OCR Combined Science: Chemistry — Recall & Retrieval

C1.1a Describe the main features of the particle model in terms of states of matter and change of state

1.

What are the three states of matter?

Solid, liquid and gas.

2.

How are particles arranged in a solid?

Particles are closely packed in a regular arrangement.

3.

How are particles arranged in a liquid?

Particles are close together but arranged randomly.

4.

How are particles arranged in a gas?

Particles are far apart and arranged randomly.

5.

How do particles move in a solid?

Particles vibrate around fixed positions.

6.

How do particles move in a liquid?

Particles move around each other.

C1.1b Explain in terms of the particle model the distinction between physical changes and chemical changes

1.

What is a physical change?

A physical change does not produce a new substance.

2.

What is a chemical change?

A chemical change produces one or more new substances.

3.

What happens to the particles during a physical change?

The particles remain the same but their arrangement or movement changes.

4.

What happens to the particles during a chemical change?

Atoms are rearranged to form different substances.

5.

Is melting a physical change or a chemical change?

A physical change.

6.

Is the formation of a new substance a physical change or a chemical change?

A chemical change.

C1.1c Explain the limitations of the particle model in relation to changes of state when particles are represented by inelastic spheres (e.g. like bowling balls)

1.

What type of forces between particles are not included in the simple particle model?

Forces of attraction between particles.

2.

Why is the size of particles not accurately represented by the simple particle model?

The particles are shown as much larger than they really are.

3.

Why is the space between particles not accurately represented by the simple particle model?

The spaces between particles are not shown accurately.

4.

What does the simple model assume about the particles when they are represented as inelastic spheres?

The particles are assumed to be hard, inelastic spheres that do not change shape.

5.

Why can the simple particle model not fully explain changes of state?

It does not accurately represent the forces and spacing between particles.

6.

What three features of particles are not taken into account by the simple inelastic-sphere model?

Particle size, particle spacing and forces between particles.

C2.1a Explain what is meant by the purity of a substance, distinguishing between the scientific and everyday use of the term ‘pure’

1.

What does the term pure mean in chemistry?

A pure substance contains only one substance with no other substances mixed with it.

2.

What is a pure substance made up of?

One element or one compound.

3.

What does pure mean in everyday language?

Free from contamination or unwanted substances.

4.

How is the scientific meaning of pure different from the everyday meaning?

In chemistry, pure has a precise meaning: only one substance is present.

5.

What is an impurity?

A substance present in a material that is not wanted.

6.

How can impurities affect the properties of a substance?

They can change properties such as melting point and boiling point.

C2.1b Use melting point data to distinguish pure from impure substances

1.

What happens to the melting point of a pure substance when it melts?

It melts at one fixed temperature.

2.

What happens to the melting point of an impure substance compared with the pure substance?

Its melting point is usually lower and occurs over a range of temperatures.

3.

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

Compare the measured melting point with the known melting point of the pure substance.

4.

What would a sharp melting point suggest about the purity of a substance?

The substance is likely to be pure.

5.

What would a range of melting temperatures suggest about the purity of a substance?

The substance is likely to be impure.

6.

A substance melts over a range of temperatures rather than at one temperature. What does this suggest about its purity?

It suggests that the substance is impure.

C2.1c Calculate relative formula masses of species separately and in a balanced chemical equation

1.

What is relative atomic mass (Ar)?

The relative atomic mass is the average mass of an atom compared with a standard carbon atom.

2.

What is relative molecular mass (Mr)?

Relative molecular mass is the total of the relative atomic masses of all atoms in a molecule.

3.

What is relative formula mass?

The total of the relative atomic masses of all atoms shown in a formula.

4.

How do you calculate the relative formula mass of a compound?

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

5.

Calculate the relative formula mass of H₂O. (Ar: H = 1, O = 16)

Mr(H₂O) = (2 × 1) + 16 = 18.

6.

Calculate the relative formula mass of CaCO₃. (Ar: Ca = 40, C = 12, O = 16)

Mr(CaCO₃) = 40 + 12 + (3 × 16) = 100.

C2.1d Deduce the empirical formula of a compound from the relative numbers of atoms present or from a model or diagram and vice versa

1.

What is an empirical formula?

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

2.

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

The simplest ratio of each type of atom.

3.

How do you determine an empirical formula from the numbers of atoms present?

Divide the numbers of each atom by their highest common factor.

4.

What is the empirical formula of a compound containing 2 carbon atoms and 4 hydrogen atoms?

C₂H₄ → CH₂.

5.

What is the empirical formula of a compound containing 1 carbon atom and 2 oxygen atoms?

CO₂ → CO₂.

6.

How can you determine the relative numbers of atoms in a compound from its empirical formula?

Use the subscripts in the empirical formula to find the simplest ratio of the atoms.

C2.1e Explain that many useful materials are formulations of mixtures

1.

What is a formulation?

A mixture designed to have particular properties.

2.

What is a mixture?

