001 mol dm⁻³?

0.1 mol dm⁻³.

6.

Which alkaline solution has the higher pH: one with a hydroxide ion concentration of 0.01 mol dm⁻³ or one with 0.001 mol dm⁻³?

0.01 mol dm⁻³.

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.

What happens to the pH when the hydrogen ion concentration increases by a factor of 10?

The pH decreases by 1.

2.

What happens to the pH when the hydrogen ion concentration decreases by a factor of 10?

The pH increases by 1.

3.

A solution has a pH of 5. What is its pH if its hydrogen ion concentration increases by a factor of 10?

pH 4.

4.

A solution has a pH of 3. What is its pH if its hydrogen ion concentration increases by a factor of 100?

pH 1.

5.

A solution has a pH of 6. What is its pH if its hydrogen ion concentration decreases by a factor of 100?

pH 8.

6.

By what factor has the hydrogen ion concentration increased if the pH decreases from 6 to 3?

A factor of 1000.

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.

What apparatus can be used to measure the pH of hydrochloric acid during the reaction with calcium hydroxide?

A pH meter/probe or universal indicator.

2.

Why should the volume and concentration of hydrochloric acid be kept constant when investigating the effect of adding calcium hydroxide?

To ensure the amount and concentration of acid are controlled variables, making the investigation a fair test.

3.

What should be measured after each addition of powdered calcium hydroxide to hydrochloric acid?

The pH of the solution.

4.

Why should powdered calcium hydroxide be added in small measured amounts during this investigation?

To measure the effect of known amounts of calcium hydroxide and obtain accurate results.

5.

What happens to the pH of hydrochloric acid as calcium hydroxide is added?

The pH increases.

6.

What graph could be plotted to show how the pH changes as calcium hydroxide is added?

A graph of pH against mass/amount of calcium hydroxide 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?

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

2.

What is meant by a concentrated solution?

A solution containing a relatively 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 water.

5.

How does adding water affect the concentration of a solution?

It decreases the concentration.

6.

Do the terms dilute and concentrated describe the strength of an acid?

No. They describe concentration, not acid strength.

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?

An acid that completely dissociates into ions in aqueous solution.

2.

What is meant by a weak acid?

An acid that only partially dissociates into ions in aqueous solution.

3.

What happens to acid molecules in water when a strong acid dissociates?

Nearly all acid molecules dissociate into ions.

4.

What happens to acid molecules in water when a weak acid dissociates?

Only some acid molecules dissociate into ions.

5.

Which has a greater degree of dissociation into ions: a strong acid or a weak acid?

A strong acid.

6.

Can a dilute acid be strong, and can a concentrated acid be weak?

Yes. Strength and concentration are different properties.

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?

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

2.

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

A salt and water.

3.

What type of substance reacts with an acid to form only a salt and water?

A base.

4.

Is copper oxide a base if it reacts with hydrochloric acid to form copper chloride and water?

Yes.

5.

Is sodium hydroxide a base if it reacts with hydrochloric acid to form sodium chloride and water?

Yes.

6.

What product other than water is formed when a base reacts with an acid?

A salt.

3.10 Recall that alkalis are soluble bases

1.

What is an alkali?

A soluble base.

2.

What property must a base have to be classified as an alkali?

It must be soluble in water.

3.

Are all bases alkalis?

No.

4.

Why is sodium hydroxide classified as an alkali?

Sodium hydroxide is a soluble base.

5.

Why is copper oxide not classified as an alkali?

Copper oxide is insoluble in water.

6.

What happens when an alkali dissolves in water?

It dissolves and produces OH⁻ ions in solution.

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 products when an acid reacts with a metal?

Salt and hydrogen.

2.

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

Salt and water.

3.

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

Salt and water.

4.

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

Salt, water and carbon dioxide.

5.

What gas is produced when an acid reacts with a metal?

Hydrogen.

6.

What gas is produced when an acid reacts with a metal carbonate?

Carbon dioxide.

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

1.

How is hydrogen tested for in the laboratory?

Place a lighted splint at the mouth of the container holding the gas.

2.

What observation confirms that a gas is hydrogen?

A squeaky pop is heard.

3.

How is carbon dioxide tested for using limewater?

Bubble the gas through limewater.

4.

What observation confirms that a gas is carbon dioxide using limewater?

The limewater turns milky/cloudy.

5.

What sound is produced when a lighted splint is placed in hydrogen?

A squeaky pop.

6.

What happens to limewater when carbon dioxide is bubbled through it?

It turns milky/cloudy.

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.

Which two types of substances react in a neutralisation reaction?

An acid and a base.

3.

What products are formed in a neutralisation reaction between an acid and a base?

A salt and water.

4.

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

Yes.

5.

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

A neutralisation reaction.

6.

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

The acidic and alkaline properties are removed as the acid and base react.

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 react together during an acid-alkali neutralisation?

H⁺ and OH⁻ ions.

2.

Where do the hydrogen ions in an acid-alkali neutralisation come from?

The acid.

3.

Where do the hydroxide ions in an acid-alkali neutralisation come from?

The alkali.

4.

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

Water.

5.

Write the ionic equation for an acid-alkali neutralisation.

H⁺ + OH⁻ → H₂O.

6.

Why does an acid-alkali neutralisation remove the acidic and alkaline properties of the solutions?

H⁺ ions and OH⁻ ions react to form water.

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.

How is excess insoluble reactant removed after reacting it with an acid?

By filtration.

3.

Why can excess insoluble reactant be removed by filtration?

The insoluble reactant remains as a solid and can be separated from the solution.

4.

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

Salt and water.

5.

Why must the excess insoluble reactant be removed before crystallising the soluble salt?

To prevent the insoluble reactant contaminating the salt crystals.

6.

Why is an insoluble reactant suitable for preparing a soluble salt without using titration?

The excess insoluble reactant can be removed by filtration.

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 is titration used when preparing a soluble salt from an acid and a soluble reactant?

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

2.

Why cannot an excess soluble reactant simply be removed by filtration?

It would remain dissolved in the solution.

3.

What does titration determine when preparing a soluble salt from an acid and a soluble reactant?

The volumes of reactants needed to exactly neutralise each other.

4.

Why are the acid and soluble reactant mixed in the correct proportions after titration?

To ensure neither reactant is left in excess.

5.

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

Salt and water.

6.

Why is titration necessary when both reactants used to prepare a soluble salt are soluble?

Both reactants are soluble, so filtration cannot remove excess reactant.

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 copper sulfate crystals?

To ensure all the sulfuric acid reacts.

3.

How is excess copper oxide removed from the copper sulfate solution?

By filtration.

4.

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

To gently heat the solution and evaporate some water without overheating it.

5.

How are hydrated copper sulfate crystals obtained from the copper sulfate solution?

Concentrate the solution, then allow it to cool so crystals form.

6.

Why must the copper sulfate crystals be dried after they have formed?

To remove surface water and obtain a dry sample.

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 volume of acid accurately during a titration?

Burette.

2.

What piece of apparatus is used to measure a fixed volume of alkali accurately during a titration?

Pipette.

3.

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

To show when the endpoint/neutralisation has been reached.

4.

What is done when the indicator shows that neutralisation has been reached?

Record the titre and use the volume to determine the exact proportions of reactants.

5.

Why is a second titration carried out without indicator when preparing a pure salt?

To avoid contaminating the final salt with indicator.

6.

How is a pure, dry salt obtained from the neutralised solution after titration?

Evaporate some water from the solution, allow the salt to crystallise, then filter and dry the crystals.

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.

Are all common sodium, potassium and ammonium salts soluble in water?

Yes.

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 soluble in water?

Sodium, potassium and ammonium carbonates.

6.

Which common hydroxides are soluble in water?

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.

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

Identify the possible products and use the solubility rules to determine whether either product is insoluble.

2.

What is a precipitate?

An insoluble solid formed from solutions during a reaction.

3.

What happens when two solutions are mixed and an insoluble product forms?

A precipitate forms.

4.

When sodium chloride solution is mixed with silver nitrate solution, what precipitate is formed?

Silver chloride, AgCl.

5.

When barium chloride solution is mixed with sodium sulfate solution, what precipitate is formed?

Barium sulfate, BaSO₄.

6.

When potassium nitrate solution is mixed with sodium chloride solution, why is no precipitate formed?

The possible products are soluble, so no precipitate forms.

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

1.

What type of reaction can be used to prepare an insoluble salt?

A precipitation reaction between two soluble solutions.

2.

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

To form an insoluble salt as a precipitate.

3.

How is the insoluble salt separated from the reaction mixture?

By filtration.

4.

Why is the insoluble salt washed with distilled water?

To remove soluble impurities.

5.

How is an insoluble salt dried after it has been filtered and washed?

Leave it to dry or dry it using suitable drying equipment.

6.

Why must the insoluble salt be washed before it is dried?

