AQA Combined Science: Chemistry — Recall & Retrieval

5.1.1.1 Atoms, Elements and Compounds

1.

What is an atom?

The smallest particle of an element that can exist and retain the properties of that element.

2.

What is an element?

A substance made of only one type of atom.

3.

How are atoms of different elements represented?

By their chemical symbols.

4.

How many different elements are there approximately?

Approximately 100 different elements.

5.

What is a compound?

A substance made from two or more different elements chemically bonded together.

6.

Explain why compounds can only be separated into elements by chemical reactions.

Compounds contain elements that are chemically bonded, so chemical reactions are needed to break the bonds and separate them.

5.1.1.2 Mixtures

1.

What is a mixture?

Two or more substances mixed together that are not chemically bonded.

2.

How are the properties of substances in a mixture affected when they are mixed?

The properties of substances in a mixture are not changed because no chemical bonds are formed.

3.

What is the difference between a mixture and a compound?

A mixture contains substances that are not chemically bonded and can be separated by physical methods; a compound contains elements chemically bonded and requires a chemical reaction to separate.

4.

Which separation technique is used to separate an insoluble solid from a liquid?

Filtration.

5.

Which separation technique is used to separate substances based on differences in boiling points?

Fractional distillation.

6.

Explain why physical separation methods do not produce new substances.

Physical separation methods do not produce new substances because no chemical bonds are broken or formed.

5.1.1.3 The Development of the Model of the Atom

1.

What did scientists originally believe atoms were like before the discovery of electrons?

Atoms were originally thought to be tiny solid spheres that could not be divided.

2.

What discovery led to the development of the plum pudding model?

The discovery of electrons.

3.

Describe the plum pudding model of the atom.

A positive sphere containing negatively charged electrons.

4.

What did the alpha particle scattering experiment show about the atom?

The atom was mostly empty space with a small, dense, positively charged nucleus.

5.

Why was the nuclear model developed to replace the plum pudding model?

The nuclear model was developed because the plum pudding model could not explain the results of the alpha particle scattering experiment.

6.

Explain how scientific evidence can lead to changes in scientific models.

Scientific evidence can show that existing models are incorrect, leading to new models being developed.

5.1.1.4 Relative Electrical Charges of Subatomic Particles

1.

What is the relative charge of a proton?

+1

2.

What is the relative charge of a neutron?

0

3.

What is the relative charge of an electron?

-1

4.

Why do atoms have no overall electrical charge?

Atoms have equal numbers of protons and electrons, so the positive and negative charges cancel out.

5.

What is the atomic number of an element?

The number of protons in an atom.

6.

Explain why all atoms of the same element have the same number of protons.

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

5.1.1.5 Size and Mass of Atoms

1.

What is the approximate radius of an atom?

Approximately 0.1 nm.

2.

Where is almost all of the mass of an atom found?

The nucleus.

3.

What particles are found in the nucleus of an atom?

Protons and neutrons.

4.

What is the mass number of an atom?

The total number of protons and neutrons in an atom.

5.

How are isotopes of an element different?

Isotopes have the same number of protons but different numbers of neutrons.

6.

An atom has an atomic number of 12 and a mass number of 24. Calculate the number of protons, neutrons and electrons in the atom.

Protons = 12, neutrons = 12, electrons = 12.

5.1.1.6 Relative Atomic Mass

1.

What is meant by relative atomic mass?

The weighted mean mass of atoms of an element compared with one twelfth of the mass of carbon-12.

2.

Why is relative atomic mass usually not a whole number?

Because it is an average value that includes isotopes with different masses.

3.

What information is needed to calculate relative atomic mass?

The mass and abundance of each isotope.

4.

Why does relative atomic mass take account of isotopes?

Because elements contain different isotopes.

5.

Calculate the relative atomic mass of an element with 75% isotope A of mass 20 and 25% isotope B of mass 22.

(20 × 75/100) + (22 × 25/100) = 20.5.

6.

Explain why the relative atomic mass of chlorine is approximately 35.5.

Chlorine has different isotopes, mainly chlorine-35 and chlorine-37, giving an average relative atomic mass of approximately 35.5.

5.1.1.7 Electronic Structure

1.

What are electrons arranged into around the nucleus?

Energy levels or electron shells.

2.

Which energy level do electrons fill first?

The lowest energy level first.

3.

How many electrons can the first shell hold?

Two electrons.

4.

Write the electronic structure of sodium.

Sodium: 2,8,1.

5.

Write the electronic structure of oxygen.

Oxygen: 2,6.

6.

Explain why the outer shell of an atom affects its chemical properties.

The outer shell determines the chemical properties and how an atom reacts.

5.1.2.1 The Periodic Table

1.

How are elements arranged in the periodic table?

Elements are arranged in order of increasing atomic number.

2.

What is the atomic number of an element?

The number of protons in an atom.

3.

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

Elements in the same group have the same number of electrons in their outer shell.

4.

What does the group number tell you about the electrons in an atom?

The group number tells you the number of electrons in the outer shell.

5.

How is the position of an element in the periodic table related to its electronic structure?

The period shows the number of occupied shells and the group shows the number of outer electrons.

6.

Predict the reactivity of an element from its position in the periodic table.

Reactivity can be predicted from the element's position and trends within its group.

5.1.2.2 Development of the Periodic Table

1.

How did early scientists arrange elements in the periodic table?

Early scientists arranged elements by atomic weight and properties.

2.

What problem occurred when elements were arranged only by atomic weight?

Some elements were placed incorrectly because atomic weight did not always match their properties.

3.

Who developed a periodic table that left gaps for undiscovered elements?

Dmitri Mendeleev.

4.

Why did Mendeleev leave gaps in his periodic table?

He left gaps for undiscovered elements and predicted their properties.

5.

How did the discovery of new elements support Mendeleev's predictions?

New elements were discovered that matched Mendeleev's predictions.

6.

Explain why knowledge of isotopes showed that atomic weight was not always the best way to arrange elements.

Isotopes showed that elements with different masses can have the same chemical properties, so atomic weight was not always suitable.

5.1.2.3 Metals and Non-metals

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 side of the periodic table.

3.

What type of ions do metals form?

Positive ions.

4.

What type of ions do non-metals form?

Negative ions.

5.

State one physical property of metals.

Metals are good conductors of electricity.

6.

Explain how the electronic structure of metals and non-metals affects their reactions.

Metals lose electrons to form positive ions, while non-metals gain electrons to form negative ions.

5.1.2.4 Group 0

1.

What are the elements in Group 0 called?

Noble gases.

2.

Why are noble gases unreactive?

They have a full outer electron shell.

3.

How many electrons do noble gases usually have in their outer shell?

Eight electrons, except helium which has two.

4.

Why is helium different from other noble gases?

Helium has one shell that is full with two electrons.

5.

How does boiling point change down Group 0?

Boiling points increase down Group 0.

6.

Explain why the properties of noble gases are related to their outer electrons.

Their full outer shells make them stable and unreactive.

5.1.2.5 Group 1

1.

What are Group 1 elements called?

Alkali metals.

2.

How many electrons do Group 1 elements have in their outer shell?

One electron.

3.

How does the reactivity of Group 1 metals change down the group?

Reactivity increases down the group.

4.

What gas is produced when Group 1 metals react with water?

Hydrogen.

5.

What type of ions do Group 1 metals form?

Positive ions.

6.

Explain why Group 1 metals become more reactive down the group.

The outer electron is further from the nucleus and easier to lose.

5.1.2.6 Group 7

1.

What are Group 7 elements called?

Halogens.

2.

How many electrons do halogens have in their outer shell?

Seven electrons.

3.

Are halogens metals or non-metals?

Non-metals.

4.

How does reactivity change down Group 7?

Reactivity decreases down the group.

5.

What happens when a more reactive halogen reacts with a salt of a less reactive halogen?

A more reactive halogen displaces a less reactive halogen from its salt.

6.

Explain why iodine has a higher boiling point than chlorine.

Iodine has stronger intermolecular forces because it has larger molecules and more electrons.

Topic 1 Review

1.

Explain how the atomic model changed as new evidence was discovered.

The atomic model changed as new evidence was discovered, including the discovery of electrons and the nucleus.

2.

Describe the structure of an atom using protons, neutrons and electrons.

An atom contains protons and neutrons in the nucleus with electrons arranged in shells around it.

3.

Explain how electronic structure determines the position of an element in the periodic table.

Electronic structure determines position because the number of shells gives the period and the number of outer electrons gives the group.

4.

Compare the properties and reactions of Group 1 metals and Group 7 halogens.

Group 1 metals become more reactive down the group, while Group 7 halogens become less reactive down the group.

5.

Explain how the periodic table allows scientists to predict the properties of elements.

The periodic table allows scientists to predict properties because elements in the same group have similar electron arrangements.

6.

Explain the difference between atoms, elements, compounds and mixtures

Atoms are single particles, elements contain one type of atom, compounds contain chemically bonded elements, and mixtures contain substances that are not chemically bonded.

5.2.1.1 Chemical Bonds

1.

What are the three types of strong chemical bonds?

Ionic, covalent and metallic bonds.

2.

In ionic bonding, what particles are held together by electrostatic forces?

Positive and negative ions.

3.

In covalent bonding, what do atoms do with their electrons?

Share electrons.

4.

In metallic bonding, what type of electrons are shared between metal atoms?

Delocalised electrons.

5.

Which type of bonding occurs between metals and non-metals?

Ionic bonding.

6.

Explain how the transfer or sharing of electrons creates chemical bonds.

Chemical bonds form when atoms transfer or share electrons to achieve stable electron arrangements.

5.2.1.2 Ionic Bonding

1.

What happens to electrons when a metal atom reacts with a non-metal atom?

Electrons are transferred from the metal atom to the non-metal atom.

2.

Why does a metal atom become a positive ion when it forms an ionic bond?