Two or more substances that are not chemically bonded together.

3.

Why are formulations useful?

They can be designed to have useful properties for a particular purpose.

4.

What is an alloy?

A mixture containing a metal and one or more other elements.

5.

Why are alloys described as mixtures?

Their components are not chemically bonded in fixed ratios.

6.

Give one example of an alloy.

Brass.

C2.1f Describe, explain and exemplify the processes of filtration, crystallisation, simple distillation, and fractional distillation

1.

What type of mixture can be separated by filtration?

An insoluble solid and a liquid.

2.

How does filtration separate a solid from a liquid?

The mixture passes through filter paper; the liquid passes through while the insoluble solid remains as residue.

3.

What is crystallisation used to separate?

A dissolved solid from a solution.

4.

How does crystallisation produce a solid from a solution?

The solvent is evaporated and the dissolved substance forms crystals.

5.

What is simple distillation used to separate?

A solvent from a solution, or liquids with very different boiling points.

6.

What is the difference between simple distillation and fractional distillation?

Fractional distillation uses a fractionating column and separates liquids with similar boiling points more effectively.

C2.1g Describe the techniques of paper and thin layer chromatography using aqueous and non-aqueous solvents and locating agents

1.

What is paper chromatography used to separate?

Mixtures of dissolved substances.

2.

What is the stationary phase in paper chromatography?

Paper.

3.

What is used as the stationary phase in thin layer chromatography?

A thin layer of a solid such as silica.

4.

What is the solvent used for in chromatography?

It carries the substances through the stationary phase.

5.

What is a locating agent used for in chromatography?

It makes substances that are otherwise invisible visible.

6.

How can aqueous and non-aqueous solvents be used in chromatography?

Either type of solvent can be selected depending on which substances need to be separated.

C2.1h Recall that chromatography involves a stationary and a mobile phase and that separation depends on the distribution between the phases

1.

What is the stationary phase in chromatography?

The phase that stays fixed, such as paper or silica.

2.

What is the mobile phase in chromatography?

The solvent that moves through the stationary phase.

3.

What happens to a substance that has a greater attraction to the stationary phase?

It moves more slowly and travels a shorter distance.

4.

What happens to a substance that has a greater attraction to the mobile phase?

It moves further with the mobile phase.

5.

What causes different substances to separate during chromatography?

Different attractions between substances and the two phases.

6.

What determines how far a substance travels during chromatography?

How strongly it interacts with the stationary and mobile phases.

C2.1i Interpret chromatograms, including measuring Rf values

1.

What is a chromatogram?

The pattern of separated substances produced by chromatography.

2.

What does a single spot at one position on a chromatogram show?

It usually indicates one substance is present.

3.

What do several spots on a chromatogram show?

It indicates that more than one substance is present.

4.

What does the Rf value measure?

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

5.

What is the formula for calculating Rf?

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

6.

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

Rf = 4 ÷ 8 = 0.50.

C2.1j Suggest suitable purification techniques given information about the substances involved

1.

Which separation technique is suitable for separating an insoluble solid from a liquid?

Filtration.

2.

Which purification technique is suitable for obtaining crystals of a dissolved solid?

Crystallisation.

3.

Which distillation technique is suitable for separating a solvent from a solution?

Simple distillation.

4.

Which distillation technique is suitable for separating two miscible liquids with different boiling points?

Fractional distillation.

5.

What information about substances should be considered when choosing a purification technique?

Properties such as solubility, boiling point and whether a substance is soluble or insoluble.

6.

Why can differences in boiling point be used to separate liquids by distillation?

Liquids with different boiling points vaporise and condense at different temperatures.

C2.1k Suggest chromatographic methods for distinguishing pure from impure substances

1.

How can paper chromatography be used to distinguish a pure substance from an impure substance?

A pure substance usually gives one spot; an impure substance usually gives more than one spot.

2.

How can thin layer chromatography be used to distinguish a pure substance from an impure substance?

The same principle applies using a thin layer of stationary material.

3.

What would a single spot on a chromatogram usually suggest about a substance?

A pure substance.

4.

What would several spots on a chromatogram suggest about a substance?

An impure substance or mixture.

5.

What type of chromatography can be used to separate and identify components in a gas mixture?

Gas chromatography can be used to separate and identify components in a gas mixture.

6.

How can chromatography provide evidence that a substance is impure?

Multiple spots show that more than one substance is present.

C2.2a Describe metals and non-metals and explain the differences between them on the basis of their characteristic physical and chemical properties

1.

What are two typical physical properties of metals?

Good electrical conductivity and high melting points.

2.

What are two typical physical properties of non-metals?

Poor electrical conductivity and generally low melting points.

3.

What type of ion do metals usually form?

Positive ions.

4.

What type of ion do non-metals usually form?

Negative ions.

5.

What is formed when a metal reacts with oxygen?

A metal oxide.

6.

What is the general word equation for a metal reacting with oxygen?