To remove soluble impurities remaining in the precipitate.

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 compounds 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 it is dissolved in water?

The ions are free to move through the solution.

5.

Why does solid sodium chloride not conduct electricity?

Its ions are held in fixed positions in the lattice.

6.

What must be present and free to move for an ionic substance to conduct electricity?

Free-moving charged ions.

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 energy is supplied during electrolysis?

Electrical energy.

3.

What type of electrical supply is used for electrolysis?

Direct current (DC).

4.

What happens to an electrolyte during electrolysis?

It decomposes into simpler substances.

5.

Why does electrolysis cause an electrolyte to decompose?

Electrical energy causes ions to move to the electrodes, where they undergo chemical reactions.

6.

What is the role of a direct current supply in electrolysis?

It provides the electrical energy and maintains a constant direction of ion movement.

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.

During electrolysis, which electrode do positively charged cations migrate towards?

The negatively charged cathode.

2.

During electrolysis, which electrode do negatively charged anions migrate towards?

The positively charged anode.

3.

Why do cations migrate towards the cathode during electrolysis?

Opposite charges attract.

4.

Why do anions migrate towards the anode during electrolysis?

Opposite charges attract.

5.

What is the charge of a cation and which electrode attracts it during electrolysis?

Positive; the cathode.

6.

What is the charge of an anion and which electrode attracts it during electrolysis?

Negative; the anode.

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 cathode and anode when aqueous copper chloride is electrolysed using inert electrodes?

Cathode: copper. Anode: chlorine.

2.

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

Cathode: hydrogen. Anode: chlorine.

3.

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

Cathode: hydrogen. Anode: oxygen.

4.

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

Cathode: hydrogen. Anode: oxygen.

5.

What products are formed at the cathode and anode when molten lead bromide is electrolysed?

Cathode: lead. Anode: bromine.

6.

Why is hydrogen produced at the cathode instead of sodium during the electrolysis of aqueous sodium chloride?

Hydrogen is preferentially discharged from water because sodium is more reactive than hydrogen.

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

1.

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

The metal and non-metal elements.

2.

How can the products of electrolysis of a molten metal halide be predicted?

The metal forms at the cathode and the halogen forms at the anode.

3.

What product forms from the metal cation at the cathode during electrolysis of a molten ionic compound?

The metal.

4.

What product forms from the non-metal anion at the anode during electrolysis of a molten ionic compound?

The non-metal.

5.

What products are formed when molten sodium chloride is electrolysed?

Sodium and chlorine.

6.

What products are formed when molten magnesium bromide is electrolysed?

Magnesium and bromine.

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

1.

What is a half equation?

An equation showing the electron transfer in one half of a redox reaction.

2.

How can a half equation show what happens to an ion during electrolysis?

It shows the ions, electrons and products involved.

3.

What type of reaction occurs at the cathode during electrolysis, and how is it represented in a half equation?

Reduction; electrons are gained.

4.

What type of reaction occurs at the anode during electrolysis, and how is it represented in a half equation?

Oxidation; electrons are lost.

5.

What half equation represents the reduction of Cu²⁺ ions to copper atoms at the cathode?

Cu²⁺ + 2e⁻ → Cu.

6.

What half equation represents the oxidation of chloride ions to chlorine gas at the anode?

2Cl⁻ → Cl₂ + 2e⁻.

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 electrons when a substance is oxidised?

They lose electrons.

4.

What happens to electrons when a substance is reduced?

They gain electrons.

5.

How can the mnemonic OIL RIG be used to remember oxidation and reduction?

Oxidation Is Loss; Reduction Is Gain.

6.

Is 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 happens to ions that are reduced at the cathode?

They gain electrons.

4.

What happens to ions that are oxidised at the anode?

They lose electrons.

5.

Which electrode is associated with electron gain during electrolysis?

Cathode.

6.

Which electrode is associated with electron loss during electrolysis?

Anode.

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 to Cu²⁺ ions at the cathode when copper sulfate solution is electrolysed using copper electrodes?

They gain electrons and are deposited as copper atoms: Cu²⁺ + 2e⁻ → Cu.

2.

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

Copper atoms lose electrons and form Cu²⁺ ions: Cu → Cu²⁺ + 2e⁻.

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 sulfate solution remain approximately unchanged 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 copper?

Impure copper is used as the anode and pure copper is deposited at the cathode. Insoluble impurities fall to the bottom as sludge.

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

1.

How can the electrolysis of copper sulfate solution be investigated using inert electrodes?

Place copper sulfate solution in a suitable container, insert inert electrodes, connect them to a DC supply and observe the electrodes during electrolysis.

2.

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

Copper is deposited at the cathode and bubbles of oxygen form at the anode.

3.

How can the electrolysis of copper sulfate solution be investigated using copper electrodes?

Repeat using copper electrodes connected to a DC supply and observe and measure changes at the electrodes.

4.

What happens to the mass of each copper electrode during electrolysis of copper sulfate solution?

The cathode gains mass and the anode loses mass.

5.

What measurements could be taken to compare the changes occurring at copper electrodes during electrolysis?

Measure the mass of each electrode before and after electrolysis and compare the changes.

6.

What safety precautions should be taken when investigating the electrolysis of copper sulfate solution?

Wear safety goggles, use a low-voltage DC supply, avoid contact with the solution, and work in suitable ventilation when gases are produced.

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

1.

How can the relative reactivity of metals be deduced from their reactions with water?

A more reactive metal reacts more readily with water, often producing hydrogen.

2.

How can the relative reactivity of metals be deduced from their reactions with dilute acids?

A more reactive metal reacts more vigorously with dilute acids, usually producing hydrogen.

3.

How can the relative reactivity of two metals be compared using a salt solution?

Place each metal in the other metal's salt solution and observe whether a displacement reaction occurs.

4.

What does it indicate about the reactivity of a metal if it displaces another metal from its salt solution?

The metal doing the displacing is more reactive.

5.

If magnesium displaces copper from copper sulfate solution, which metal is more reactive?

Magnesium is more reactive than copper.

6.

If copper does not react with dilute hydrochloric acid but magnesium does, what does this show about their relative reactivities?

Magnesium is more reactive than copper.

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

1.

What is a displacement reaction between a metal and a metal compound?

A reaction in which a more reactive metal displaces a less reactive metal from its compound.

2.

Why is a metal displacement reaction a redox reaction?

Electrons are transferred between the metals, so oxidation and reduction occur.

3.

What happens to the atoms of the more reactive metal during a displacement reaction?

They lose electrons and are oxidised.

4.

What happens to the ions of the less reactive metal during a displacement reaction?

They gain electrons and are reduced.

5.

In the reaction between zinc and copper sulfate, which substance is oxidised and which is reduced?

Zinc is oxidised; copper ions are reduced.

6.

In a displacement reaction, how can electron transfer be used to identify oxidation and reduction?

Loss of electrons is oxidation; gain of electrons is reduction.

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 order of the metals in the reactivity series from most to least reactive?

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

2.

How does the reactivity series show the relative tendency of metal atoms to form cations?

Metals higher in the series form cations more readily than metals lower in the series.

3.

Why do metals higher in the reactivity series tend to form cations more readily?

They lose electrons more readily.

4.

How does the reaction of a metal with water provide evidence of its position in the reactivity series?

More reactive metals react more readily with water.

5.

How does the reaction of a metal with dilute acid provide evidence of its position in the reactivity series?

More reactive metals react more readily and vigorously with dilute acids.

6.

Where are carbon and hydrogen placed in the reactivity series, and why are they useful reference points?

Carbon and hydrogen are reference points used to compare the reactivity of metals and determine extraction methods.

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 before they are extracted?

In ores in the Earth's crust.

2.

What is an ore?

A naturally occurring rock containing enough metal compound for the metal to be extracted economically.

3.

Why are most metals found as compounds in ores rather than as uncombined elements?

Most metals are reactive and react with other elements to form compounds.

4.

Which type of metals can be found naturally in the Earth's crust as uncombined elements?

Unreactive metals.

5.

Why can unreactive metals such as gold be found as uncombined elements?

They do not readily react with other substances.

6.

How does a metal's reactivity affect whether it is likely to be found as an ore or as an uncombined element?

More reactive metals are generally found as compounds in ores, while very unreactive metals may occur as uncombined elements.

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 gains oxygen?

It is oxidised.

4.

What happens to a substance when it loses oxygen?

It is reduced.

5.

In the reaction CuO + H₂ → Cu + H₂O, which substance is oxidised?

H₂ is oxidised.

6.

In the reaction CuO + H₂ → Cu + H₂O, which substance is reduced?

CuO is reduced.

4.6 Recall that the extraction of metals involves reduction of ores

1.

Why does extracting a metal from its ore involve reduction?

The metal compound must be reduced to obtain the metal.

2.

What must happen to a metal compound in an ore for the metal to be extracted?