It loses electrons and forms a positive ion.

3.

Why does a non-metal atom become a negative ion when it forms an ionic bond?

It gains electrons and forms a negative ion.

4.

Why do ions formed from Groups 1, 2, 6 and 7 elements have stable electron arrangements?

They gain or lose electrons to achieve a full outer shell.

5.

Draw a dot and cross diagram showing bonding in magnesium chloride

Magnesium loses two electrons and each chlorine atom gains one electron, forming Mg²⁺ and Cl⁻ ions with a giant ionic lattice structure.

6.

How can the charge of an ion be predicted from its group number in the periodic table?

The group number shows the number of electrons lost or gained to form a stable ion.

5.2.1.3 Ionic Compounds

1.

What type of structure do ionic compounds have?

A giant ionic lattice structure.

2.

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

Strong electrostatic forces of attraction between oppositely charged ions.

3.

Why do ionic compounds have high melting points?

A lot of energy is needed to overcome the strong forces between ions.

4.

Why can molten ionic compounds conduct electricity?

The ions are free to move and carry charge.

5.

Why can dissolved ionic compounds conduct electricity?

The ions are free to move through the solution and carry charge.

6.

How can the empirical formula of an ionic compound be found from a diagram showing its ions?

Use the charges of the ions to find the simplest ratio of ions.

5.2.1.4 Covalent Bonding

1.

What happens when atoms form covalent bonds?

They share electrons.

2.

Why are covalent bonds described as strong bonds?

Strong electrostatic forces of attraction exist between the nuclei and shared electrons.

3.

How can a substance be identified as a small molecule from its chemical formula?

A small molecule usually contains only a few atoms.

4.

What is the difference between a polymer structure and a small molecule structure?

Polymers are very large molecules made from repeating units; small molecules contain only a few atoms.

5.

What are examples of giant covalent structures?

Diamond, graphite and graphene.

6.

Draw a dot and cross diagram showing the bonding in carbon dioxide.

Carbon dioxide has a central carbon atom double-bonded to two oxygen atoms. Carbon shares two pairs of electrons with each oxygen, creating two double bonds to achieve a full outer shell

5.2.1.5 Metallic Bonding

1.

What type of structure do metals have?

A giant metallic structure.

2.

What happens to the outer electrons of metal atoms in metallic bonding?

Outer electrons become delocalised.

3.

What are delocalised electrons?

Electrons that are free to move through the structure.

4.

What causes the strong attraction in metallic bonding?

Strong attraction between positive metal ions and delocalised electrons.

5.

Show a labelled diagram of metallic bonding in calcium metal.

Positive calcium ions arranged in a lattice surrounded by delocalised electrons.

6.

Why can delocalised electrons move through a metal structure?

They are not attached to one atom and can move through the metal.

5.2.2.1 The Three States of Matter

1.

What are the three states of matter?

Solid, liquid and gas.

2.

What happens to particles when a solid melts into a liquid?

Particles gain energy, move faster and become less ordered.

3.

What happens to particles when a liquid boils into a gas?

Particles gain enough energy to overcome forces and spread apart.

4.

Why do substances with stronger forces between particles have higher melting points?

More energy is needed to overcome stronger forces between particles.

5.

How does particle theory explain changes of state?

Changes of state occur when particles gain or lose energy, changing their movement and arrangement.

6.

What are the limitations of the simple particle model of matter?

It assumes particles are solid spheres and does not explain particle interactions or forces accurately.

5.2.2.2 State Symbols

1.

What state symbol represents a solid substance in a chemical equation?

(s)

2.

What state symbol represents a liquid substance in a chemical equation?

(l)

3.

What state symbol represents a gaseous substance in a chemical equation?

(g)

4.

What state symbol represents a substance dissolved in water?

(aq)

5.

Why are state symbols included in chemical equations?

To show the physical state of substances in a reaction.

6.

Which state symbol would be used for sodium chloride dissolved in water?

(aq)

5.2.2.3 Properties of Ionic Compounds

1.

Why do ionic compounds have giant ionic lattice structures?

They contain many oppositely charged ions arranged in a regular lattice.

2.

Why do ionic compounds have high melting and boiling points?

Strong electrostatic forces between ions.

3.

Why are strong electrostatic forces present in ionic compounds?

Because the ions have strong attractions between them.

4.

Why can ionic compounds conduct electricity when they are melted?

Ions are free to move when molten.

5.

Why can ionic compounds conduct electricity when dissolved in water?

Ions are free to move when dissolved in water.

6.

Why can solid ionic compounds not conduct electricity?

The ions cannot move because they are fixed in position.

5.2.2.4 Properties of Small Molecules

1.

Why do small molecular substances usually have low melting points?

Weak intermolecular forces require little energy to overcome.

2.

What type of forces exist between molecules in small molecular substances?

Weak intermolecular forces.

3.

Why are intermolecular forces weaker than covalent bonds?

Covalent bonds involve strong forces within molecules, while intermolecular forces are weaker forces between molecules.

4.

Why do larger molecules have higher boiling points than smaller molecules?

Larger molecules have stronger intermolecular forces.

5.

Why do small molecular substances not conduct electricity?

They do not contain charged particles that can move.

6.

What forces are overcome when a small molecular substance melts or boils?

Intermolecular forces.

5.2.2.5 Polymers

1.

What are polymers made from?

Small molecules called monomers.

2.

Why are polymer molecules described as very large molecules?

They contain thousands of atoms joined together.

3.

What type of bonds hold atoms together in polymer molecules?

Covalent bonds.

4.

Why are polymers usually solids at room temperature?

They contain long chains of atoms with strong covalent bonds.

5.

How can a polymer be identified from a bonding diagram?

A polymer has repeating units in its structure.

6.

How are polymers different from substances made of small molecules?

Polymers are very large molecules, while small molecules contain fewer atoms.

5.2.2.6 Giant Covalent Structures

1.

What is a giant covalent structure?

A structure containing many atoms joined by covalent bonds in a giant network.

2.

Why do giant covalent structures have very high melting points?

Many strong covalent bonds require large amounts of energy to break.

3.

Why are strong covalent bonds present throughout giant covalent structures?

Covalent bonds extend throughout the whole structure.

4.

Which three substances are examples of giant covalent structures?

Diamond, graphite and graphene.

5.

Why does diamond have a very high melting point?

It contains many strong covalent bonds between carbon atoms.

6.

How can a giant covalent structure be recognised from a diagram?

A giant covalent structure has a large network of atoms joined by covalent bonds.

5.2.2.7 Properties of Metals and Alloys

1.

Why do metals usually have high melting and boiling points?

Strong metallic bonds require a lot of energy to break.

2.

How are atoms arranged in a pure metal?

Regular layers of metal atoms.

3.

Why can pure metals be bent and shaped?

Layers of atoms can slide over each other.

4.

Why are alloys harder than pure metals?

Different sized atoms disrupt the layers, making them harder to move.

5.

How does adding other metals affect the layers of atoms in a metal?

They make the layers less able to slide over each other.

6.

Why are alloys often used instead of pure metals?

Alloys are stronger and harder than pure metals.

5.2.2.8 Metals as Conductors

1.

Why do metals conduct electricity?

Metals contain delocalised electrons that can move.

2.

Which particles carry electrical charge through metals?

Delocalised electrons.

3.

Why do metals conduct thermal energy?

Delocalised electrons transfer energy through the metal.

4.

What is the role of delocalised electrons in metallic conduction?

They carry electrical charge and transfer thermal energy.

5.

Why are metals useful materials for electrical wires?

They are good electrical conductors because electrons can move freely.

6.

How does metallic bonding explain the conductivity of metals?

Metallic bonding allows electrons to move through the structure.

5.2.3.1 Diamond

1.

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

Four covalent bonds.

2.

What type of structure does diamond have?

A giant covalent structure.

3.

Why is diamond very hard?

Each carbon atom is strongly bonded to four others.

4.

Why does diamond have a high melting point?

Many strong covalent bonds require lots of energy to break.

5.

Why does diamond not conduct electricity?

It has no free electrons to carry charge.

6.

How does the structure of diamond explain its properties?

Its giant covalent structure explains its hardness and high melting point.

5.2.3.2 Graphite

1.

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

Three covalent bonds.

2.

What structure do carbon atoms form in graphite?

Layers of carbon atoms arranged in hexagonal structures.

3.

Why can graphite conduct electricity?

It has delocalised electrons that can move.

4.

Why is graphite soft?

Layers are held together by weak forces.

5.

Why can layers in graphite slide over each other?

Weak forces between layers allow them to slide.

6.

How does the structure of graphite explain its properties?

Its layered structure explains why it is soft and conducts electricity.

5.2.3.3 Graphene and Fullerenes

1.

What is graphene?

A single layer of graphite made of carbon atoms.

2.

How is graphene related to graphite?

It is one layer of graphite.

3.

Why is graphene useful in electronics?

It conducts electricity and is very strong.

4.

What are fullerenes?

Molecules made from carbon atoms arranged in hollow shapes.

5.

What is the structure of buckminsterfullerene (C₆₀)?

A spherical molecule containing 60 carbon atoms.

6.

What are carbon nanotubes used for?

Used in nanotechnology and strengthening materials.

Topic 2 Review

1.

How does bonding determine the properties of a substance?

Bonding and structure determine properties such as melting point, conductivity and hardness.

2.

What are the differences between ionic, covalent and metallic bonding?

Ionic bonding involves electron transfer, covalent bonding involves electron sharing, and metallic bonding involves positive ions and delocalised electrons.

3.

Why do ionic compounds conduct electricity when molten but not when solid?

Solid ionic compounds cannot conduct because ions cannot move; molten ionic compounds conduct because ions are free to move.

4.

How do the structures of diamond and graphite cause their different properties?