Metal + oxygen → metal oxide.

C2.2b Explain how the atomic structure of metals and non-metals relates to their position in the Periodic Table

1.

Where are metals found in the Periodic Table?

Mainly on the left and centre.

2.

Where are non-metals found in the Periodic Table?

Mainly on the right.

3.

What happens to the number of outer-shell electrons across a period from left to right?

It generally increases by one across a period.

4.

Why do metals tend to lose electrons when they react?

They usually have relatively few outer-shell electrons and can lose them to form positive ions.

5.

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

They tend to gain or share electrons to obtain a stable outer shell.

6.

How does the arrangement of electrons help explain whether an element is a metal or non-metal?

The number of outer-shell electrons influences how atoms react.

C2.2c Explain how the position of an element in the Periodic Table is related to the arrangement of electrons in its atoms and hence to its atomic number, group number and period number

1.

What does the atomic number of an element tell you?

The number of protons in the nucleus.

2.

What does the group number tell you about the electrons in the outer shell of a main-group element?

The number of electrons in the outer shell for main-group elements.

3.

What does the period number tell you about an atom?

The number of occupied electron shells.

4.

How is the atomic number related to the number of electrons in a neutral atom?

They have the same number of electrons.

5.

How can the electron arrangement of an atom be used to identify its group?

Count the electrons in the outer shell.

6.

How can the electron arrangement of an atom be used to identify its period?

Count the occupied electron shells.

C2.2d Describe and compare the nature and arrangement of chemical bonds in: i. ionic compounds ii. simple molecules iii. giant covalent structures iv. polymers v. metals

1.

What type of bonding occurs in ionic compounds?

Ionic bonding.

2.

How are particles arranged in an ionic compound?

A giant lattice of oppositely charged ions.

3.

What type of bonding holds the atoms together in simple molecules?

Covalent bonding.

4.

How are atoms arranged in a giant covalent structure?

A giant network of atoms joined by covalent bonds.

5.

How are atoms joined together in polymers?

By strong covalent bonds in long chains.

6.

What holds the positive metal ions and delocalised electrons together in a metal?

Electrostatic attraction between positive metal ions and delocalised electrons.

C2.2e Explain chemical bonding in terms of electrostatic forces and the transfer or sharing of electrons

1.

What is meant by the transfer of electrons?

Electrons move from one atom to another.

2.

What happens to electrons when a positive ion forms?

It loses electrons.

3.

What happens to electrons when a negative ion forms?

It gains electrons.

4.

What is meant by the sharing of electrons?

Atoms share electrons.

5.

What electrostatic attraction holds oppositely charged ions together?

Strong electrostatic attraction between oppositely charged ions.

6.

What electrostatic attraction holds the particles together in metallic bonding?

Attraction between positive metal ions and delocalised electrons.

C2.2f Construct dot and cross diagrams for simple covalent and binary ionic substances

1.

What do the dots and crosses represent in a dot-and-cross diagram?

Electrons from different atoms.

2.

What does a dot-and-cross diagram show about the electrons involved in bonding?

Which electrons are shared or transferred during bonding.

3.

How many electrons are shared in a single covalent bond?

Two electrons.

4.

How would you show the electron transfer when sodium forms a sodium ion and chlorine forms a chloride ion?

Sodium loses one electron to form Na⁺ and chlorine gains one electron to form Cl⁻.

5.

How would you use a dot-and-cross diagram to show bonding in a chlorine molecule, Cl₂?

Show one shared pair of electrons between the two chlorine atoms, with each chlorine having a full outer shell.

6.

How would you use a dot-and-cross diagram to show bonding in a magnesium oxide compound, MgO?

Show two electrons transferred from Mg to O, forming Mg²⁺ and O²⁻ with full outer shells.

C2.2g Describe the limitations of particular representations and models

1.

What is one limitation of a dot-and-cross diagram?

It does not show the actual three-dimensional shape or electron density.

2.

What is one limitation of a ball-and-stick model?

It does not accurately represent the sizes or distances of atoms.

3.

Why does a ball-and-stick model not show the actual sizes of atoms?

Atoms are shown as large spheres for simplicity.

4.

What is one limitation of a two-dimensional representation of a molecule?

It cannot accurately show the three-dimensional structure.

5.

What is one limitation of a three-dimensional representation of a molecule?

It may hide some details or be difficult to interpret.

6.

Why can no single model give a complete representation of a chemical structure?

Each model simplifies some features and cannot show everything at once.

C2.2h Explain how the reactions of elements are related to the arrangement of electrons in their atoms and hence to their atomic number

1.

Which electrons are mainly involved when elements react?

Outer-shell electrons.

2.

Why are outer-shell electrons important in chemical reactions?

They are involved in forming chemical bonds.

3.

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

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

4.

How does an element's atomic number determine the number of electrons in a neutral atom?

A neutral atom has the same number of electrons as protons.

5.

How does electron arrangement affect the way an element reacts?