The metal ions must be converted into metal atoms.

3.

How is oxygen removed from a metal oxide during metal extraction?

By reacting the metal oxide with a reducing agent such as carbon.

4.

Why is reduction of a metal compound necessary to obtain the metal itself?

To remove oxygen from the metal compound and leave the metal.

5.

What happens to the metal ions in a metal ore during extraction?

They gain electrons and become metal atoms.

6.

How does the reduction of a metal oxide produce the metal?

Oxygen is removed from the metal oxide, leaving 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.

How does a metal's position in the reactivity series determine how it can be extracted from its ore?

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

2.

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

Carbon is more reactive than these metals and can remove oxygen from their oxides.

3.

Why can aluminium not be extracted from aluminium oxide by heating with carbon?

Aluminium is more reactive than carbon, so carbon cannot remove oxygen from aluminium oxide.

4.

Why is electrolysis used to extract aluminium from its ore?

Aluminium oxide must be electrolysed because aluminium is more reactive than carbon.

5.

Why is the cost of extraction an important factor when choosing a method for extracting a metal?

Extraction methods require different amounts of energy and resources, affecting the overall cost.

6.

Why is extracting a very reactive metal by electrolysis generally more expensive than extracting iron by heating its ore with carbon?

Electrolysis requires large amounts of electrical energy, whereas carbon reduction requires less expensive energy and reducing agents.

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

1.

What is bacterial extraction of metals?

Using bacteria to convert metal compounds into soluble forms that can be extracted.

2.

How can bacteria help extract metals from low-grade ores?

Bacteria can carry out reactions that make metal compounds soluble, allowing the metal to be recovered.

3.

What is phytoextraction?

Using plants to absorb metal compounds from contaminated soil.

4.

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

Plants absorb metal ions through their roots and concentrate them in their tissues, which can then be harvested.

5.

What are the advantages of using biological methods to extract metals compared with traditional extraction methods?

They can use less energy and produce less environmental damage than traditional mining and extraction.

6.

What are the disadvantages of bacterial extraction and phytoextraction?

They can be slow, may produce lower yields, and depend on suitable environmental conditions.

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

1.

How is a metal's position in the reactivity series related to its resistance to oxidation?

Metals higher in the reactivity series are less resistant to oxidation.

2.

Why are metals high in the reactivity series more easily oxidised?

They lose electrons readily and therefore oxidise easily.

3.

Why are metals low in the reactivity series more resistant to oxidation?

They lose electrons less readily.

4.

Which is more resistant to oxidation, magnesium or copper, and why?

Copper, because it is lower in the reactivity series and is more resistant to oxidation.

5.

Why does gold have a high resistance to oxidation?

Gold is very unreactive and does not readily react with oxygen.

6.

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

A greater tendency to form cations means a greater tendency to lose electrons and become 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.

What are the environmental advantages of recycling metals?

It reduces mining, habitat destruction, waste and pollution.

2.

How does recycling metals help conserve valuable raw materials?

It reduces the amount of new ore that needs to be extracted.

3.

How can recycling metals reduce the environmental damage caused by mining?

Less mining is required, reducing habitat destruction and pollution.

4.

How can recycling metals reduce energy use compared with extracting metals from ores?

Recycling often requires less energy than extracting metals from ores.

5.

What are the economic advantages of recycling metals?

It saves raw materials, can reduce energy costs and provides useful materials for reuse.

6.

What factors could make recycling metals less economically or environmentally beneficial?

Collection, sorting and processing can be expensive and recycling may require significant energy.

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 impacts of a product throughout its life.

2.

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

Mining, extraction, habitat destruction, energy use and pollution.

3.

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

Energy use, emissions, waste and pollution.

4.

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

Energy consumption, emissions and other environmental effects during use.

5.

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

Waste produced, recycling, reuse and disposal impacts.

6.

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

Obtaining raw materials, manufacturing, using the product and disposing of it.

4.12 Evaluate data from a life cycle assessment of a product

1.

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

Compare the environmental impacts at each stage and the overall impact.

2.

What factors should be considered when evaluating life-cycle assessment data?

The environmental impacts and resource use at all stages of the product's life.

3.

Why should the environmental impact of obtaining raw materials be considered when evaluating a product?

Obtaining raw materials can involve mining, energy use, habitat destruction and pollution.

4.

Why should the environmental impact of manufacturing, using and disposing of a product be considered together?

The overall environmental impact depends on all stages of the product's life.

5.

How can life-cycle assessment data be used to identify which product has the lower overall environmental impact?

Compare the total environmental impacts shown by the data.

6.

Why might a product with a lower environmental impact at one stage of its life still have a greater overall environmental impact?

A greater impact at another stage may outweigh the lower impact at one stage.

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 in which the products can react to reform the original reactants.

2.

What symbol is used to represent a reversible reaction?

⇌.

3.

What does the symbol ⇌ indicate in a chemical equation?

The reaction can proceed in both directions.

4.

What happens in the forward reaction of a reversible reaction?

Reactants form products.

5.

What happens in the reverse reaction of a reversible reaction?

Products react to form reactants.

6.

How can changing reaction conditions alter the direction in which a reversible reaction proceeds?

Changing temperature, pressure or concentration can alter the position/direction of equilibrium.

4.14 Explain what is meant by dynamic equilibrium

1.

What is meant by dynamic equilibrium?

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

2.

What happens to the forward reaction at dynamic equilibrium?

It continues to occur.

3.

What happens to the reverse reaction at dynamic equilibrium?

It continues to occur.

4.

Why do the concentrations of reactants and products remain constant at dynamic equilibrium?

The reactants and products are being formed and used at equal rates.

5.

Why does a reversible reaction not stop when dynamic equilibrium is reached?

Both reactions continue, but at equal rates.

6.

What must be true about the rates of the forward and reverse reactions at dynamic equilibrium?

The forward and reverse reaction rates are 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 two elements react to form 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 reaction between nitrogen and hydrogen to form ammonia?

N₂(g) + 3H₂(g) ⇌ 2NH₃(g).

5.

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

The reaction is reversible, so ammonia can react to reform nitrogen and hydrogen.

6.

What happens to the concentrations of nitrogen, hydrogen and ammonia when dynamic equilibrium is reached?

Their concentrations remain constant because the forward and reverse reactions occur at equal rates.

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.

What 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 without being used up.

6.

Why is the Haber process carried out at 450 °C and 200 atmospheres rather than under extreme conditions?

These conditions provide a compromise between reaction rate, ammonia yield and industrial cost.

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

1.

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

It shifts to the left, towards nitrogen and hydrogen.

2.

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

It shifts to the right, towards ammonia.

3.

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

It shifts to the right, towards ammonia.

4.

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

It shifts to the left, towards nitrogen and hydrogen.

5.

For a reversible reaction at dynamic equilibrium, how does increasing the concentration of a reactant affect the position of equilibrium?

It shifts towards the products, using up some of the added reactant.

6.

For a reversible reaction at dynamic equilibrium, how does decreasing the concentration of a product affect the position of equilibrium?

It shifts towards the products, replacing some of the product that was removed.

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.

Which group of the periodic table contains the alkali metals?

Group 1.

2.

Which group of the periodic table contains the halogens?

Group 7.

3.

Which group of the periodic table contains the noble gases?

Group 0.

4.

Why can elements in the same group of the periodic table be classified as a particular group of elements?

Elements in the same group have similar chemical properties.

5.

What does an element's position in the periodic table tell you about which group it belongs to?

It tells you which group the element belongs to and therefore its classification.

6.

Why are lithium, sodium and potassium classified as alkali metals?

They are all elements in Group 1.

6.2 Recall that alkali metals a are soft b have relatively low melting points

1.

What is the physical property of alkali metals that allows them to be cut relatively easily?

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 are alkali metals described as soft metals?

They can be cut relatively easily.

4.

What happens to the physical state of an alkali metal when it is heated above its relatively low melting point?

It melts and changes from a solid to a liquid.

5.

Which two physical properties are characteristic of alkali metals?

They are soft and have relatively low melting points.

6.

How do the softness and melting points of alkali metals distinguish them from many other metals?

They are softer and have lower melting points than many other metals.

6.3 Describe the reactions of lithium, sodium and potassium with water

1.

What happens when lithium is added to water?

Lithium floats, moves around on the surface and fizzes as it reacts, forming lithium hydroxide and hydrogen.

2.

What happens when sodium is added to water?

Sodium melts into a ball, moves rapidly on the surface and fizzes, forming sodium hydroxide and hydrogen.

3.

What happens when potassium is added to water?

Potassium reacts very vigorously, often igniting with a lilac flame, forming potassium hydroxide and hydrogen.

4.

What products are formed when an alkali metal reacts with water?

A metal hydroxide and hydrogen.

5.

What observations would you make when sodium reacts with water?