Diamond has a giant structure with four bonds per carbon atom, while graphite has layers and delocalised electrons.

5.

Why are metals good conductors of electricity and thermal energy?

Metals conduct because they contain mobile delocalised electrons.

6.

How do scientists use knowledge of bonding and structure to create new materials?

Scientists use bonding and structure knowledge to design materials with specific properties.

5.3.1.1 Conservation of Mass and Balanced Chemical Equations

1.

What does the law of conservation of mass state?

Mass is not created or destroyed during a chemical reaction.

2.

Why is the mass of products equal to the mass of reactants in a chemical reaction?

Atoms are rearranged during reactions, but no atoms are lost or gained.

3.

Why must chemical equations be balanced?

Chemical equations must be balanced to show the same number of atoms of each element on both sides.

4.

What do the numbers in front of chemical formulas in a balanced equation represent?

The numbers in front of chemical formulas show the ratio of reacting substances and products.

5.

What is the difference between a multiplier and a subscript in a chemical equation?

A multiplier is the number in front of a formula showing the number of molecules or moles; a subscript shows the number of atoms of an element in a formula.

6.

Balance the equation: H₂ + O₂ → H₂O.

H₂ + O₂ → H₂O becomes 2H₂ + O₂ → 2H₂O.

5.3.1.2 Relative Formula Mass

1.

What is meant by relative formula mass (Mr)?

The sum of the relative atomic masses of all atoms in a compound.

2.

How is the relative formula mass of a compound calculated?

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

3.

Which values from the periodic table are needed to calculate Mr?

Relative atomic masses from the periodic table.

4.

Calculate the Mr of H₂O.

H₂O: Mr = (2 × 1) + 16 = 18.

5.

Why does the total Mr of reactants equal the total Mr of products in a balanced equation?

The total mass of reactants equals the total mass of products because mass is conserved.

6.

How can relative formula mass be used to calculate the percentage by mass of an element in a compound?

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

5.3.1.3 Mass Changes When a Reactant or Product is a Gas

1.

Why can a chemical reaction appear to have a change in mass?

A reaction can appear to change mass because gases can enter or leave the system.

2.

Why does a metal gain mass when it reacts with oxygen?

The metal reacts with oxygen from the air, adding oxygen atoms to the metal.

3.

What happens to carbon dioxide gas produced during thermal decomposition of a metal carbonate?

Carbon dioxide gas escapes into the surroundings.

4.

Why can an open system give an inaccurate measurement of mass during a reaction?

Gas entering or leaving an open system changes the measured mass.

5.

Explain why the mass of a metal oxide is greater than the mass of the original metal.

The metal oxide contains the original metal plus oxygen atoms, increasing the mass.

6.

How can the particle model explain changes in mass during reactions involving gases?

Gas particles can move into or out of the system, causing changes in measured mass.

5.3.1.4 Chemical Measurements

1.

What is meant by uncertainty in a measurement?

Uncertainty is the amount by which a measurement may differ from the true value.

2.

Why does every measurement have some uncertainty?

Every measurement has uncertainty because equipment and human measurements have limitations.

3.

How can the range of measurements be used to estimate uncertainty?

Uncertainty can be estimated by: (highest value − lowest value) ÷ 2.

4.

What does the spread of results show about an experiment?

The spread of results shows how close together the measurements are.

5.

How can repeated measurements improve confidence in a result?

Repeated measurements improve confidence by allowing a mean to be calculated.

6.

What does a smaller range of results suggest about the reliability of measurements?

A smaller range suggests measurements are more precise and reliable.

5.3.2.1 Moles (HT Only)

1.

What is the unit used to measure chemical amounts?

The mole (mol).

2.

What is the mass of one mole of a substance equal to?

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

3.

What is the Avogadro constant?

6.02 × 10²³ particles per mole.

4.

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

6.02 × 10²³ particles.

5.

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

Moles = mass ÷ Mr. Water: 18 ÷ 18 = 1 mol.

6.

Calculate the mass of 0.5 moles of chlorine gas.

Mr of Cl₂ = 35.5 × 2 = 71. Mass = 0.5 × 71 = 35.5 g.

5.3.2.2 Amounts of Substances in Equations (HT Only)

1.

What information does a balanced symbol equation show about reacting substances?

A balanced equation shows the ratio of moles of reactants and products.

2.

How can a balanced chemical equation be interpreted in terms of moles?

The coefficients in a balanced equation represent the mole ratio.

3.

In the equation Mg + 2HCl → MgCl₂ + H₂, how many moles of hydrochloric acid react with one mole of magnesium?

Mg + 2HCl → MgCl₂ + H₂. One mole of magnesium reacts with 2 moles of hydrochloric acid.

4.

How can you calculate the mass of a product from a balanced equation?

Calculate moles of the known substance, use the mole ratio, then convert the product moles into mass.

5.

Why are mole ratios important when calculating reacting masses?

Mole ratios show the exact amounts of substances reacting.

6.

Calculate the mass of aluminium oxide produced when 135g of aluminium is burned in air.

4Al + 3O₂ → 2Al₂O₃. Moles of Al = 135 ÷ 27 = 5 mol. Ratio Al:Al₂O₃ = 4:2, so moles of Al₂O₃ = 5 ÷ 2 = 2.5 mol. Mr Al₂O₃ = (2 × 27) + (3 × 16) = 102. Mass = 2.5 × 102 = 255 g.

5.3.2.3 Using Moles to Balance Equations (HT Only)

1.

Why are moles used to balance chemical equations?

Moles are used because equations must show the correct ratios of reacting particles.

2.

What must masses be converted into before balancing by calculation?

Masses must be converted into moles.

3.

What type of ratio is used to balance equations?

Mole ratios.

4.

Balance the chemical equation for photosynthesis using the following masses: 264 g carbon dioxide reacts with 108 g water to produce 180 g glucose and 192 g oxygen.

CO₂: 264 ÷ 44 = 6 mol. H₂O: 108 ÷ 18 = 6 mol. Glucose: 180 ÷ 180 = 1 mol. O₂: 192 ÷ 32 = 6 mol. Equation: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂.

5.

Balance the chemical equation for the complete combustion of methane using the following masses: 16 g methane reacts with 64 g oxygen to produce 44 g carbon dioxide and 36 g water.

CH₄: 16 ÷ 16 = 1 mol. O₂: 64 ÷ 32 = 2 mol. CO₂: 44 ÷ 44 = 1 mol. H₂O: 36 ÷ 18 = 2 mol. Equation: CH₄ + 2O₂ → CO₂ + 2H₂O.

6.

Balance a chemical equation between magnesium and oxygen using the following masses: 24 g magnesium reacts with 16 g oxygen to produce 40 g magnesium oxide.

Mg: 24 ÷ 24 = 1 mol. O₂: 16 ÷ 32 = 0.5 mol. MgO: 40 ÷ 40 = 1 mol. Ratio = 2:1:2. Equation: 2Mg + O₂ → 2MgO.

5.3.2.4 Limiting Reactants (HT Only)

1.

What is meant by a limiting reactant?

The reactant that is completely used up first.

2.

Why is one reactant often used in excess during a chemical reaction?

It ensures the reaction uses the available reactants efficiently and determines the maximum product formed.

3.

How does a limiting reactant affect the amount of product formed?

It limits the amount of product that can be produced.

4.

What happens to the excess reactant after the limiting reactant is used up?

The excess reactant remains after the reaction has finished.

5.

How can moles be used to identify the limiting reactant?

Compare the number of moles of reactants with the balanced equation ratio.

6.

Explain why increasing the amount of the limiting reactant increases the amount of product made.

Increasing the amount of the limiting reactant allows more product to form.

5.3.2.5 Concentration of Solutions

1.

What is meant by the concentration of a solution?

The amount of solute dissolved in a given volume of solution.

2.

What units can be used to measure concentration?

g/dm³ or mol/dm³.

3.

How is concentration calculated using mass and volume?

Concentration = mass ÷ volume.

4.

Calculate the concentration of a solution containing 20 g of solute in 2 dm³ of solution.

Concentration = 20 ÷ 2 = 10 g/dm³.

5.

How does increasing the mass of solute affect concentration?

Increasing the mass of solute increases concentration.

6.

How does increasing the volume of solution affect concentration?

Increasing the volume of solution decreases concentration.

Topic 3 Review

1.

Explain how the law of conservation of mass applies to chemical reactions.

The law of conservation of mass means atoms are rearranged but total mass remains constant.

2.

Describe how balanced equations represent the quantities of substances in reactions.

Balanced equations show the quantities and ratios of reactants and products.

3.

Explain how relative formula mass is calculated and used.

Relative formula mass is calculated by adding atomic masses and is used in chemical calculations.

4.

Explain how moles allow chemists to measure amounts of substances.

Moles allow chemists to measure and compare amounts of substances.

5.

Describe how limiting reactants affect the amount of product formed.

The limiting reactant determines the maximum amount of product formed.

6.

Explain how concentration calculations are used in chemical reactions.

Concentration calculations show the amount of solute present in a solution.

5.4.1.1 Metal Oxides

1.

What are formed when metals react with oxygen?

Metal oxides.

2.

Why are reactions between metals and oxygen classed as oxidation reactions?

They involve the gain of oxygen by the metal.

3.

In terms of oxygen, what is oxidation?

Oxidation is the gain of oxygen.

4.

In terms of oxygen, what is reduction?

Reduction is the loss of oxygen.

5.

What happens to a metal atom when it reacts with oxygen?

The metal atom loses electrons and forms a positive ion.

6.

Explain why the formation of a metal oxide is an oxidation reaction.

The metal gains oxygen, so it is oxidised.

5.4.1.2 The Reactivity Series

1.

What is meant by the reactivity of a metal?

The ability of a metal to undergo reactions.

2.

Why do metals form positive ions when they react?

Metals lose electrons to form positive ions.

3.