It determines whether atoms are likely to lose, gain or share electrons.

6.

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

They have the same number of outer-shell electrons.

C2.2i Explain in terms of atomic number how Mendeleev’s arrangement was refined into the modern Periodic Table

1.

What did Mendeleev use to arrange the elements in his Periodic Table?

Increasing atomic mass, while considering similarities in properties.

2.

What important information did Mendeleev use to identify patterns in the elements?

Atomic masses and chemical properties.

3.

What is the modern Periodic Table arranged in order of?

Increasing atomic number.

4.

What does atomic number represent?

The number of protons in the nucleus.

5.

Why did the modern Periodic Table need to be refined from Mendeleev's arrangement?

Some elements did not fit perfectly when arranged by atomic mass.

6.

How does arranging elements by atomic number produce repeating patterns in their properties?

Elements with similar electron arrangements and properties appear at regular intervals.

C2.3a Recall that carbon can form four covalent bonds

1.

How many covalent bonds can one carbon atom form?

Four.

2.

How many electrons are in the outer shell of a carbon atom?

Four.

3.

Why can carbon form four covalent bonds?

It has four outer-shell electrons and needs four more to achieve a full outer shell.

4.

How many covalent bonds does carbon form in methane, CH₄?

Four.

5.

How many covalent bonds does each carbon atom form in ethane, C₂H₆?

Four.

6.

How does carbon's ability to form four covalent bonds allow it to form many different compounds?

It can bond with itself and many other elements in different arrangements.

C2.3b Explain that the vast array of natural and synthetic organic compounds occur due to the ability of carbon to form families of similar compounds, chains and rings

1.

What is an organic compound?

A compound containing carbon, usually with hydrogen.

2.

Why can carbon form chains of atoms?

Carbon atoms can form covalent bonds with other carbon atoms.

3.

Why can carbon form rings of atoms?

Carbon atoms can join together in closed chains.

4.

What is a family of similar organic compounds?

A group of compounds with similar structures and chemical properties.

5.

How does carbon's ability to form chains and rings produce many different organic compounds?

Carbon can form many different structures and combinations.

6.

What is the difference between a natural organic compound and a synthetic organic compound?

Natural compounds occur naturally; synthetic compounds are made by humans.

C2.3c Explain the properties of diamond, graphite, fullerenes and graphene in terms of their structures and bonding

1.

What type of bonding is found in diamond?

Strong covalent bonding.

2.

Why is diamond very hard?

Each carbon atom is strongly bonded to four other carbon atoms.

3.

How are carbon atoms arranged in graphite?

Layers of carbon atoms arranged in hexagonal structures.

4.

Why can graphite conduct electricity?

It has delocalised electrons that can move through the layers.

5.

What are fullerenes?

Carbon structures made from carbon atoms arranged in hollow shapes.

6.

How is graphene structured and why is it strong?

It is a single layer of carbon atoms in a hexagonal arrangement; the strong covalent bonds make it strong.

C2.3d Use ideas about energy transfers and the relative strength of chemical bonds and intermolecular forces to explain the different temperatures at which changes of state occur

1.

What happens to particles when a substance changes state?

The arrangement and movement of particles change.

2.

What happens to energy during melting?

Energy is absorbed.

3.

What happens to energy during boiling?

Energy is absorbed.

4.

What is an intermolecular force?

An attractive force between molecules.

5.

How does the strength of intermolecular forces affect the temperature at which a substance changes state?

Stronger intermolecular forces require more energy to overcome.

6.

Why do substances with stronger intermolecular forces generally have higher boiling points?

More energy is needed to separate the molecules.

C2.3e Use data to predict states of substances under given conditions

1.

What information can be used to predict the state of a substance?

Melting point, boiling point and temperature.

2.

What happens to a substance when its temperature reaches its melting point?

It changes from solid to liquid.

3.

What happens to a substance when its temperature reaches its boiling point?

It changes from liquid to gas.

4.

A substance has a melting point of 20°C. What state is it likely to be at 10°C?

Solid.

5.

A substance has a boiling point of 80°C. What state is it likely to be at 100°C?

Gas.

6.

A substance has a melting point of 0°C and a boiling point of 100°C. What state is it at 50°C?

Liquid.

C2.3f Explain how the bulk properties of materials (ionic compounds; simple molecules; giant covalent structures; polymers and metals) are related to the different types of bonds they contain, their bond strengths in relation to intermolecular forces and the ways in which their bonds are arranged

1.

What is meant by the bulk properties of a material?

Properties of a material that can be observed or measured on a large scale.

2.

Why do ionic compounds generally have high melting points?

Strong electrostatic attractions between ions require lots of energy to overcome.

3.

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

Only weak intermolecular forces need to be overcome.

4.

Why do giant covalent structures generally have very high melting points?

Many strong covalent bonds need to be broken.

5.

Why can metals conduct electricity?

They contain delocalised electrons that can carry charge.

6.