It floats, melts into a ball, moves around, fizzes and gradually disappears.

6.

What observations would you make when potassium reacts with water?

It reacts very vigorously, moves rapidly, fizzes and usually burns with a lilac flame.

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 pattern of lithium, sodium and potassium be used to predict the reactivity of rubidium?

Rubidium would be more reactive than potassium.

5.

How can the reactivity pattern of Group 1 metals be used to predict the reactivity of caesium?

Caesium would be even more reactive than rubidium.

6.

What happens to the speed and violence of the reaction with water as Group 1 metals become more reactive?

The reaction becomes faster and more violent.

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.

2.

What happens to the outer electron of a Group 1 atom as the atomic number increases down the group?

It is further from the nucleus.

3.

Why is the outer electron easier to remove from potassium than from lithium?

Potassium has a larger atomic radius and greater electron shielding, so its outer electron is less strongly attracted to the nucleus.

4.

How does increased atomic radius down Group 1 affect the attraction between the nucleus and the outer electron?

The attraction between the nucleus and outer electron decreases.

5.

Why does increased electron shielding down Group 1 make the outer electron easier to remove?

Shielding reduces the attraction between the nucleus and the outer electron.

6.

How does the electronic configuration of Group 1 metals explain their increasing reactivity down the group?

Group 1 metals have one outer electron, which becomes easier to lose down the group, increasing reactivity.

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?

Pale 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.

What colour is chlorine at room temperature?

Pale green.

5.

What colour is bromine at room temperature?

Red-brown.

6.

What colour is iodine at room temperature?

Grey-black.

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 from chlorine to bromine to iodine?

They change from gas to liquid to solid.

2.

How does the colour of the halogens change down Group 7?

The colour becomes darker.

3.

What happens to the melting and boiling points of the halogens down Group 7?

Their melting and boiling points increase.

4.

How can the physical property pattern of chlorine, bromine and iodine be used to predict the physical state of astatine at room temperature?

Astatine is predicted to be a solid.

5.

How can the physical property pattern of chlorine, bromine and iodine be used to predict the colour of another halogen?

The halogen would be predicted to become darker down the group.

6.

What general trend occurs in the physical properties of the halogens as you move down Group 7?

Melting point, boiling point and density generally increase, while the physical state changes from gas to liquid to solid.

6.8 Describe the chemical test for chlorine

1.

How can chlorine gas be chemically tested?

Expose damp litmus paper to the gas.

2.

What happens when damp litmus paper is exposed to chlorine gas?

It is bleached.

3.

What colour change does damp blue litmus paper undergo when exposed to chlorine?

Blue litmus turns red and is then bleached white.

4.

Why does chlorine eventually bleach damp litmus paper?

Chlorine reacts with water to form substances that bleach the dye.

5.

Why must the litmus paper be damp when testing for chlorine?

Chlorine must dissolve in water for the bleaching reaction to occur.

6.

What observation confirms the presence of chlorine gas using damp litmus paper?

Damp litmus paper is bleached.

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 compounds are formed when halogens react with metals?

Metal halides.

2.

What happens when chlorine reacts with a metal?

It reacts with the metal to form a metal chloride.

3.

What happens when bromine reacts with a metal?

It reacts with the metal to form a metal bromide.

4.

What happens when iodine reacts with a metal?

It reacts with the metal to form a metal iodide.

5.

What type of ion does a halogen form when it reacts with a metal?

A halide ion with a −1 charge.

6.

How can the reactions of chlorine, bromine and iodine with metals be used to predict the reactions of other halogens?

Other halogens are predicted to 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 compounds are formed when chlorine, bromine and iodine react with hydrogen?

Hydrogen chloride, hydrogen bromide and hydrogen iodide.

2.

What are hydrogen chloride, hydrogen bromide and hydrogen iodide collectively known as?

Hydrogen halides.

3.

What happens when hydrogen chloride dissolves in water?

It dissolves and forms an acidic solution.

4.

Why do hydrogen halides produce acidic solutions when they dissolve in water?

They produce hydrogen ions in water.

5.

What acidic solution is formed when hydrogen bromide dissolves in water?

Hydrobromic acid.

6.

How can the reactions of chlorine, bromine and iodine with hydrogen be used to predict the reaction of another halogen?

Another halogen is predicted to react with hydrogen to form a hydrogen halide that dissolves in water to form an acidic solution.

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 compare?

Reactivity decreases down the group: chlorine > bromine > iodine.

2.

What happens when chlorine is added to a solution containing bromide ions?

Chlorine displaces bromide ions, producing bromine.

3.

What happens when chlorine is added to a solution containing iodide ions?

Chlorine displaces iodide ions, producing iodine.

4.

What happens when bromine is added to a solution containing iodide ions?

Bromine displaces iodide ions, producing iodine.

5.

Why can chlorine displace bromide ions and iodide ions from aqueous solutions?

Chlorine is more reactive and gains electrons more readily than bromine and iodine.

6.

Using the Group 7 reactivity pattern, how would astatine be expected to react with chloride, bromide and iodide ions?

Astatine would be less reactive than iodine and would not displace chloride, bromide or iodide ions.

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 the displacement of a halide ion by a more reactive halogen a redox reaction?

Electrons are transferred, so oxidation and reduction both occur.

2.

What happens to the halide ions when they are displaced by a more reactive halogen?

They lose electrons and are oxidised.

3.

What happens to the halogen molecules when they displace halide ions?

They gain electrons and are reduced.

4.

In the reaction Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, which substance is oxidised?

Br⁻ is oxidised.

5.

In the reaction Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, which substance is reduced?

Cl₂ is reduced.

6.

In a halogen displacement reaction, how can electron gain and loss be used to identify oxidation and reduction?

Loss of electrons is oxidation; gain of electrons is reduction.

6.13 Explain the relative reactivity of the halogens in terms of electronic configurations

1.

Why does the reactivity of the halogens decrease down Group 7?

The outer shell is further from the nucleus and electron shielding increases, making it harder to attract an incoming electron.

2.

What happens to the atomic radius of halogen atoms as you move down Group 7?

It increases.

3.

How does electron shielding change down Group 7?

It increases.

4.

Why is a chlorine atom better able to attract an electron than an iodine atom?

Chlorine has a smaller atomic radius and less electron shielding, so its nucleus attracts an incoming electron more strongly.

5.

How does the increasing distance between the nucleus and outer shell affect the attraction for an incoming electron?

The attraction decreases.

6.

How does electronic configuration explain why chlorine is more reactive than bromine and iodine?

Chlorine attracts an incoming electron more strongly because its outer shell is closer to the nucleus and there is less shielding.

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 gas atoms?

Their outer electron shell is full.

3.

Why does a full outer electron shell make a noble gas chemically stable?

A full outer shell is stable, so atoms have little tendency to gain, lose or share electrons.

4.

Why do noble gases generally not gain electrons?

Their outer shell is already full.

5.

Why do noble gases generally not lose or share electrons?

Their outer shell is already stable and full.

6.

How does the electronic configuration of noble gases explain their lack of chemical reactivity?

Their full outer shells make them very stable and therefore 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 suitable for use in balloons?

Helium has a low density, so it provides lift.

2.

Why is helium used instead of hydrogen in many balloons?

Helium is non-flammable, making it safer than hydrogen.

3.

Why is argon used in some light bulbs?

Argon is inert, so it prevents the hot filament from reacting.

4.

Why is the inertness of argon useful when it is used in light bulbs?

It prevents reactions with the hot filament and extends the bulb's life.

5.

How does the non-flammability of noble gases make them useful in particular applications?

They can be used where a gas must not readily burn or support combustion.

6.

How do the inertness, low density and non-flammability of noble gases determine their uses?

Their properties determine suitable uses: low density for balloons, inertness for protective atmospheres and non-flammability where safety is important.

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 do the boiling points of noble gases change down Group 0?

They increase down Group 0.

2.

How does the density of noble gases change down Group 0?

They increase down Group 0.

3.

How does the physical state of noble gases change with increasing atomic number at room temperature?

They remain gases at room temperature.

4.

How can the physical property pattern of helium, neon and argon be used to predict the properties of krypton?

Krypton would be a gas with a higher boiling point and density than argon.

5.

How can the physical property pattern of noble gases be used to predict the density of a heavier noble gas?

Its density would be greater than that of lighter noble gases.

6.

What general trend occurs in the physical properties of noble gases as you move down Group 0?

Boiling point and density generally increase 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 could you investigate the effect of changing the concentration of hydrochloric acid on the rate of reaction with marble chips?

Use different concentrations of hydrochloric acid, keeping the mass and surface area of marble chips, volume, temperature and apparatus constant. Measure the volume of CO₂ produced over time.

2.

What measurement would you take when investigating the rate of reaction between hydrochloric acid and marble chips by measuring gas production?

Measure the volume of carbon dioxide gas produced at regular time intervals.