What is the reactivity series used to show?

The order of metals from most reactive to least reactive.

4.

Which metal is more reactive: a metal that displaces another metal from its compound or the metal being displaced?

The metal that displaces another metal from its compound is more reactive.

5.

How can reactions with water and dilute acids be used to compare metal reactivity?

Reactions with water and acids show how easily metals react and allow comparison of reactivity.

6.

Explain why a more reactive metal can displace a less reactive metal from its compound.

A more reactive metal removes oxygen or other elements from a less reactive metal compound.

5.4.1.3 Extraction of Metals and Reduction

1.

Why are some metals found as compounds in the Earth rather than as pure metals?

Some metals are found as compounds because they have reacted with other elements.

2.

Which type of metals are found naturally as the metal itself?

Metals below carbon in the reactivity series can be found naturally as the metal.

3.

Why can metals below carbon in the reactivity series be extracted using carbon?

Carbon can remove oxygen from less reactive metal oxides.

4.

What does reduction mean in terms of oxygen?

Reduction is the loss of oxygen.

5.

In the extraction of a metal oxide using carbon, which substance is reduced?

The metal oxide is reduced.

6.

Explain why more reactive metals need different extraction methods.

More reactive metals cannot be extracted using carbon because they form stronger bonds with oxygen.

5.4.1.4 Oxidation and Reduction in Terms of Electrons (HT Only)

1.

What is oxidation in terms of electrons?

Oxidation is the loss of electrons.

2.

What is reduction in terms of electrons?

Reduction is the gain of electrons.

3.

What happens to electrons during oxidation?

Electrons are lost during oxidation.

4.

What happens to electrons during reduction?

Electrons are gained during reduction.

5.

How can half equations show oxidation and reduction?

Half equations show electrons being lost or gained.

6.

Identify which substance is oxidised and which is reduced in a given redox reaction.

The substance losing electrons is oxidised; the substance gaining electrons is reduced.

5.4.2.1 Reactions of Acids with Metals

1.

What products are formed when acids react with metals?

A salt and hydrogen gas.

2.

How can you measure the rate of reaction of a metal and acid reaction?

Measure the volume of gas produced over time or measure mass loss over time.

3.

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

Hydrogen.

4.

Why are reactions between acids and metals considered redox reactions?

Electrons are transferred between substances.

5.

Which substance is oxidised when a metal reacts with an acid?

The metal is oxidised.

6.

Explain how electrons are transferred during a reaction between magnesium and hydrochloric acid.

Magnesium loses electrons to form Mg²⁺ ions, while hydrogen ions gain electrons to form hydrogen gas.

5.4.2.2 Neutralisation of Acids and Salt Production

1.

What is produced when an acid reacts with an alkali?

A salt and water.

2.

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

A salt, water and carbon dioxide.

3.

Which acid produces chloride salts?

Hydrochloric acid.

4.

Which acid produces nitrate salts?

Nitric acid.

5.

Which acid produces sulfate salts?

Sulfuric acid.

6.

Explain how the reactants determine the salt produced in a neutralisation reaction.

The acid and reactant determine the ions present in the salt produced.

5.4.2.3 Soluble Salts

1.

How can soluble salts be made from acids?

React an acid with a suitable metal, metal oxide, metal hydroxide or metal carbonate.

2.

Why is excess insoluble solid added during salt preparation?

To ensure all the acid reacts.

3.

Why is the excess solid filtered off after reacting with an acid?

To remove the unreacted excess solid.

4.

How are salt crystals produced from a salt solution?

Evaporate some water and allow crystals to form.

5.

What is the purpose of evaporating water from a salt solution?

Water is removed so the salt becomes more concentrated and crystals form.

6.

Describe how to prepare a pure, dry sample of a soluble salt from an insoluble oxide or carbonate.

Add excess insoluble solid to acid, filter to remove excess solid, evaporate water and dry the salt crystals.

5.4.2.4 The pH Scale and Neutralisation

1.

What ion do acids produce in aqueous solutions?

Hydrogen ions (H⁺).

2.

What ion do alkalis produce in aqueous solutions?

Hydroxide ions (OH⁻).

3.

What does the pH scale measure?

The concentration of hydrogen ions in a solution.

4.

What pH value represents a neutral solution?

pH 7.

5.

What pH values do acidic and alkaline solutions have?

Acids have pH values below 7; alkalis have pH values above 7.

6.

Explain what happens when hydrogen ions react with hydroxide ions.

Hydrogen ions react with hydroxide ions to form water.

5.4.2.5 Strong and Weak Acids (HT Only)

1.

What is a strong acid?

An acid that completely ionises in water.

2.

What is a weak acid?

An acid that only partially ionises in water.

3.

What happens to strong acids when dissolved in water?

Strong acids release all their hydrogen ions.

4.

What happens to weak acids when dissolved in water?

Weak acids only release some hydrogen ions.

5.

If the pH of a solution decreases from pH 5 to pH 2, how does the hydrogen ion concentration change?

The hydrogen ion concentration increases by 1000 times.

6.

Explain the difference between a strong acid and a concentrated acid.

Strength describes ionisation; concentration describes the amount of acid dissolved in a solution.

5.4.3.1 The Process of Electrolysis

1.

What is electrolysis?

The decomposition of an ionic compound using electricity.

2.

What is an electrolyte?

A substance containing free-moving ions that conducts electricity.

3.

Why can molten ionic compounds conduct electricity?

Ions are free to move and carry charge.

4.

Which electrode do positive ions move towards during electrolysis?

The cathode.

5.

Which electrode do negative ions move towards during electrolysis?

The anode.

6.

Explain how electrolysis causes ions to form elements.

Electricity causes ions to gain or lose electrons, forming elements.

5.4.3.2 Electrolysis of Molten Ionic Compounds

1.

What products are formed when molten ionic compounds are electrolysed?

The metal and non-metal elements.

2.

What is produced at the cathode during electrolysis of molten lead bromide?

Lead.

3.

What is produced at the anode during electrolysis of molten lead bromide?

Bromine.

4.

Why must molten ionic compounds contain free-moving ions?

Ions must be free to move to carry charge.

5.

Why are inert electrodes used in electrolysis?

They do not react with the products formed.

6.

Predict the products formed when a binary ionic compound is electrolysed in the molten state.

Positive ions form metals at the cathode; negative ions form non-metals at the anode.

5.4.3.3 Using Electrolysis to Extract Metals

1.

When is electrolysis used to extract metals?

For metals more reactive than carbon.

2.

Why cannot very reactive metals be extracted using carbon?

They cannot be reduced by carbon because they are too strongly bonded to oxygen.

3.

Why is a mixture of aluminium oxide and cryolite used in aluminium extraction?

Cryolite lowers the melting point of aluminium oxide and reduces energy costs.

4.

What metal is extracted by electrolysis of aluminium oxide?

Aluminium.

5.

Why does electrolysis require large amounts of energy?

Large amounts of energy are needed to melt compounds and move ions.

6.

Explain why the positive electrode in aluminium extraction must be replaced regularly.

The positive electrode reacts with oxygen to form carbon dioxide and wears away.

5.4.3.4 Electrolysis of Aqueous Solutions

1.

What determines which ions are discharged during electrolysis of an aqueous solution?

The reactivity of the ions and their position in the reactivity series.

2.

What gas is produced at the cathode if the metal is more reactive than hydrogen?

Hydrogen gas.

3.

What gas is usually produced at the anode during electrolysis?

Oxygen gas.

4.

What happens at the anode if halide ions are present?

Halogen gases are produced.

5.

Why can water molecules affect the products formed during electrolysis?

Water molecules can provide hydrogen and hydroxide ions.

6.

Predict the products formed when an aqueous ionic solution is electrolysed.

Predict products by comparing the reactivity of ions present.

5.4.3.5 Representation of Reactions at Electrodes as Half Equations (HT Only)

1.

What happens to positive ions at the cathode?

Positive ions gain electrons at the cathode.

2.

Why are reactions at the cathode reductions?

They are reduction reactions because electrons are gained.

3.

What happens to negative ions at the anode?

Negative ions lose electrons at the anode.

4.

Why are reactions at the anode oxidations?

They are oxidation reactions because electrons are lost.

5.

What information does a half equation show?

A half equation shows electrons transferred during a reaction.

6.

Write a half equation to represent the reaction at an electrode during electrolysis.

Example: Mg²⁺ + 2e⁻ → Mg.

Topic 4 Review

1.

Explain how the reactivity series allows scientists to predict metal reactions.

The reactivity series shows how easily metals react and allows predictions of displacement reactions.

2.

Describe the difference between oxidation and reduction.

Oxidation is loss of electrons or gain of oxygen; reduction is gain of electrons or loss of oxygen.

3.

Explain how metals can be extracted from their compounds.

Metals can be extracted by reduction with carbon or by electrolysis depending on their reactivity.

4.

Describe how acids react with metals, bases and carbonates.

Acids react with metals to produce salts and hydrogen, with bases to produce salts and water, and with carbonates to produce salts, water and carbon dioxide.

5.

Explain how electrolysis separates ionic compounds into elements.

Electrolysis separates ionic compounds by moving ions to electrodes where they gain or lose electrons.

6.

Compare extraction of metals using carbon reduction and electrolysis.

Carbon reduction is used for less reactive metals; electrolysis is needed for more reactive metals.

5.5.1.1 Energy Transfer During Exothermic and Endothermic Reactions

1.

What happens to energy in a chemical reaction according to the law of conservation of energy?

Energy cannot be created or destroyed; it is transferred between stores.

2.

What is an exothermic reaction?

A reaction that transfers energy to the surroundings.

3.

What happens to the temperature of the surroundings during an exothermic reaction?

The temperature of the surroundings increases.

4.

Give two examples of exothermic reactions.

Combustion and neutralisation.

5.

What is an endothermic reaction?