Why do the bulk properties of a material depend on its bonding and structure rather than the properties of individual atoms?

Bonding and structure determine how particles interact on a large scale.

C3.1a Use chemical symbols to write the formulae of elements and simple covalent and ionic compounds

1.

What is the chemical symbol for sodium?

Na.

2.

What is the chemical formula for oxygen gas?

O₂.

3.

What is the chemical formula for water?

H₂O.

4.

What is the chemical formula for carbon dioxide?

CO₂.

5.

What is the chemical formula for sodium chloride?

NaCl.

6.

What is the chemical formula for magnesium oxide?

MgO.

C3.1b Use the names and symbols of common elements and compounds and the principle of conservation of mass to write formulae and balanced chemical equations and half equations

1.

What is the chemical symbol for calcium?

Ca.

2.

What is the chemical formula for calcium chloride?

CaCl₂.

3.

Write the balanced equation for magnesium reacting with oxygen to form magnesium oxide.

2Mg + O₂ → 2MgO.

4.

Write the balanced equation for hydrogen reacting with oxygen to form water.

2H₂ + O₂ → 2H₂O.

5.

What must be conserved when balancing a chemical equation?

The number of atoms of each element.

6.

What does a half equation show about the electrons transferred during a reaction?

The electrons transferred during a redox reaction.

C3.1c Use the names and symbols of common elements from a supplied Periodic Table to write formulae and balanced chemical equations where appropriate

1.

What is the chemical symbol for potassium?

K.

2.

What is the chemical symbol for chlorine?

Cl.

3.

What is the chemical symbol for iron?

Fe.

4.

What is the chemical symbol for copper?

Cu.

5.

Write the formula for aluminium oxide.

Al₂O₃.

6.

Write the balanced equation for calcium reacting with water to form calcium hydroxide and hydrogen.

Ca + 2H₂O → Ca(OH)₂ + H₂.

C3.1d Use the formula of common ions to deduce the formula of a compound

1.

What is the charge on a sodium ion, Na⁺?

+1.

2.

What is the charge on a chloride ion, Cl⁻?

−1.

3.

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

NaCl.

4.

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

MgCl₂.

5.

What is the formula of the compound formed from Ca²⁺ and O²⁻?

CaO.

6.

What is the formula of the compound formed from Al³⁺ and O²⁻?

Al₂O₃.

C3.1e Construct balanced ionic equations

1.

What is an ionic equation?

An equation showing only the ions and substances directly involved in a reaction.

2.

What should be removed from a complete ionic equation to produce a net ionic equation?

Spectator ions.

3.

Write the ionic equation for the reaction between hydrogen ions and hydroxide ions.

H⁺ + OH⁻ → H₂O.

4.

Write the ionic equation for the formation of a precipitate of silver chloride.

Ag⁺ + Cl⁻ → AgCl.

5.

How must the total charge compare on both sides of a balanced ionic equation?

The total charge must be equal on both sides.

6.

How must the number of each type of atom compare on both sides of a balanced ionic equation?

The number of each type of atom must be equal on both sides.

C3.1f Describe the physical states of products and reactants using state symbols (s, l, g and aq)

1.

What does the state symbol (s) mean?

Solid.

2.

What does the state symbol (l) mean?

Liquid.

3.

What does the state symbol (g) mean?

Gas.

4.

What does the state symbol (aq) mean?

Aqueous; dissolved in water.

5.

What state symbol would be used for solid sodium chloride?

NaCl(s).

6.

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

NaCl(aq).

C3.1g Describe tests to identify selected gases

1.

What test is used to identify oxygen gas?

A glowing splint.

2.

What result is seen when a glowing splint is placed in oxygen?

It relights.

3.

What test is used to identify hydrogen gas?

A lit splint.

4.

What result is seen when a lit splint is placed in hydrogen?

It produces a squeaky pop.

5.

What test is used to identify carbon dioxide?

Bubble the gas through limewater.

6.

What test and result can be used to identify chlorine gas?

Damp blue litmus paper is bleached white.

C3.1h Recall and use the definitions of the Avogadro constant (in standard form) and of the mole

1.

What is the Avogadro constant?

The number of particles in one mole.

2.

What is the value of the Avogadro constant?

6.02 × 10²³ mol⁻¹.

3.

What is the unit of the Avogadro constant?

mol⁻¹.

4.

What is a mole?

The amount of substance containing the Avogadro number of particles.

5.

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

6.02 × 10²³.

6.

How can the Avogadro constant be used to calculate the number of particles in a given amount of substance?

Number of particles = moles × Avogadro constant.

C3.1i Explain how the mass of a given substance is related to the amount of that substance in moles and vice versa

1.

What is the formula linking mass, moles and molar mass?

moles = mass ÷ molar mass.

2.

What is the unit of amount of substance?

mol.

3.

What is the unit of mass commonly used in mole calculations?

g.

4.

What is the unit of molar mass?

g mol⁻¹.