3.

How could the rate of the reaction between sodium thiosulfate and hydrochloric acid be investigated using a colour change?

Mix sodium thiosulfate with hydrochloric acid over a marked cross and measure the time taken for the cross to become obscured.

4.

What causes the visible colour change when sodium thiosulfate reacts with hydrochloric acid?

Sulfur precipitate forms, making the solution cloudy.

5.

How could you use experimental results to determine which of two conditions produces the faster reaction?

Compare the gradients of graphs or calculate the rate using change in measured quantity ÷ time. The greater rate is the faster reaction.

6.

What variables should be controlled when investigating the effect of one condition on the rate of a reaction?

Keep all variables constant except the condition being investigated, for example temperature, volumes, concentrations, mass/surface area and apparatus.

7.2 Suggest practical methods for determining the rate of a given reaction

1.

How could you measure the rate of a reaction that produces a gas?

Collect the gas in a gas syringe or measure its volume by displacement of water over time.

2.

How could you measure the rate of a reaction that produces a precipitate?

Measure the time taken for a precipitate to form to a specified opacity or to obscure a marked cross.

3.

How could you measure the rate of a reaction that causes a colour change?

Measure the time taken for a specified colour change to occur.

4.

How could you measure the rate of a reaction that causes a change in mass?

Measure the mass of the reaction mixture at regular time intervals using a balance.

5.

How could you determine the rate of a reaction from a graph of a measured quantity against time?

Calculate the gradient of the graph at the required point. A steeper gradient represents a faster rate.

6.

What factors should be considered when choosing a practical method for measuring the rate of a reaction?

Consider what measurable change occurs, the required accuracy, reaction speed, suitable apparatus and whether the method gives reliable results.

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 collide so that bonds can break and new bonds can form.

2.

What is meant by a successful collision?

A successful collision is one with sufficient energy, and the appropriate conditions/orientation, to cause a reaction.

3.

Why do some collisions between reacting particles not result in a reaction?

They may have insufficient energy to overcome the activation energy.

4.

How does increasing the frequency of collisions affect the rate of reaction?

More frequent collisions increase the number of opportunities for successful collisions, increasing the rate.

5.

How does increasing the energy of collisions affect the rate of reaction?

Higher-energy collisions are more likely to have enough energy to overcome the activation energy, increasing the rate.

6.

How does collision theory explain why a reaction becomes faster when there are more frequent successful collisions?

A faster reaction has more successful collisions per unit time.

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.

Explain why increasing the temperature increases the rate of a reaction in terms of particle collisions.

Particles have more kinetic energy, move faster and collide more frequently. A greater proportion also have enough energy to overcome the activation energy.

2.

Explain why increasing the concentration of a solution increases the rate of a reaction.

There are more reactant particles per unit volume, so collisions occur more frequently.

3.

Explain why increasing the surface area to volume ratio of a solid increases the rate of reaction.

More particles are exposed at the surface, so collisions between reactant particles occur more frequently.

4.

Explain why increasing the pressure of reacting gases increases the rate of reaction.

Gas particles are closer together, increasing the frequency of collisions.

5.

How does increasing temperature affect both the frequency and energy of collisions between particles?

Increasing temperature increases the frequency of collisions and the average energy of the particles, so more collisions have enough energy to be successful.

6.

Why does increasing concentration, surface area to volume ratio or gas pressure increase the frequency of collisions?

The particles are closer together or more exposed, increasing the frequency of collisions.

7.5 Interpret graphs of mass, volume or concentration of reactant or product against time

1.

What does the gradient of a graph of product volume against time represent?

It represents the rate of gas production.

2.

What does a steep gradient on a reaction graph indicate about the rate of reaction?

A steep gradient indicates a fast rate of reaction.

3.

What does a horizontal section of a graph of reactant concentration against time indicate?

The reaction has stopped because the reactant concentration is no longer changing.

4.

A reaction produces 60 cm³ of gas in 30 seconds. Calculate the mean rate of gas production.

Mean rate = 60 ÷ 30 = 2 cm³ s⁻¹.

5.

A reaction produces 90 cm³ of gas in 45 seconds. Calculate the mean rate of gas production and give the unit.

Mean rate = 90 ÷ 45 = 2 cm³ s⁻¹.

6.

Two reactions produce the same final volume of gas, but one graph has a steeper initial gradient. Which reaction has the greater initial rate, and why?

The reaction with the steeper initial gradient has the greater initial rate because the measured quantity changes more rapidly at the start.

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 catalyst is a substance that increases the rate of a reaction without being chemically changed or used up.

2.

How does a catalyst affect the rate of a chemical reaction?

It increases the rate of reaction.

3.

Does a catalyst alter the products formed in a reaction?

No. It does not alter the products formed.

4.

What happens to a catalyst chemically 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.

6.

Why can a catalyst be used repeatedly in principle without being used up?

Because it is not used up or chemically changed overall during the reaction.

7.7 Explain how the addition of a catalyst increases the rate of a reaction in terms of activation energy

1.

What is meant by the activation energy of a reaction?

Activation energy is the minimum energy particles must have for a reaction to occur.

2.

How does a catalyst affect the activation energy of a reaction?

A catalyst lowers the activation energy.

3.

Why does lowering the activation energy increase the rate of a reaction?

More particles have enough energy to overcome the lower activation energy, increasing the number of successful collisions.

4.

How does a catalyst provide an alternative reaction pathway?

It provides an alternative reaction pathway with a lower activation energy.

5.

Why does lowering activation energy result in more successful collisions?

More collisions have sufficient energy to overcome the lower activation energy, so more are successful.

6.

How would the activation energy shown on a reaction profile differ for a catalysed and uncatalysed reaction?

The catalysed reaction profile has a lower activation-energy peak than the uncatalysed profile.

7.8 Recall that enzymes are biological catalysts and that enzymes are used in the production of alcoholic drinks

1.

What are enzymes?

Enzymes are biological catalysts.

2.

Why are enzymes described as biological catalysts?

They speed up reactions in living organisms without being used up.

3.

How do enzymes affect the rate of reactions in living organisms?

They increase the rates of biochemical reactions.

4.

Why are enzymes involved in the production of alcoholic drinks?

Enzymes catalyse reactions involved in fermentation.

5.

What role do enzymes from yeast play during the production of alcoholic drinks?

Yeast enzymes catalyse the conversion of sugars into ethanol and carbon dioxide.

6.

Why are enzymes useful catalysts in biological and industrial processes?

They allow biological reactions to occur rapidly under relatively mild conditions.

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 some salts dissolve in water?

Heat energy may be absorbed or released.

2.

Why can a temperature change be measured during a neutralisation reaction?

Chemical reactions involve energy changes, which can cause the temperature of the solution to change.

3.

Why can a temperature change be measured during a displacement reaction?

Energy is transferred during the reaction, causing the temperature of the solution to change.

4.

Why can a temperature change be measured during a precipitation reaction?

Energy is transferred when the precipitate forms, causing a temperature change.

5.

How can a temperature change be used to determine whether a reaction in solution releases or absorbs heat energy?

A temperature increase indicates heat has been released; a temperature decrease indicates heat has been absorbed.

6.

What four types of chemical change in solution can be investigated by measuring temperature changes?

Dissolving salts, neutralisation, displacement and precipitation reactions.

7.10 Describe an exothermic change or reaction as one in which heat energy is given out

1.

What is an exothermic reaction?

An exothermic reaction is a reaction in which heat energy is given out.

2.

What happens to heat energy during an exothermic reaction?

Heat energy is transferred to the surroundings.

3.

What happens to the temperature of the surroundings during an exothermic reaction?

The temperature of the surroundings increases.

4.

Why does an exothermic reaction cause the surroundings to become warmer?

Heat energy is released from the reacting system into the surroundings.

5.

How can a temperature increase provide evidence that a reaction is exothermic?

A temperature increase shows that heat energy has been transferred to the surroundings.

6.

Give one example of a chemical reaction or change that could be exothermic.

Neutralisation of an acid with an alkali.

7.11 Describe an endothermic change or reaction as one in which heat energy is taken in

1.

What is an endothermic reaction?

An endothermic reaction is a reaction in which heat energy is taken in.

2.

What happens to heat energy during an endothermic reaction?

Heat energy is absorbed from the surroundings.

3.

What happens to the temperature of the surroundings during an endothermic reaction?

The temperature of the surroundings decreases.

4.

Why does an endothermic reaction cause the surroundings to become cooler?

Heat energy is transferred from the surroundings into the reacting system.

5.

How can a temperature decrease provide evidence that a reaction is endothermic?

A temperature decrease indicates that heat energy has been absorbed.

6.

Give one example of a chemical reaction or change that could be endothermic.

The thermal decomposition of calcium carbonate.

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 forces holding the atoms together.

2.

Why is breaking a chemical bond an endothermic process?