A reaction that takes in energy from the surroundings.

6.

What happens to the temperature of the surroundings during an endothermic reaction?

The temperature of the surroundings decreases.

5.5.1.2 Reaction Profiles

1.

What must happen between reacting particles for a chemical reaction to occur?

Particles must collide with enough energy and the correct orientation.

2.

What is activation energy?

The minimum amount of energy needed for a reaction to occur.

3.

What does a reaction profile show?

The energy changes during a chemical reaction.

4.

How can you identify an exothermic reaction from a reaction profile?

An exothermic reaction has products at a lower energy level than the reactants.

5.

Draw a labelled reaction profile for the endothermic reaction A + B → C + D, showing the reactants, products, activation energy and overall energy change.

Endothermic profile: products are higher in energy than reactants; activation energy is shown as the energy needed to start the reaction; overall energy change is positive.

6.

Draw a labelled reaction profile for the exothermic reaction A + B → C + D, showing the reactants, products, activation energy and overall energy change.

Exothermic profile: products are lower in energy than reactants; activation energy is shown as the energy needed to start the reaction; overall energy change is negative.

5.5.1.3 The Energy Change of Reactions (HT Only)

1.

Why is energy needed during a chemical reaction?

Energy is needed to break existing bonds and form new bonds.

2.

Is energy absorbed or released when chemical bonds are broken?

Energy is absorbed when chemical bonds are broken.

3.

Is energy absorbed or released when new chemical bonds are formed?

Energy is released when new chemical bonds are formed.

4.

In an exothermic reaction, is more energy released forming bonds or absorbed breaking bonds?

More energy is released when bonds form than is absorbed when bonds break.

5.

In an endothermic reaction, is more energy absorbed breaking bonds or released forming bonds?

More energy is absorbed breaking bonds than is released forming bonds.

6.

How can bond energies be used to calculate the overall energy change of a reaction?

Overall energy change = energy needed to break bonds − energy released forming bonds.

Topic 5 Review

1.

Explain the difference between exothermic and endothermic reactions.

Exothermic reactions release energy to the surroundings; endothermic reactions absorb energy from the surroundings.

2.

Explain how energy changes occur when bonds are broken and formed.

Energy is absorbed when bonds break and released when bonds form.

3.

Describe how reaction profiles can be used to identify exothermic and endothermic reactions.

Reaction profiles show activation energy and whether energy is released or absorbed.

4.

Explain why activation energy is needed for chemical reactions.

Activation energy is needed to start reactions by allowing particles to collide successfully.

5.

Compare the energy changes in exothermic and endothermic reactions.

Exothermic reactions have a negative energy change; endothermic reactions have a positive energy change.

6.

Explain how energy changes in chemical reactions are used in everyday applications.

Energy changes are used in applications such as fuels, hand warmers and cooling packs.

5.6.1.1 Calculating rates of reactions

1.

How is the mean rate of a reaction calculated using the quantity of reactant used?

Mean rate = quantity of reactant used ÷ time taken.

2.

How is the mean rate of a reaction calculated using the quantity of product formed?

Mean rate = quantity of product formed ÷ time taken.

3.

What units can be used for the rate of reaction when measuring mass change?

g/s or g/min.

4.

What units can be used for the rate of reaction when measuring gas volume produced?

cm³/s or cm³/min.

5.

How can a graph showing the amount of product formed over time be used to determine the rate of reaction?

The gradient of the graph shows the rate; a steeper gradient means a faster reaction.

6.

How is the gradient of a tangent to a reaction graph used to calculate the rate of reaction at a specific time?

Draw a tangent at the point and calculate its gradient: change in amount ÷ change in time.

5.6.1.2 Factors which affect the rates of chemical reactions

1.

What five factors can affect the rate of a chemical reaction?

Concentration, pressure, surface area, temperature and catalysts.

2.

How does increasing the concentration of reactants in solution affect the rate of reaction?

Increasing concentration increases the rate because there are more particles in the same volume.

3.

How does decreasing the pressure of reacting gases affect the rate of reaction?

Decreasing pressure decreases the rate because gas particles collide less often.

4.

How does increasing the surface area of a solid reactant affect the rate of reaction?

Increasing surface area increases the rate because more particles are exposed for collisions.

5.

How does increasing the temperature affect the rate of reaction?

Increasing temperature increases the rate because particles move faster (due to more kinetic energy) and collide more successfully.

6.

How do catalysts affect the rate of chemical reactions?

Catalysts increase the rate by providing an alternative pathway with lower activation energy.

5.6.1.3 Collision theory and activation energy

1.

What does collision theory state about when chemical reactions can occur?

Particles must collide with enough energy and the correct orientation for a reaction to occur.

2.

What is activation energy?

The minimum energy needed for particles to react.

3.

Why does increasing the concentration of reactants increase the rate of reaction?

More particles are present, causing more frequent successful collisions.

4.

Why does increasing the temperature increase the rate of reaction?

Particles move faster and have more energy, increasing successful collisions.

5.

Why does increasing the surface area of a solid reactant increase the rate of reaction?

More surface particles are exposed, increasing collision frequency.

6.

Explain how pressure affects the rate of reaction between reacting gases using collision theory.

Higher pressure pushes gas particles closer together, causing more frequent collisions.

5.6.1.4 Catalysts

1.

What is a catalyst?

A substance that increases the rate of reaction without being used up.

2.

Why are catalysts not included in chemical equations?

Catalysts do not take part in the overall reaction and are not changed.

3.

How do catalysts increase the rate of a reaction?

They provide an alternative reaction pathway with lower activation energy.

4.

What happens to the activation energy when a catalyst is used?

The activation energy decreases.

5.

Why do different chemical reactions require different catalysts?

Different reactions require different catalysts because they involve different reactants and mechanisms.

6.

How do enzymes act as catalysts in biological systems?

Enzymes are biological catalysts that speed up reactions in living organisms.

5.6.2.1 Reversible reactions

1.

What is a reversible reaction?

A reaction where products can react to form the original reactants.

2.

How are reversible reactions represented using a symbol equation?

They are represented using a reversible arrow (⇌).

3.

What happens to the products in a reversible reaction?

Products can react to form reactants again.

4.

How can the direction of a reversible reaction be changed?

By changing the conditions such as temperature, pressure or concentration.

5.

Give an example of a condition that can affect the direction of a reversible reaction.

Changing temperature, pressure or concentration.

6.

Why can reversible reactions produce both reactants and products?

They can proceed in both directions, producing both reactants and products.

5.6.2.2 Energy changes and reversible reactions

1.

If a reaction is exothermic in one direction, what type of reaction is it in the opposite direction?

It is endothermic in the opposite direction.

2.

How does the energy change compare in the forward and reverse reactions?

The energy change is equal but opposite.

3.

What happens to energy during an exothermic reaction?

Energy is transferred to the surroundings.

4.

What happens to energy during an endothermic reaction?

Energy is taken in from the surroundings.

5.

Why does reversing a reaction reverse the energy change?

The reverse reaction breaks and forms bonds in the opposite way.

6.

A forward reaction releases energy. Is the reverse reaction exothermic or endothermic?

Endothermic.

5.6.2.3 Equilibrium

1.

What is meant by equilibrium in a reversible reaction?

A state where the forward and reverse reactions occur at the same rate.

2.

What conditions are needed for equilibrium to be reached?

A closed system and constant conditions.

3.

Why must reactants and products be prevented from escaping for equilibrium to occur?

Reactants and products must not escape or the amounts will change.

4.

At equilibrium, how do the rates of the forward and reverse reactions compare?

The forward and reverse reactions occur at the same rate.

5.

Why is equilibrium called dynamic?

It is dynamic because reactions continue happening.

6.

What happens to the amounts of reactants and products once equilibrium is reached?

The amounts of reactants and products remain constant.

5.6.2.4 The effect of changing conditions on equilibrium (HT only)

1.

What happens to an equilibrium system when a change is made to its conditions?

The equilibrium position shifts to oppose the change.

2.

What principle is used to predict the effect of changes on equilibrium?

Le Chatelier's Principle.

3.

What does Le Chatelier's Principle state?

If conditions are changed, the equilibrium moves to reduce the effect of that change.

4.

How does changing conditions affect the position of equilibrium?

Temperature, pressure and concentration changes shift equilibrium.

5.

Why does a system at equilibrium respond to changes in conditions?

The system responds to restore equilibrium.

6.

How can equilibrium be shifted to increase the amount of a desired product?

Increase the amount of desired product by changing conditions to favour the forward reaction.

5.6.2.5 The effect of changing concentration (HT only)

1.

What happens when the concentration of a reactant is increased in an equilibrium system?

The equilibrium shifts to use up the increased reactant.

2.

Why does increasing the concentration of a reactant produce more products?

More reactant particles are available, increasing successful collisions and producing more products.

3.

What happens when the concentration of a product is decreased in an equilibrium system?

The equilibrium shifts to replace the removed product.

4.

Why do concentrations of all substances change after the concentration of one substance is altered?

The system adjusts to restore the balance between reactants and products.

5.

How can changing concentration affect the position of equilibrium?

Increasing reactant concentration shifts equilibrium towards products; increasing product concentration shifts it towards reactants.

6.

For the equilibrium reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), explain what happens to the position of equilibrium when the concentration of hydrogen is increased.

N₂(g) + 3H₂(g) ⇌ 2NH₃(g): increasing hydrogen concentration shifts equilibrium to the right, producing more ammonia.

5.6.2.6 The effect of temperature changes on equilibrium (HT only)

1.

What happens to the amount of products at equilibrium when the temperature is increased for an endothermic reaction?

More products are formed because the equilibrium shifts in the endothermic direction.

2.

What happens to the amount of products at equilibrium when the temperature is increased for an exothermic reaction?

Less product is formed because the equilibrium shifts in the endothermic direction.

3.