5.

Calculate the amount of substance in 18 g of water. (Mr of H₂O = 18)

18 ÷ 18 = 1 mol.

6.

Calculate the mass of 0.5 mol of carbon dioxide. (Mr of CO₂ = 44)

0.5 × 44 = 22 g.

C3.1j Explain how the mass of a solute and the volume of the solution is related to the concentration of the solution

1.

What is a solute?

A substance dissolved in a solvent.

2.

What is a solution?

A mixture formed when a solute dissolves in a solvent.

3.

What is concentration?

The amount of solute in a given volume of solution.

4.

What is the formula linking concentration, mass of solute and volume of solution?

Concentration = mass ÷ volume.

5.

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

20 ÷ 2 = 10 g/dm³.

6.

Calculate the mass of solute needed to make 500 cm³ of a solution with a concentration of 40 g/dm³.

40 × 0.5 = 20 g.

C3.1k Recall and use the law of conservation of mass

1.

What does the law of conservation of mass state?

Mass cannot be created or destroyed.

2.

What happens to the total mass during a chemical reaction?

It stays the same.

3.

Why must the total mass of the reactants equal the total mass of the products?

Atoms are rearranged but not created or destroyed.

4.

What happens to atoms during a chemical reaction that explains conservation of mass?

The same atoms are present before and after the reaction.

5.

A reaction starts with 15 g of reactants. What is the total mass of the products?

15 g.

6.

A reaction produces 24 g of products from two reactants. If one reactant has a mass of 9 g, what is the mass of the other reactant?

24 − 9 = 15 g.

C3.1l Explain any observed changes in mass in non-enclosed systems during a chemical reaction and explain them using the particle model

1.

Why can the measured mass change during a reaction in an open container?

Particles can enter or leave the system.

2.

What happens to the measured mass if a gas escapes from an open reaction?

The measured mass decreases.

3.

What happens to the measured mass if a gas from the air enters an open reaction?

The measured mass increases.

4.

Why does the total mass of all particles remain constant even when the measured mass changes?

Particles have moved into or out of the system rather than being created or destroyed.

5.

Why can burning a substance in an open container cause an increase in measured mass?

Oxygen from the air combines with the substance.

6.

How can the particle model explain an apparent change in mass in an open system?

Gas particles can escape or enter, changing the measured mass.

C3.1m Deduce the stoichiometry of an equation from the masses of reactants and products and explain the effect of a limiting quantity of a reactant

1.

What does the stoichiometry of a chemical equation describe?

The ratio in which substances react according to a balanced equation.

2.

How can a balanced equation show the ratio in which substances react?

The coefficients give the mole ratio of the substances.

3.

What is meant by the limiting reactant?

The reactant that is completely used up first.

4.

What happens to the limiting reactant during a reaction?

It is completely used up.

5.

What happens to an excess reactant when the limiting reactant is completely used up?

It remains after the reaction stops.

6.

What mass of aluminium oxide is formed when 135g of aluminium is burned in air?

The balanced equation is: 4Al + 3O₂ → 2Al₂O₃. Using Ar: Al = 27, O = 16: Mr of Al₂O₃ = (2 × 27) + (3 × 16) = 102. Moles of Al = 135 ÷ 27 = 5 mol. From the equation, 4 mol Al → 2 mol Al₂O₃. Therefore, 5 mol Al → 2.5 mol Al₂O₃. Mass = 2.5 × 102 = 255 g.

C3.1n Use a balanced equation to calculate masses of reactants or product

1.

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

The mole ratio from the balanced equation and the molar masses.

2.

What calculation can be used to convert mass into moles?

moles = mass ÷ Mr.

3.

What calculation can be used to convert moles into mass?

mass = moles × Mr.

4.

Calculate the mass of magnesium oxide produced when 0.5 mol of MgO is formed. (Mr of MgO = 40)

0.5 × 40 = 20 g.

5.

Calculate the amount in moles of carbon dioxide in 88 g of CO₂. (Mr of CO₂ = 44)

88 ÷ 44 = 2 mol.

6.

Why must a balanced chemical equation be used when calculating the mass of a reactant or product?

It gives the correct mole ratio between reactants and products.

C3.2a Distinguish between endothermic and exothermic reactions on the basis of the temperature change of the surroundings

1.

What is an exothermic reaction?

A reaction that transfers energy to the surroundings.

2.

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

It increases.

3.

What is an endothermic reaction?

A reaction that takes in energy from the surroundings.

4.

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

It decreases.

5.

A reaction causes the temperature of the surroundings to increase. Is the reaction exothermic or endothermic?

Exothermic.

6.

A reaction causes the temperature of the surroundings to decrease. Is the reaction exothermic or endothermic?

Endothermic.

C3.2b Draw and label a reaction profile for an exothermic and an endothermic reaction

1.

What does a reaction profile show?

The energy changes during a chemical reaction.

2.

What does the vertical axis represent on a reaction profile?