Energy is taken in to break the bond.

3.

What happens to energy when a new chemical bond is formed?

Energy is released.

4.

Why is making a chemical bond an exothermic process?

Forming a bond releases energy as the atoms become more stable.

5.

Which releases energy: breaking bonds or making bonds?

Making bonds.

6.

Which requires energy: breaking bonds or making bonds?

Breaking bonds.

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 a reaction exothermic in terms of the energy required to break bonds and the energy released when bonds form?

More energy is released forming product bonds than is required to break reactant bonds.

2.

When is a reaction endothermic in terms of the energy required to break bonds and the energy released when bonds form?

Less energy is released forming product bonds than is required to break reactant bonds.

3.

A reaction requires 800 kJ mol⁻¹ to break bonds and releases 1100 kJ mol⁻¹ when new bonds form. Is the reaction exothermic or endothermic?

Exothermic, because 1100 kJ mol⁻¹ is released and only 800 kJ mol⁻¹ is required.

4.

A reaction requires 1400 kJ mol⁻¹ to break bonds and releases 1000 kJ mol⁻¹ when new bonds form. Is the reaction exothermic or endothermic?

Endothermic, because 1400 kJ mol⁻¹ is required and only 1000 kJ mol⁻¹ is released.

5.

Why is a reaction exothermic when more energy is released forming bonds than is required to break bonds?

The excess energy is transferred to the surroundings as heat.

6.

Why is a reaction endothermic when less energy is released forming bonds than is required to break bonds?

More energy is required to break the reactant bonds than is released when the product bonds form, so energy is taken in overall.

7.14 Calculate the energy change in a reaction given the energies of bonds (in kJ mol–1)

1.

The bond energies are H–H = 436 kJ mol⁻¹, Cl–Cl = 242 kJ mol⁻¹ and H–Cl = 431 kJ mol⁻¹. Calculate the energy change for H₂ + Cl₂ → 2HCl.

Bonds broken = 436 + 242 = 678 kJ mol⁻¹. Bonds formed = 2 × 431 = 862 kJ mol⁻¹. Energy change = 678 − 862 = −184 kJ mol⁻¹.

2.

The bond energies are H–H = 436 kJ mol⁻¹, O=O = 498 kJ mol⁻¹ and O–H = 463 kJ mol⁻¹. Calculate the energy change for 2H₂ + O₂ → 2H₂O.

Bonds broken = (2 × 436) + 498 = 1370 kJ mol⁻¹. Bonds formed = 4 × 463 = 1852 kJ mol⁻¹. Energy change = 1370 − 1852 = −482 kJ mol⁻¹.

3.

A reaction requires 1250 kJ mol⁻¹ to break the bonds in the reactants and releases 1500 kJ mol⁻¹ when bonds form in the products. Calculate the overall energy change and state whether the reaction is exothermic or endothermic.

Energy change = 1250 − 1500 = −250 kJ mol⁻¹. The negative value means the reaction is exothermic.

4.

A reaction requires 980 kJ mol⁻¹ to break bonds and releases 760 kJ mol⁻¹ when new bonds form. Calculate the overall energy change and state whether the reaction is exothermic or endothermic.

Energy change = 980 − 760 = +220 kJ mol⁻¹. The positive value means the reaction is endothermic.

5.

The energy required to break bonds is 1650 kJ mol⁻¹ and the energy released when new bonds form is 1920 kJ mol⁻¹. Calculate the overall energy change and explain what the sign of your answer means.

Energy change = 1650 − 1920 = −270 kJ mol⁻¹. The negative sign means 270 kJ mol⁻¹ of heat energy is released overall, so the reaction is exothermic.

6.

A reaction has an overall energy change of +240 kJ mol⁻¹. The energy required to break the reactant bonds is 1350 kJ mol⁻¹. Calculate the energy released when the product bonds form.

Energy released = 1350 + 240 = 1590 kJ mol⁻¹.

7.15 Explain the term activation energy

1.

What is meant by the activation energy of a chemical reaction?

The minimum energy required for a reaction to occur.

2.

Why must reacting particles have at least the activation energy for a reaction to occur?

Particles need sufficient energy to overcome the activation energy and allow bonds to break and new bonds to form.

3.

How is activation energy related to successful collisions?

A successful collision must have at least the activation energy.

4.

What happens to the rate of a reaction if fewer particles have enough energy to overcome the activation energy?

The rate decreases because fewer collisions have sufficient energy to be successful.

5.

Why does increasing temperature usually increase the number of particles able to overcome the activation energy?

Particles have greater kinetic energy, so a greater proportion have energy equal to or above the activation energy.

6.

How does a catalyst affect the activation energy of a reaction?

A catalyst lowers the activation energy.

7.16 Draw and label reaction profiles for endothermic and exothermic reactions, identifying activation energy

1.

What does a reaction profile show?

A reaction profile shows the energy level of reactants and products during a reaction.

2.

How is the activation energy shown on a reaction profile?

It is shown as the energy difference between the reactants and the peak of the profile.

3.

How can you identify an exothermic reaction from its reaction profile?

The products are at a lower energy level than the reactants.

4.

How can you identify an endothermic reaction from its reaction profile?

The products are at a higher energy level than the reactants.

5.

Where are the reactants, products and activation energy labelled on a reaction profile?

Reactants are labelled at the initial energy level, products at the final energy level and activation energy from the reactant level to the peak.

6.

How does the energy level of the products compare with the energy level of the reactants in an exothermic reaction and in an endothermic reaction?

Exothermic: products have lower energy than reactants. Endothermic: products have higher energy than reactants.

8.1 Recall that hydrocarbons are compounds that contain carbon and hydrogen only

1.

What elements are present in a hydrocarbon?

Carbon and hydrogen only.

2.

What is the defining feature of a hydrocarbon?

It contains only carbon and hydrogen atoms.

3.

Is methane, CH₄, a hydrocarbon? Explain why.

Yes. CH₄ contains only carbon and hydrogen.

4.

Is ethanol, C₂H₅OH, a hydrocarbon? Explain why.

No. Ethanol also contains oxygen.

5.

Is carbon dioxide, CO₂, a hydrocarbon? Explain why.

No. Carbon dioxide contains carbon and oxygen, not hydrogen.

6.

Classify C₃H₈, C₂H₄ and C₆H₆ as hydrocarbons or not hydrocarbons, giving a reason for each.

C₃H₈ — hydrocarbon; C₂H₄ — hydrocarbon; C₆H₆ — hydrocarbon. Each 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?

Crude oil is a complex mixture of hydrocarbons.

2.

Why is crude oil described as a complex mixture of hydrocarbons?

It contains many different hydrocarbon molecules with different sizes and structures.

3.

How can the carbon atoms in crude oil molecules be arranged?

Carbon atoms can be arranged in chains or rings.

4.

Why is crude oil an important source of useful substances?

It provides fuels and feedstock for the petrochemical industry.

5.

What is meant by saying that crude oil is a finite resource?

It is a limited resource that will eventually run out.

6.

Explain why crude oil is both an important resource and a finite resource.

It is important because it provides fuels and chemical feedstocks, but it is finite because its supply is limited and cannot be replaced on a human timescale.

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 fractions?

Fractional distillation.

2.

Why can the hydrocarbons in crude oil be separated by fractional distillation?

They have different boiling points.

3.

What happens to crude oil when it is heated during fractional distillation?

It is heated so that most of the hydrocarbons vaporise.

4.

What happens to hydrocarbons with different boiling points as they pass through the fractionating column?

They condense at different temperatures as the vapours rise through the column.

5.

Why do hydrocarbons condense at different levels in the fractionating column?

The column is hotter at the bottom and cooler at the top, so hydrocarbons condense at different levels according to their boiling points.

6.

Explain how fractional distillation produces simpler, more useful mixtures from crude oil.

Crude oil is vaporised and passed into a fractionating column, where hydrocarbons condense at different temperatures, separating them into fractions with similar 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?

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 are the uses of fuel oil?

Fuel for large ships and some power stations.

6.

What is bitumen used for?

To surface roads and roofs.

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 and hydrogen atoms generally change as the hydrocarbons become larger?

The molecules generally contain more carbon and hydrogen atoms.

2.

How does boiling point change as the hydrocarbon molecules become larger?

Boiling points increase.

3.

How does ease of ignition change as the hydrocarbon molecules become larger?

Larger hydrocarbons are harder to ignite.

4.

How does viscosity change as the hydrocarbon molecules become larger?

Viscosity increases.

5.

Which homologous series do most hydrocarbons in crude oil fractions belong to?

The alkane homologous series.

6.

A fraction contains hydrocarbons with higher boiling points and greater viscosity than another fraction. What can you predict about the relative sizes of their molecules?

The fraction with higher boiling points and greater viscosity contains larger molecules.

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 meant by the term homologous series?