What happens to the amount of products at equilibrium when the temperature is decreased for an endothermic reaction?

Less product is formed because the equilibrium shifts away from the endothermic direction.

4.

What happens to the amount of products at equilibrium when the temperature is decreased for an exothermic reaction?

More product is formed because the equilibrium shifts in the exothermic direction.

5.

Why does increasing temperature affect the position of equilibrium?

Temperature changes alter the balance between energy absorbed and released.

6.

For the (exothermic) equilibrium reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), predict how increasing the temperature affects the equilibrium position and explain why.

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) is exothermic, so increasing temperature shifts equilibrium left, reducing ammonia yield.

5.6.2.7 The effect of pressure changes on equilibrium (HT only)

1.

How does increasing pressure affect the equilibrium position of a gaseous reaction?

Increasing pressure shifts equilibrium towards the side with fewer gas molecules.

2.

Why does increasing pressure shift equilibrium towards the side with fewer gas molecules?

This reduces pressure because fewer gas particles occupy less space.

3.

How does decreasing pressure affect the equilibrium position of a gaseous reaction?

Decreasing pressure shifts equilibrium towards the side with more gas molecules.

4.

Why does decreasing pressure shift equilibrium towards the side with more gas molecules?

More gas molecules increase pressure and oppose the change.

5.

How can the symbol equation be used to predict the effect of pressure changes on equilibrium?

Count the number of gas molecules on each side of the equation.

6.

For the equilibrium reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), predict how increasing the pressure affects the equilibrium position and explain why.

N₂(g) + 3H₂(g) ⇌ 2NH₃(g): increasing pressure shifts equilibrium right because there are fewer gas molecules on the product side.

Topic 6 Review

1.

Explain how changing concentration, pressure, temperature and surface area affects the rate of a chemical reaction.

Increasing concentration, pressure, temperature and surface area increases reaction rate; catalysts also increase rate.

2.

Explain how collision theory can be used to explain why increasing temperature increases the rate of reaction.

Higher temperature gives particles more energy, causing more successful collisions.

3.

Explain how catalysts increase the rate of reaction and why they are not used up during a reaction.

Catalysts lower activation energy and are not used up during reactions.

4.

Describe the conditions needed for dynamic equilibrium and explain what happens when conditions are changed.

Dynamic equilibrium occurs in a closed system where forward and reverse reactions happen at equal rates.

5.

For the reaction N₂ + 3H₂ ⇌ 2NH₃, explain how changing temperature, pressure and concentration affects the equilibrium position.

Increasing pressure favours the side with fewer gas molecules; temperature changes favour the endothermic or exothermic direction; concentration changes oppose the change.

6.

Explain the difference between reversible reactions, equilibrium and irreversible reactions.

Reversible reactions can go both directions; equilibrium is when the forward and reverse rates are equal; irreversible reactions only go one direction.

5.7.1.1 Crude oil, hydrocarbons and alkanes

1.

What is crude oil?

Crude oil is a mixture of hydrocarbons formed from the remains of ancient organisms.

2.

Why is crude oil described as a finite resource?

Crude oil is finite because it takes millions of years to form and cannot be replaced quickly.

3.

What are hydrocarbons?

Hydrocarbons are compounds made only of hydrogen and carbon.

4.

What elements are found in hydrocarbon molecules?

Carbon and hydrogen.

5.

What is the general formula for alkanes?

CₙH₂ₙ₊₂.

6.

An alkane has the formula C₄H₁₀. Explain how this shows that it is an alkane.

C₄H₁₀ follows the formula CₙH₂ₙ₊₂ because 2(4)+2 = 10, so it is an alkane.

5.7.1.2 Fractional distillation and petrochemicals

1.

What process is used to separate crude oil into fractions?

Fractional distillation.

2.

Explain how fractional distillation separates hydrocarbons in crude oil.

Crude oil is heated and separated into fractions based on different boiling points.

3.

Why do different fractions condense at different temperatures?

Different hydrocarbons have different chain lengths and intermolecular forces, giving different boiling points.

4.

What are fractions from crude oil used to produce?

Fuels, solvents and chemicals for making products.

5.

Give two examples of fuels produced from crude oil.

Petrol and diesel.

6.

Explain why crude oil is useful as a feedstock for the petrochemical industry.

Crude oil contains many useful hydrocarbons that can be used to make petrochemicals.

5.7.1.3 Properties of hydrocarbons

1.

How does boiling point change as the size of hydrocarbon molecules increases?

Boiling point increases as hydrocarbon molecules get larger.

2.

How does viscosity change as the size of hydrocarbon molecules increases?

Viscosity increases as hydrocarbon molecules get larger.

3.

How does flammability change as the size of hydrocarbon molecules increases?

Flammability decreases as hydrocarbon molecules get larger.

4.

Why are small hydrocarbons useful as fuels?

Small hydrocarbons ignite easily and release useful energy.

5.

What are the products of complete combustion of a hydrocarbon?

Carbon dioxide and water.

6.

Write the balanced symbol equation for the complete combustion of methane (CH₄).

CH₄ + 2O₂ → CO₂ + 2H₂O.

5.7.1.4 Cracking and alkenes

1.

What is cracking?

Cracking is breaking large hydrocarbon molecules into smaller, more useful molecules.

2.

Why are hydrocarbons cracked?

To produce more useful fuels and chemicals from larger hydrocarbons.

3.

What two types of hydrocarbons are produced during cracking?

Alkanes and alkenes.

4.

What test is used to identify an alkene?

Bromine water.

5.

What colour change occurs when bromine water reacts with an alkene?

Orange/brown bromine water turns colourless.

6.

Explain why cracking is important for producing useful chemicals and fuels.

Cracking produces useful smaller hydrocarbons and alkenes needed for fuels and polymers.

Topic 7 Review

1.

Explain how crude oil is separated into useful fractions by fractional distillation.

Crude oil is separated by fractional distillation using differences in boiling points.

2.

Explain how the properties of hydrocarbons change as their molecules become larger.

Larger hydrocarbons have higher boiling points, higher viscosity and lower flammability.

3.

Describe the difference between alkanes and alkenes.

Alkanes have only single carbon-carbon bonds; alkenes contain at least one carbon-carbon double bond.

4.

Explain how cracking produces more useful hydrocarbons from larger molecules.

Cracking breaks large hydrocarbons into smaller alkanes and alkenes.

5.

Write and explain the equation for the complete combustion of a hydrocarbon.

Complete combustion produces carbon dioxide and water: CH₄ + 2O₂ → CO₂ + 2H₂O.

6.

Explain why hydrocarbons are important in modern life.

Hydrocarbons are important because they are used as fuels and as raw materials for making many products.

5.8.1.1 Pure substances

1.

What is meant by a pure substance in chemistry?

A substance containing only one element or compound.

2.

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

A pure substance contains only one substance; a mixture contains two or more substances not chemically bonded.

3.

How can melting point data be used to identify a pure substance?

A pure substance has a fixed melting point; impurities cause melting over a range of temperatures.

4.

How can boiling point data be used to identify a pure substance?

A pure substance has a fixed boiling point; impurities change the boiling point.

5.

Why do pure substances have fixed melting points and boiling points?

Pure substances contain particles that are all the same, so they change state at a fixed temperature.

6.

A substance melts over a range of temperatures rather than at one temperature. Explain what this suggests about the substance.

It suggests the substance is impure because impurities cause a range of melting temperatures.

5.8.1.2 Formulations

1.

What is a formulation?

A mixture designed as a useful product with specific amounts of components.

2.

Why are the components of a formulation mixed in carefully measured quantities?

To ensure the product has the correct properties and works effectively.

3.

Give two examples of products that are formulations.

Paint and medicines.

4.

Explain why a medicine is an example of a formulation.

Medicines contain carefully measured ingredients to provide the correct effect and dosage.

5.

How do formulations differ from simple mixtures?

Formulations contain carefully selected amounts of several substances, while simple mixtures do not have controlled compositions.

6.

A paint contains pigments, solvents and other chemicals. Explain why paint is a formulation.

Paint is a formulation because pigments, solvents and other chemicals are combined in specific quantities to produce the desired properties.

5.8.1.3 Chromatography

1.

What is chromatography used for?

To separate and identify substances in a mixture.

2.

What are the stationary phase and mobile phase in paper chromatography?

Stationary phase: the paper. Mobile phase: the solvent.

3.

Explain how paper chromatography separates substances in a mixture.

Different substances travel at different speeds depending on their solubility and attraction to the paper.

4.

What does an Rf value show in chromatography?

The Rf value shows the distance travelled by a substance compared with the solvent front.

5.

Calculate the Rf value of a substance that moves 4 cm when the solvent moves 8 cm.

Rf = distance travelled by substance ÷ distance travelled by solvent. Rf = 4 ÷ 8. Rf = 0.5

6.

A pure compound produces one spot on a chromatogram. Explain what this shows.

It shows the compound contains only one substance.

5.8.2.1 Test for hydrogen

1.

How is hydrogen gas tested for?

Place a burning splint at the gas sample.

2.

What happens when a burning splint is placed into hydrogen gas?

The gas burns with a squeaky pop.

3.

What observation confirms the presence of hydrogen?

A squeaky pop confirms hydrogen.

4.

Why must the burning splint be held at the open end of the test tube when testing hydrogen?

To prevent the gas mixture escaping and to make the test safer.

5.

Which gas produces a squeaky pop sound when tested with a burning splint?

Hydrogen.

6.

A gas burns with a pop sound when tested with a burning splint. Identify the gas.

The gas is hydrogen.

5.8.2.2 Test for oxygen

1.

How is oxygen gas tested for?

Place a glowing splint into the gas.

2.

What happens when a glowing splint is placed into oxygen?

The glowing splint relights.

3.

What observation confirms the presence of oxygen?

Relighting confirms oxygen.