Energy.

3.

What does the horizontal axis represent on a reaction profile?

Progress of reaction.

4.

What is activation energy shown as on a reaction profile?

The energy needed to start the reaction.

5.

How does the energy level of the products compare with the reactants in an exothermic reaction?

Products have less energy than the reactants.

6.

How does the energy level of the products compare with the reactants in an endothermic reaction?

Products have more energy than the reactants.

C3.2c Explain activation energy as the energy needed for a reaction to occur

1.

What is activation energy?

The minimum energy needed for a reaction to occur.

2.

Why is activation energy needed for a chemical reaction?

Particles need enough energy for bonds to break and the reaction to begin.

3.

Where is activation energy shown on a reaction profile?

From the reactants' energy level to the highest point of the profile.

4.

What happens if reacting particles do not have enough energy to overcome the activation energy?

The reaction does not occur successfully.

5.

Does an exothermic reaction still require activation energy?

Yes.

6.

Does an endothermic reaction still require activation energy?

Yes.

C3.2d Calculate energy changes in a chemical reaction by considering bond making and bond breaking energies

1.

What happens to energy when chemical bonds are broken?

Energy is absorbed.

2.

What happens to energy when chemical bonds are formed?

Energy is released.

3.

What calculation is used to determine the overall energy change from bond energies?

Energy change = energy needed to break bonds − energy released when bonds form.

4.

A reaction requires 500 kJ/mol to break bonds and releases 700 kJ/mol when new bonds form. What is the overall energy change?

500 − 700 = −200 kJ/mol, so the reaction is exothermic.

5.

If more energy is released when bonds form than is absorbed when bonds break, is the reaction exothermic or endothermic?

Exothermic.

6.

If more energy is absorbed when bonds break than is released when bonds form, is the reaction exothermic or endothermic?

Endothermic.

C3.3a Explain reduction and oxidation in terms of loss or gain of oxygen, identifying which species are oxidised and which are reduced

1.

What is oxidation in terms of oxygen?

Gain of oxygen.

2.

What is reduction in terms of oxygen?

Loss of oxygen.

3.

What happens to a substance when it is oxidised?

It gains oxygen.

4.

What happens to a substance when it is reduced?

It loses oxygen.

5.

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

Copper oxide.

6.

In the same reaction, which substance is oxidised?

Carbon.

C3.3b Explain reduction and oxidation in terms of gain or loss of electrons, identifying which species are oxidised and which are reduced

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?

It loses electrons.

4.

What happens to electrons when a substance is reduced?

It gains electrons.

5.

Which substance is oxidised when magnesium reacts to form Mg²⁺?

Magnesium.

6.

Which substance is reduced when Cu²⁺ gains electrons to form copper?

Cu²⁺.

C3.3c Recall that acids form hydrogen ions when they dissolve in water and solutions of alkalis contain hydroxide ions

1.

Which ions do acids form when they dissolve in water?

Hydrogen ions, H⁺.

2.

Which ions are present in aqueous alkali solutions?

Hydroxide ions, OH⁻.

3.

What is the formula for a hydrogen ion?

H⁺.

4.

What is the formula for a hydroxide ion?

OH⁻.

5.

What ion makes an aqueous solution acidic?

H⁺.

6.

What ion makes an aqueous solution alkaline?

OH⁻.

C3.3d Describe neutralisation as acid reacting with alkali or a base to form a salt plus water

1.

What is neutralisation?

An acid reacting with an alkali or base to form a salt and water.

2.

What are the products of an acid reacting with an alkali?

Salt and water.

3.

What are the products of an acid reacting with a base?

Salt and water.

4.

What is the general word equation for an acid reacting with an alkali?

Acid + alkali → salt + water.

5.

What salt is formed when hydrochloric acid reacts with sodium hydroxide?

Sodium chloride.

6.

What salt is formed when sulfuric acid reacts with potassium hydroxide?

Potassium sulfate.

C3.3e Recognise that aqueous neutralisation reactions can be generalised to hydrogen ions reacting with hydroxide ions to form water

1.

Which ions react during aqueous neutralisation?

Hydrogen ions and hydroxide ions.

2.

What is the formula of the hydrogen ion?

H⁺.

3.

What is the formula of the hydroxide ion?

OH⁻.

4.

What is the formula of the water molecule formed during neutralisation?

H₂O.

5.

Write the balanced ionic equation for neutralisation.

H⁺ + OH⁻ → H₂O.

6.

What happens to hydrogen ions and hydroxide ions during neutralisation?

They react together to form water.

C3.3f Recall that carbonates and some metals react with acids and write balanced equations predicting products from given reactants

1.

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

Carbon dioxide.

2.

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

Salt, water and carbon dioxide.

3.

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

Hydrogen.

4.

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

Salt and hydrogen.

5.

Write the word equation for calcium carbonate reacting with hydrochloric acid.

Calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide.

6.

Write the balanced equation for magnesium reacting with hydrochloric acid.