A homologous series is a series of compounds with the same general formula and similar chemical properties, with neighbouring members differing by CH₂.

2.

How do the molecular formulae of neighbouring members of a homologous series differ?

They differ by CH₂.

3.

What formula is shared by members of an homologous series?

They share the same general formula.

4.

How do physical properties change across a homologous series?

Physical properties change gradually, for example boiling point increases with molecular size.

5.

Why do members of the same homologous series have similar chemical properties?

They have the same functional group and similar bonding, giving them similar chemical behaviour.

6.

The boiling points of four members of a homologous series are 36°C, 69°C, 98°C and 126°C. What does this trend suggest about their physical properties as molecular size increases?

The gradual increase in boiling point shows a gradual change in physical properties as molecular size increases.

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 during the complete combustion of a hydrocarbon?

Carbon dioxide and water.

2.

What gas must be present for complete combustion to occur?

Oxygen.

3.

What happens to the carbon atoms during complete combustion?

They are oxidised to form carbon dioxide.

4.

What happens to the hydrogen atoms during complete combustion?

They are oxidised to form water.

5.

Is complete combustion exothermic or endothermic?

Exothermic.

6.

Write a word equation for the complete combustion of a hydrocarbon.

Hydrocarbon + oxygen → carbon dioxide + water.

8.8 Explain why the incomplete combustion of hydrocarbons can produce carbon and carbon monoxide

1.

What causes incomplete combustion of a hydrocarbon?

There is insufficient oxygen for complete combustion.

2.

Why can incomplete combustion produce carbon monoxide?

There is insufficient oxygen to fully oxidise the carbon to carbon dioxide.

3.

Why can incomplete combustion produce carbon?

Some carbon is left unoxidised when there is insufficient oxygen.

4.

How does the oxygen supply affect whether combustion is complete or incomplete?

A plentiful oxygen supply favours complete combustion; a restricted supply causes incomplete combustion.

5.

What are the two carbon-containing products that can form during incomplete combustion?

Carbon and carbon monoxide.

6.

Explain why a restricted oxygen supply can cause a hydrocarbon fuel to undergo incomplete combustion.

There is not enough oxygen to completely oxidise all the carbon to carbon dioxide, so carbon monoxide and/or carbon can form.

8.9 Explain how carbon monoxide behaves as a toxic gas

1.

Why is carbon monoxide toxic?

It binds strongly to haemoglobin, reducing its ability to carry oxygen.

2.

What happens to haemoglobin when carbon monoxide enters the blood?

It binds to haemoglobin and forms carboxyhaemoglobin.

3.

Why does carbon monoxide reduce the amount of oxygen transported around the body?

Carbon monoxide occupies haemoglobin binding sites that would normally carry oxygen.

4.

Why can exposure to carbon monoxide be dangerous even when the gas cannot easily be detected?

It is colourless and odourless, so dangerous concentrations may go unnoticed.

5.

What effect can carbon monoxide have on cells that require oxygen for respiration?

Cells receive less oxygen for aerobic respiration.

6.

Explain why carbon monoxide poisoning can be fatal.

Severe oxygen deprivation can prevent vital organs and cells from respiring properly, potentially causing death.

8.10 Describe the problems caused by incomplete combustion producing carbon monoxide and soot in appliances that use carbon compounds as fuels

1.

What toxic gas can be produced by incomplete combustion?

Carbon monoxide.

2.

What solid pollutant can be produced by incomplete combustion?

Soot/carbon particles.

3.

Why is carbon monoxide particularly dangerous in poorly ventilated rooms?

Carbon monoxide can build up to dangerous concentrations and reduce oxygen transport in the blood.

4.

What health problems can be caused by breathing in soot particles?

Soot particles can damage the respiratory system and contribute to breathing problems.

5.

Why can incomplete combustion make fuel-burning appliances dangerous?

They can release toxic carbon monoxide and harmful soot.

6.

Explain why adequate ventilation is important when using appliances that burn carbon compounds.

Ventilation provides enough oxygen for combustion and prevents dangerous carbon monoxide from accumulating.

8.11 Explain how impurities in some hydrocarbon fuels result in the production of sulfur dioxide

1.

What impurity in some hydrocarbon fuels can lead to sulfur dioxide formation?

Sulfur-containing impurities.

2.

What happens to sulfur-containing impurities when the fuel is burned?

They react with oxygen during combustion.

3.

Which gas is produced when sulfur impurities are oxidised during combustion?

Sulfur dioxide, SO₂.

4.

Write a word equation for the formation of sulfur dioxide when sulfur burns in oxygen.

Sulfur + oxygen → sulfur dioxide.

5.

Why can burning some fossil fuels release sulfur dioxide into the atmosphere?

Sulfur-containing impurities are oxidised during combustion.

6.

Explain the link between sulfur impurities in fuels and sulfur dioxide pollution.

Sulfur impurities in fuels burn and react with oxygen to form sulfur dioxide, which is released into the atmosphere.

8.12 Explain some problems associated with acid rain caused when sulfur dioxide dissolves in rain water

1.

What happens when sulfur dioxide dissolves in rainwater?

It dissolves and forms an acidic solution.

2.

Why can sulfur dioxide contribute to acid rain?

The resulting acidic water contributes to acid rain.

3.

What effect can acid rain have on aquatic ecosystems?

It can lower the pH of water, harming aquatic organisms.

4.

What effect can acid rain have on plants and trees?

It can damage plants and trees.

5.

How can acid rain damage buildings and statues made from carbonate-containing materials?

Acids react with and dissolve carbonate-containing materials.

6.

Explain how burning sulfur-containing fuels can eventually cause environmental damage through acid rain.

Sulfur-containing fuels produce SO₂; SO₂ dissolves in rainwater, producing acidic rain that can damage ecosystems, vegetation and buildings.

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.

Which two gases from the air can react in a hot engine?

Nitrogen and oxygen.

2.

Why can nitrogen and oxygen react together inside an engine?

The high temperature provides enough energy for them to react.

3.

What are the products called when nitrogen and oxygen react at high temperatures?

Oxides of nitrogen, or nitrogen oxides.

4.

Why are oxides of nitrogen considered pollutants?

They are harmful air pollutants and can contribute to problems such as acid rain and respiratory irritation.

5.

Why does the high temperature inside an engine promote the formation of nitrogen oxides?

More particles have sufficient energy to overcome the activation energy.

6.

Explain how burning fuels in engines can lead to the formation of oxides of nitrogen.

Burning fuel produces very high temperatures inside the engine, allowing nitrogen and oxygen from the air to react and form nitrogen oxides.

8.14 Evaluate the advantages and disadvantages of using hydrogen, rather than petrol, as a fuel in cars

1.

What is the main environmental advantage of using hydrogen instead of petrol in a car?

Hydrogen produces no carbon dioxide at the point of use.

2.

What product is formed when hydrogen reacts with oxygen in a fuel cell?

Water.

3.

Why does hydrogen not produce carbon dioxide when it is used as a fuel?

Hydrogen contains no carbon, so its reaction with oxygen produces water rather than carbon dioxide.

4.

What are some disadvantages associated with producing and storing hydrogen?

Hydrogen can be difficult and energy-intensive to produce, and it is difficult to store because it has low density.

5.

Why does the environmental impact of hydrogen depend on how the hydrogen is produced?

If hydrogen is produced using fossil fuels, carbon dioxide can be produced during its manufacture; renewable production has lower associated emissions.

6.

Evaluate whether hydrogen is a better car fuel than petrol, considering emissions, production, storage and availability.

Hydrogen can reduce direct carbon emissions and produces water in a fuel cell, but production, storage, transport, infrastructure and availability present disadvantages. Its overall environmental benefit depends on how it is produced.

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.

Which fossil resource is petrol obtained from?

Crude oil.

2.

Which fossil resource is kerosene obtained from?

Crude oil.

3.

Which fossil resource is diesel oil obtained from?

Crude oil.

4.

Where is methane found naturally as a fossil fuel?

Natural gas.

5.

Why are petrol, kerosene, diesel oil and methane described as non-renewable?

They are finite resources formed over very long periods and cannot be replaced quickly.

6.

Identify whether petrol, kerosene, diesel oil and methane are obtained from crude oil or natural gas.

Petrol — crude oil; kerosene — crude oil; diesel oil — crude oil; methane — natural gas.

8.16 Explain why cracking involves the breaking down of larger, saturated hydrocarbon molecules (alkanes) into smaller, more useful ones, some of which are unsaturated (alkenes)

1.

What type of hydrocarbon molecules are broken down during cracking?

Large, saturated hydrocarbons called alkanes.

2.

What happens to large alkane molecules during cracking?

They are broken down into smaller hydrocarbon molecules.

3.

Why are the products of cracking generally smaller molecules?

Cracking breaks large molecules into smaller molecules.

4.

Why can cracking produce alkenes?