4.

Why does oxygen relight a glowing splint?

Oxygen supports combustion.

5.

Which gas relights a glowing splint?

Oxygen.

6.

A gas causes a glowing splint to relight. Identify the gas.

The gas is oxygen.

5.8.2.3 Test for carbon dioxide

1.

How is carbon dioxide gas tested for?

Bubble the gas through limewater.

2.

What solution is used to test for carbon dioxide?

Limewater.

3.

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

Limewater turns cloudy or milky.

4.

What observation confirms the presence of carbon dioxide?

A cloudy appearance confirms carbon dioxide.

5.

Which gas turns limewater milky?

Carbon dioxide.

6.

A gas turns limewater cloudy. Identify the gas.

The gas is carbon dioxide.

5.8.2.4 Test for chlorine

1.

How is chlorine gas tested for?

Place damp litmus paper into the gas.

2.

What happens when damp litmus paper is placed into chlorine gas?

Damp litmus paper is bleached and turns white.

3.

What observation confirms the presence of chlorine?

Bleaching confirms chlorine.

4.

Why must the litmus paper be damp when testing chlorine?

Water allows chlorine to react and produce acidic substances that bleach the paper.

5.

Which gas bleaches damp litmus paper?

Chlorine.

6.

A gas turns damp litmus paper white. Identify the gas.

The gas is chlorine.

Topic 8 Review

1.

Explain how melting point and boiling point data can be used to identify pure substances.

Melting point and boiling point data can identify pure substances because pure substances have fixed temperatures.

2.

Explain how chromatography can be used to separate and identify substances in a mixture.

Chromatography separates substances because different substances move different distances through a stationary phase.

3.

Describe how to calculate an Rf value from a chromatogram.

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

4.

Explain how a formulation differs from a pure substance.

A formulation is a carefully designed mixture, while a pure substance contains only one substance.

5.

Describe the tests used to identify hydrogen, oxygen, carbon dioxide and chlorine.

Hydrogen gives a squeaky pop, oxygen relights a glowing splint, carbon dioxide turns limewater cloudy and chlorine bleaches damp litmus paper.

6.

Explain why chemical analysis is useful in areas such as forensic science and medicine.

Chemical analysis helps identify substances in areas such as forensic science, medicine and industry.

5.9.1.1 The proportions of different gases in the atmosphere

1.

What percentage of the Earth's atmosphere is nitrogen?

Approximately 78%.

2.

What percentage of the Earth's atmosphere is oxygen?

Approximately 21%.

3.

Name two other gases found in small proportions in the atmosphere.

Argon and carbon dioxide.

4.

What is the approximate ratio of nitrogen to oxygen in the atmosphere?

Approximately 4:1 nitrogen to oxygen.

5.

How have the proportions of gases in the atmosphere changed over the last 200 million years?

The proportions have changed due to processes such as photosynthesis, volcanic activity and human activity.

6.

A sample of air contains 80% nitrogen and 20% oxygen. Explain why this represents the composition of the Earth's atmosphere.

It is similar because Earth's atmosphere is mainly nitrogen and oxygen in these proportions.

5.9.1.2 The Earth's early atmosphere

1.

What gas was thought to make up most of the Earth's early atmosphere?

Carbon dioxide.

2.

How did volcanic activity contribute to the formation of the early atmosphere?

Volcanoes released gases including carbon dioxide, water vapour and nitrogen.

3.

What happened to water vapour released by volcanoes as the Earth cooled?

It condensed to form oceans.

4.

Why is there limited evidence about the Earth's early atmosphere?

There is limited direct evidence because the early atmosphere existed billions of years ago.

5.

Explain how carbon dioxide levels decreased when oceans formed.

Carbon dioxide dissolved in oceans and became locked into sedimentary rocks.

6.

Explain how the Earth's early atmosphere was different from the atmosphere today.

The early atmosphere had more carbon dioxide and little or no oxygen compared with today.

5.9.1.3 How oxygen increased

1.

Which organisms first produced oxygen in the atmosphere?

Plants and algae.

2.

What process produces oxygen during photosynthesis?

Photosynthesis.

3.

Write the word equation for photosynthesis.

Carbon dioxide + water → glucose + oxygen.

4.

How did photosynthesis change the amount of oxygen in the atmosphere?

Photosynthesis increased the amount of oxygen in the atmosphere.

5.

Why did an increase in oxygen allow animals to evolve?

More oxygen allowed aerobic organisms to evolve.

6.

Explain why plants and algae are important in maintaining oxygen levels in the atmosphere.

Plants and algae remove carbon dioxide and release oxygen through photosynthesis.

5.9.1.4 How carbon dioxide decreased

1.

How did plants and algae reduce the amount of carbon dioxide in the atmosphere?

Plants and algae removed carbon dioxide through photosynthesis.

2.

What process removes carbon dioxide from the atmosphere?

Photosynthesis.

3.

How were sedimentary rocks formed using carbon dioxide?

Carbon dioxide dissolved in oceans and formed carbonate compounds in sedimentary rocks.

4.

How did the formation of fossil fuels reduce atmospheric carbon dioxide?

Carbon dioxide was stored in fossil fuels formed from ancient organisms.

5.

Describe how limestone, coal, crude oil and natural gas formed.

Limestone formed from carbon-containing sediments; coal, crude oil and natural gas formed from buried remains of organisms.

6.

Explain why carbon dioxide levels decreased as the Earth developed.

Carbon dioxide levels decreased as it was removed by photosynthesis and stored in rocks and fossil fuels.

5.9.2.1 Greenhouse gases

1.

What are greenhouse gases?

Gases that absorb infrared radiation and contribute to the greenhouse effect.

2.

Name three greenhouse gases.

Carbon dioxide, methane and water vapour.

3.

Explain how greenhouse gases keep the Earth warm.

They absorb and re-radiate infrared radiation, keeping the Earth warm.

4.

What happens when greenhouse gases absorb long wavelength radiation?

They absorb the energy and transfer it back towards the Earth.

5.

Why is the greenhouse effect necessary for life on Earth?

It keeps the Earth warm enough for life.

6.

Explain the difference between short wavelength radiation and long wavelength radiation in the greenhouse effect.

Short wavelength radiation comes from the Sun; long wavelength radiation is emitted by the Earth.

5.9.2.2 Human activities which contribute to an increase in greenhouse gases

1.

Name two human activities that increase carbon dioxide levels.

Burning fossil fuels and deforestation.

2.

Name two human activities that increase methane levels.

Farming and decomposition of waste.

3.

How does burning fossil fuels increase carbon dioxide in the atmosphere?

Burning fossil fuels releases carbon that was stored underground as carbon dioxide.

4.

How does farming increase methane levels in the atmosphere?

Livestock and farming processes release methane.

5.

Why do scientists use peer-reviewed evidence when studying climate change?

Peer-reviewed evidence is checked by other scientists to improve reliability.

6.

Explain why some climate change predictions have uncertainty.

Predictions have uncertainty because future human activity and natural processes are difficult to predict.

5.9.2.3 Global climate change

1.

What is meant by global climate change?

Long-term changes in the average temperature and climate of Earth.

2.

What is the main cause of the recent increase in average global temperature?

Increased greenhouse gases from human activities.

3.

State four possible effects of global climate change.

Rising sea levels, changing weather patterns, loss of habitats and melting ice.

4.

How could rising temperatures affect sea levels?

Warmer temperatures cause ice to melt and oceans to expand, increasing sea levels.

5.

Explain why predicting climate change is difficult.

Climate systems are complex and affected by many variables.

6.

Discuss why global climate change has environmental and social impacts.

It affects ecosystems, agriculture, resources, health and human populations.

5.9.2.4 The carbon footprint and its reduction

1.

What is a carbon footprint?

The total amount of greenhouse gases released by a person, organisation or activity.

2.

What gases are included when calculating a carbon footprint?

Carbon dioxide and other greenhouse gases such as methane.

3.

Give two ways of reducing carbon dioxide emissions.

Using renewable energy and reducing fossil fuel use.

4.

Give two ways of reducing methane emissions.

Reducing livestock farming and improving waste management.

5.

Why can reducing greenhouse gas emissions be difficult?

Many activities depend on fossil fuels and changing them can be difficult.

6.

Explain how using renewable energy can reduce a carbon footprint.

Renewable energy produces little or no carbon dioxide during generation, reducing emissions.

5.9.3.1 Atmospheric pollutants from fuels

1.

What pollutants can be produced when fuels are burned?

Carbon monoxide, particulates, sulfur dioxide and nitrogen oxides.

2.

How is carbon monoxide produced during combustion?

It forms during incomplete combustion of fuels.

3.

How are particulates produced when fuels are burned?

They form when fuels burn incompletely.

4.

What pollutant is produced when sulfur impurities in fuels are burned?

Sulfur dioxide.

5.

What conditions cause carbon monoxide to form instead of carbon dioxide?

Limited oxygen causes incomplete combustion.

6.

Explain why burning fuels can release several different pollutants.

Fuels contain impurities and combustion conditions vary, producing different pollutants.

5.9.3.2 Properties and effects of atmospheric pollutants

1.

Why is carbon monoxide dangerous to humans?

Carbon monoxide reduces the blood's ability to carry oxygen.

2.

Why is carbon monoxide difficult to detect?

It is colourless and odourless.

3.

What problems are caused by sulfur dioxide?

Sulfur dioxide causes acid rain and respiratory problems.

4.

What problems are caused by oxides of nitrogen?

Nitrogen oxides contribute to acid rain and breathing problems.

5.

How do particulates affect human health?

Particulates can damage the lungs and cause health problems.

6.

Explain how atmospheric pollutants can cause environmental problems such as acid rain and global dimming.

Pollutants can cause environmental problems such as acid rain, global dimming and damage to ecosystems.

Topic 9 Review

1.

Explain how the composition of the Earth's atmosphere has changed over time.

Earth's atmosphere changed from being rich in carbon dioxide to mainly nitrogen and oxygen.

2.

Explain how photosynthesis increased oxygen levels and decreased carbon dioxide levels.

Photosynthesis removed carbon dioxide and released oxygen.

3.

Describe how greenhouse gases affect the temperature of the Earth.

Greenhouse gases absorb infrared radiation and keep Earth warm.

4.

Explain how human activities increase greenhouse gas levels and contribute to climate change.

Human activities such as burning fossil fuels and farming increase greenhouse gas levels.

5.

Describe the sources and effects of major atmospheric pollutants.

Major pollutants include carbon monoxide, sulfur dioxide, nitrogen oxides and particulates.

6.

Evaluate actions that could reduce the impact of human activities on the atmosphere.

Reducing emissions, using renewable energy and protecting natural carbon sinks can reduce atmospheric impacts.

5.10.1.1 Using the Earth's resources and sustainable development

1.

What are natural resources?

Natural resources are materials obtained from the Earth that humans use.

2.

Give two examples of products obtained from the Earth's resources.

Metals, crude oil, water and rocks.

3.

What is the difference between a finite resource and a renewable resource?

A finite resource is limited and will eventually run out; a renewable resource can be replaced naturally.

4.

Give one example of a natural product that can be replaced by an agricultural or synthetic product.

Wood can be replaced by synthetic materials.

5.

What is meant by sustainable development?

Sustainable development means meeting the needs of the present without preventing future generations meeting their needs.

6.

Explain why reducing the use of finite resources is important for future generations.

Reducing finite resource use helps conserve supplies and protects resources for future generations.

5.10.1.2 Potable water

1.

What is meant by potable water?

Water that is safe to drink.

2.

Why is potable water not considered chemically pure water?

It contains dissolved substances and minerals, so it is not chemically pure.

3.

What are the two main requirements for drinking water to be safe?

It must contain no harmful microorganisms and have safe levels of dissolved substances.

4.

Describe the steps used to produce potable water from fresh water supplies.

Filter the water to remove solids and sterilise it to kill microorganisms.

5.

Name three substances or methods that can be used to sterilise water.

Chlorine, ozone and ultraviolet light.

6.

Explain why desalination is used and why it requires large amounts of energy.

Desalination removes salt from seawater but requires large amounts of energy.

5.10.1.3 Waste water treatment

1.

Why does waste water need to be treated before being released into the environment?

To remove harmful substances before releasing water back into the environment.

2.

What types of waste are removed during sewage treatment?

Organic matter, solids and microorganisms.

3.

What happens during the screening and grit removal stage of sewage treatment?

Large solids are removed by screening and grit is removed by settling.

4.

What is produced during sedimentation of sewage?

Sewage sludge.

5.

What happens during anaerobic digestion of sewage sludge?

Anaerobic bacteria break down sludge and produce methane.

6.

Explain why aerobic biological treatment is used during waste water treatment.

Aerobic biological treatment uses microorganisms to break down organic waste.

5.10.1.4 Alternative methods of extracting metals (HT only)

1.

Why are alternative methods of extracting metals needed?

Traditional extraction methods may use large amounts of energy and damage the environment.

2.

What is phytomining?

Phytomining is using plants to absorb metal compounds from the ground.

3.

How does phytomining extract metal compounds from the ground?

Plants absorb metal ions and are then burned to produce ash containing metal compounds.

4.

What is bioleaching?

Bioleaching uses bacteria to produce solutions containing metal ions.

5.

How can copper be extracted from a solution of copper compounds after bioleaching?

Copper can be extracted from the solution by displacement or electrolysis.

6.

Evaluate one advantage and one disadvantage of using biological methods to extract metals.

Advantage: uses less energy. Disadvantage: slower than traditional methods.

5.10.2.1 Life cycle assessment

1.

What is a life cycle assessment (LCA)?

A life cycle assessment evaluates the environmental impacts of a product throughout its life.

2.

What stages of a product's life are considered in a life cycle assessment?

Raw material extraction, manufacture, transport, use and disposal.

3.

Why are energy use and resource use easier to quantify than pollution effects?

Energy and resource use can be measured more easily than effects such as pollution.

4.

Why are life cycle assessments not completely objective?

They involve choices about which factors are included and how important they are.

5.

How could a life cycle assessment be misused?

Results could be selected to support a particular viewpoint.

6.

Explain how a life cycle assessment could be used to compare plastic and paper shopping bags.

An LCA compares the resources used, energy needed and environmental impacts of plastic and paper bags throughout their lifetimes.

5.10.2.2 Ways of reducing the use of resources

1.

What are the three main ways of reducing the use of resources?

Reduce, reuse and recycle.

2.

Why does recycling reduce the use of limited resources?

Recycling reduces the need for extracting new raw materials.

3.

How can glass bottles be reused?

Glass bottles can be cleaned and used again.

4.

How are metals recycled?

Metals are collected, melted and reshaped into new products.

5.

Why does recycling metals reduce the need for mining?

Recycling metals reduces the need to mine new ores.

6.

Explain why reducing, reusing and recycling materials helps protect the environment.

Reducing, reusing and recycling conserve resources and reduce environmental damage.

Topic 10 Review

1.

Explain why sustainable development is important when using Earth's resources.

Sustainable development ensures resources are used without preventing future generations from accessing them.

2.

Describe how fresh water can be treated to make it potable.

Fresh water is treated by filtration and sterilisation to make it potable.

3.

Explain the processes involved in waste water treatment.

Waste water treatment removes solids, organic matter and harmful microorganisms.

4.

Compare traditional metal extraction methods with phytomining and bioleaching.

Traditional extraction uses methods such as carbon reduction or electrolysis; phytomining and bioleaching use biological processes.

5.

Explain how life cycle assessments are used to compare the environmental impact of products.

Life cycle assessments compare environmental impacts at different stages of a product's life.

6.

Evaluate different ways of reducing the use of Earth's limited resources.

Reducing, reusing and recycling resources helps conserve Earth's limited supplies and reduce pollution.

Chemistry

Paper 1

Topic 1 – Atomic Structure and the Periodic Table

5.1.1.1 Atoms, Elements and Compounds

5.1.1.2 Mixtures

5.1.1.3 The Development of the Model of the Atom

5.1.1.4 Relative Electrical Charges of Subatomic Particles

5.1.1.5 Size and Mass of Atoms

5.1.1.6 Relative Atomic Mass

5.1.1.7 Electronic Structure

5.1.2.1 The Periodic Table

5.1.2.2 Development of the Periodic Table

5.1.2.3 Metals and Non-metals

5.1.2.4 Group 0

5.1.2.5 Group 1

5.1.2.6 Group 7

Topic 1 Review

Topic 2 – Bonding, Structure and the Properties of Matter

5.2.1.1 Chemical Bonds

5.2.1.2 Ionic Bonding

5.2.1.3 Ionic Compounds

5.2.1.4 Covalent Bonding

5.2.1.5 Metallic Bonding

5.2.2.1 The Three States of Matter

5.2.2.2 State Symbols

5.2.2.3 Properties of Ionic Compounds

5.2.2.4 Properties of Small Molecules

5.2.2.5 Polymers

5.2.2.6 Giant Covalent Structures

5.2.2.7 Properties of Metals and Alloys

5.2.2.8 Metals as Conductors

5.2.3.1 Diamond

5.2.3.2 Graphite

5.2.3.3 Graphene and Fullerenes

Topic 2 Review

Topic 3 – Quantitative Chemistry

5.3.1.1 Conservation of Mass and Balanced Chemical Equations

5.3.1.2 Relative Formula Mass

5.3.1.3 Mass Changes When a Reactant or Product is a Gas

5.3.1.4 Chemical Measurements

5.3.2.1 Moles (HT Only)

5.3.2.2 Amounts of Substances in Equations (HT Only)

5.3.2.3 Using Moles to Balance Equations (HT Only)

5.3.2.4 Limiting Reactants (HT Only)

5.3.2.5 Concentration of Solutions

Topic 3 Review

Topic 4 – Chemical Changes

5.4.1.1 Metal Oxides

5.4.1.2 The Reactivity Series

5.4.1.3 Extraction of Metals and Reduction

5.4.1.4 Oxidation and Reduction in Terms of Electrons (HT Only)

5.4.2.1 Reactions of Acids with Metals

5.4.2.2 Neutralisation of Acids and Salt Production

5.4.2.3 Soluble Salts

5.4.2.4 The pH Scale and Neutralisation

5.4.2.5 Strong and Weak Acids (HT Only)

5.4.3.1 The Process of Electrolysis

5.4.3.2 Electrolysis of Molten Ionic Compounds

5.4.3.3 Using Electrolysis to Extract Metals

5.4.3.4 Electrolysis of Aqueous Solutions

5.4.3.5 Representation of Reactions at Electrodes as Half Equations (HT Only)

Topic 4 Review

Topic 5 – Energy Changes

5.5.1.1 Energy Transfer During Exothermic and Endothermic Reactions

5.5.1.2 Reaction Profiles

5.5.1.3 The Energy Change of Reactions (HT Only)

Topic 5 Review

Paper 2

Topic 6 – The Rate and Extent of Chemical Change

5.6.1.1 Calculating rates of reactions

5.6.1.2 Factors which affect the rates of chemical reactions

5.6.1.3 Collision theory and activation energy

5.6.1.4 Catalysts

5.6.2.1 Reversible reactions

5.6.2.2 Energy changes and reversible reactions

5.6.2.3 Equilibrium

5.6.2.4 The effect of changing conditions on equilibrium (HT only)

5.6.2.5 The effect of changing concentration (HT only)

5.6.2.6 The effect of temperature changes on equilibrium (HT only)

5.6.2.7 The effect of pressure changes on equilibrium (HT only)