Hydrogen atoms can be removed during cracking, leaving carbon–carbon double bonds.

5.

What type of hydrocarbon is an alkene?

An unsaturated hydrocarbon containing a carbon–carbon double bond.

6.

Explain why cracking can produce both smaller alkanes and alkenes from larger alkanes.

Breaking large alkanes can produce smaller alkanes and alkenes, with some products having fewer hydrogen atoms and therefore containing C=C bonds.

8.17 Explain why cracking is necessary

1.

Why is cracking necessary when processing crude oil?

To convert less useful large hydrocarbons into more useful smaller hydrocarbons.

2.

Why is there greater demand for some fractions than others?

There is greater demand for smaller, more easily ignited fractions than for some larger fractions.

3.

How does cracking help meet the demand for smaller hydrocarbons?

It produces additional smaller hydrocarbons.

4.

Why are alkenes produced by cracking useful to the chemical industry?

They are used to make polymers and other useful chemicals.

5.

How does cracking increase the usefulness of crude oil?

It converts surplus larger hydrocarbons into products with greater demand and usefulness.

6.

Explain why the supply of fractions obtained directly from crude oil does not always match their demand.

Fractional distillation produces fractions in proportions that do not necessarily match demand, so cracking changes larger molecules into more useful products.

8.18 Recall that the gases produced by volcanic activity formed the Earth’s early atmosphere

1.

What was the main source of gases in the Earth's early atmosphere?

Volcanic activity.

2.

How did volcanic activity contribute to the formation of the early atmosphere?

Volcanoes released gases into the atmosphere.

3.

What happened to gases released by volcanoes over time?

They accumulated around the Earth to form the early atmosphere.

4.

Why was the early atmosphere very different from today's atmosphere?

Its composition reflected the gases released by volcanic activity and differed greatly from the modern atmosphere.

5.

Which geological process released large amounts of gases to form the early atmosphere?

Volcanic activity.

6.

Explain how volcanic activity led to the formation of Earth's early atmosphere.

Volcanic eruptions released gases such as water vapour and carbon dioxide, which accumulated to form Earth's early atmosphere.

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 approximate amount of oxygen thought to be present in Earth's early atmosphere?

Little or no oxygen.

2.

Which gas was present in large amounts in the early atmosphere?

Carbon dioxide.

3.

Why was water vapour present in the early atmosphere?

Volcanic activity released large amounts of water vapour.

4.

What other gases were thought to be present in small amounts?

Small amounts of other gases, such as nitrogen.

5.

What evidence can scientists use to infer the composition of Earth's early atmosphere?

Evidence from rocks, fossils and geological records can be used to infer past atmospheric conditions.

6.

Explain how evidence can be used to support the model of an early atmosphere containing little or no oxygen and large amounts of carbon dioxide.

Geological evidence can indicate that early Earth had little or no oxygen and substantial carbon dioxide, supporting the model of the early atmosphere.

8.20 Explain how condensation of water vapour formed oceans

1.

What caused water vapour in the early atmosphere to condense?

The Earth cooled.

2.

What happens to water vapour when it condenses?

Water vapour changed into liquid water.

3.

Why did cooling of Earth allow liquid water to form?

Cooling reduced the temperature enough for water vapour to condense.

4.

How did condensation contribute to the formation of the oceans?

Liquid water accumulated on Earth's surface, forming oceans.

5.

What change of state is involved when water vapour becomes liquid water?

Gas → liquid.

6.

Explain how cooling of the early Earth led to the formation of oceans.

As the early Earth cooled, water vapour condensed into liquid water, which accumulated to form the oceans.

8.21 Explain how the amount of carbon dioxide in the atmosphere was decreased when carbon dioxide dissolved as the oceans formed

1.

What happened to some atmospheric carbon dioxide when the oceans formed?

Some atmospheric carbon dioxide dissolved in the newly formed oceans.

2.

Why did the formation of oceans decrease atmospheric carbon dioxide?

Carbon dioxide was transferred from the atmosphere into the oceans.

3.

What happens when carbon dioxide dissolves in ocean water?

It dissolves in water and can form dissolved carbon dioxide and carbonic acid.

4.

How did the formation of liquid oceans affect the composition of the atmosphere?

The atmosphere became less rich in carbon dioxide.

5.

Why did the amount of atmospheric carbon dioxide decrease as the oceans formed?

As oceans formed, increasing amounts of CO₂ dissolved in them, reducing its atmospheric concentration.

6.

Explain the link between ocean formation and the reduction in atmospheric carbon dioxide.

Cooling formed liquid oceans; atmospheric CO₂ dissolved into the water, reducing the amount of CO₂ remaining in the atmosphere.

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.

Which gas did primitive plants remove from the atmosphere during photosynthesis?

Carbon dioxide.

2.

Which gas did primitive plants release during photosynthesis?

Oxygen.

3.

How did the growth of primitive plants affect atmospheric oxygen levels?

Atmospheric oxygen gradually increased.

4.

Why did atmospheric oxygen increase gradually rather than immediately?

Photosynthesis occurred over a long period as primitive plants grew and accumulated.

5.

What process carried out by primitive plants caused oxygen to be released?

Photosynthesis.

6.

Explain how photosynthesis by primitive plants changed the composition of Earth's atmosphere.

Primitive plants removed CO₂ by photosynthesis and released O₂, causing atmospheric oxygen levels to gradually increase.

8.23 Describe the chemical test for oxygen

1.

What is the chemical test for oxygen?

Insert a glowing splint into the gas.

2.

What happens to a glowing splint when it is placed in oxygen?

The glowing splint relights.

3.

What observation indicates a positive test for oxygen?

Relighting of the glowing splint.

4.

Why is a glowing splint used rather than a lit splint?

Oxygen supports combustion, so it can relight a glowing splint without the splint already being aflame.

5.

What gas is identified if a glowing splint relights?

Oxygen.

6.

Describe how you would test an unknown gas to determine whether it contains oxygen.

Place a glowing splint into the unknown gas. If it relights, oxygen is present.

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 meant by the greenhouse effect?

The greenhouse effect is the warming of Earth caused by atmospheric gases absorbing and re-emitting radiation from Earth.

2.

Which type of radiation is absorbed by greenhouse gases after being emitted by Earth?

Infrared radiation.

3.

Name three greenhouse gases.

Carbon dioxide, methane and water vapour.

4.

What happens to the energy absorbed by greenhouse gases?

They subsequently release/re-emit the absorbed energy.

5.

How does the greenhouse effect help keep Earth's surface warm?

Energy is retained in the Earth-atmosphere system, keeping the Earth warmer.

6.

Explain how carbon dioxide, methane and water vapour contribute to the greenhouse effect.

Earth emits infrared radiation; greenhouse gases absorb some of this radiation and re-emit energy, helping keep the Earth warm.

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 exists between atmospheric carbon dioxide concentration and global temperature?

Atmospheric CO₂ concentration and global temperature have increased over the relevant period, showing a correlation.

2.

How is fossil fuel consumption linked to changes in atmospheric carbon dioxide concentration?

Burning fossil fuels releases CO₂, increasing atmospheric CO₂ concentration.

3.

Why can a correlation between carbon dioxide concentration and temperature be used as evidence for human-caused climate change?

The linked increases provide evidence that human activity may be contributing to climate change.

4.

Why does correlation alone not prove causation?

Correlation shows an association but does not by itself establish that one factor causes the other.

5.

How can the location of measurements introduce uncertainty into climate data?

Different locations can experience different temperatures and CO₂ concentrations, so measurements may not represent global conditions.

6.

Evaluate evidence for human activity causing climate change, considering carbon dioxide concentration, fossil fuel consumption, temperature records and uncertainties in historical measurements.

Increasing atmospheric CO₂, fossil fuel consumption and temperature provide evidence supporting a human contribution to climate change. However, uncertainty arises from measurement locations, limited historical data and the accuracy of older records.

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?

Approximately 78% nitrogen, 21% oxygen and about 1% other gases, including argon and carbon dioxide.

2.

Which human activities can increase atmospheric carbon dioxide levels?

Burning fossil fuels and deforestation can increase atmospheric CO₂.

3.

How can livestock farming increase atmospheric methane levels?

Livestock farming produces methane, particularly from ruminant digestion and waste.

4.

What potential effects could increased carbon dioxide and methane have on climate?

Increased greenhouse gas concentrations can enhance the greenhouse effect and contribute to global warming and climate change.

5.

What is meant by mitigating climate change?

Mitigating climate change means reducing or limiting its causes and/or effects.

6.

Evaluate possible methods of mitigating increased greenhouse gas levels, considering their scale, risks and environmental implications.

Mitigation can include reducing fossil fuel use, increasing renewable energy, improving energy efficiency and reducing methane emissions. Methods must be considered in terms of their scale, risks, costs and environmental effects.

Chemistry

Edexcel GCSE Combined 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 change

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