Mg + 2HCl → MgCl₂ + H₂.

C3.3g Use and explain the terms dilute and concentrated (amount of substance) and weak and strong (degree of ionisation) in relation to acids

1.

What does dilute mean when describing an acid?

A small amount of acid per unit volume of solution.

2.

What does concentrated mean when describing an acid?

A large amount of acid per unit volume of solution.

3.

What does strong mean when describing an acid?

An acid that ionises almost completely in water.

4.

What does weak mean when describing an acid?

An acid that only partially ionises in water.

5.

What is the difference between acid concentration and acid strength?

Concentration describes the amount of acid; strength describes how much the acid ionises.

6.

Which term describes the amount of acid per volume of solution: concentration or strength?

Concentration.

C3.3h Recall that relative acidity and alkalinity are measured by pH

1.

What does the pH scale measure?

How acidic or alkaline a solution is.

2.

What pH value represents a neutral solution?

7.

3.

What pH values represent acidic solutions?

Less than 7.

4.

What pH values represent alkaline solutions?

Greater than 7.

5.

Which is more acidic, pH 2 or pH 5?

pH 2.

6.

Which is more alkaline, pH 12 or pH 9?

pH 12.

C3.3i Describe neutrality and relative acidity and alkalinity in terms of the effect of the concentration of hydrogen ions on the numerical value of pH

1.

What happens to pH when the concentration of hydrogen ions increases?

pH decreases.

2.

What happens to pH when the concentration of hydrogen ions decreases?

pH increases.

3.

What does pH 7 indicate about a solution?

The solution is neutral.

4.

Which has the greater hydrogen ion concentration: pH 2 or pH 4?

pH 2.

5.

Which has the lower hydrogen ion concentration: pH 9 or pH 12?

pH 12.

6.

What does a lower pH value indicate about hydrogen ion concentration?

It indicates a higher hydrogen ion concentration.

C3.3j Use the idea that as hydrogen ion concentration increases by a factor of ten, the pH value of a solution decreases by one

1.

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

pH decreases by 1.

2.

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

pH increases by 1.

3.

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

It increased by a factor of 10.

4.

A solution changes from pH 3 to pH 5. By what factor has its hydrogen ion concentration decreased?

It decreased by a factor of 100.

5.

Which has 100 times the hydrogen ion concentration: pH 2 or pH 4?

pH 2.

6.

A solution has pH 6. What would its pH become if its hydrogen ion concentration increased by a factor of 100?

pH 4.

C3.3k Describe techniques and apparatus used to measure pH

1.

What indicator can be used to estimate the pH of a solution?

Universal indicator.

2.

What instrument can measure pH more precisely than universal indicator?

A pH meter.

3.

What colour range does universal indicator show across the pH scale?

Red at low pH, green around pH 7 and purple at high pH.

4.

How is universal indicator used to estimate the pH of a solution?

Add universal indicator and compare the colour with a pH colour chart.

5.

How is a pH meter used to measure the pH of a solution?

Place the electrode in the solution and read the pH.

6.

Which gives a more precise pH measurement: universal indicator or a pH meter?

A pH meter.

C3.4a Recall that metals (or hydrogen) are formed at the cathode and non-metals are formed at the anode in electrolysis using inert electrodes

1.

What is electrolysis?

The decomposition of an ionic compound using electricity.

2.

What type of ion moves towards the cathode?

Positive ions.

3.

What type of ion moves towards the anode?

Negative ions.

4.

What is formed at the cathode during electrolysis of a molten ionic compound?

A metal, or hydrogen if hydrogen is preferentially discharged.

5.

What is formed at the anode during electrolysis of a molten ionic compound?

A non-metal.

6.

What are cations and anions?

Cations are positive ions; anions are negative ions.

C3.4b Predict the products of electrolysis of binary ionic compounds in the molten state

1.

What happens to an ionic compound when it is molten?

Its ions become free to move.

2.

Which ion is discharged at the cathode during electrolysis of a molten ionic compound?

The positive ion.

3.

Which ion is discharged at the anode during electrolysis of a molten ionic compound?

The negative ion.

4.

What are the products of electrolysis of molten sodium chloride?

Sodium and chlorine.

5.

What is the half equation for the formation of sodium at the cathode?

Na⁺ + e⁻ → Na.

6.

What is the half equation for the formation of chlorine at the anode?

2Cl⁻ → Cl₂ + 2e⁻.

C3.4c Describe competing reactions in the electrolysis of aqueous solutions of ionic compounds in terms of the different species present

1.

What additional particles are present when an ionic compound is dissolved in water?

Hydrogen ions and hydroxide ions from water.

2.

Which two possible products can form at the cathode when aqueous sodium chloride is electrolysed using inert electrodes?

Hydrogen or the metal.

3.

Which two possible products can form at the anode when aqueous sodium chloride is electrolysed using inert electrodes?

A halogen or oxygen.

4.

What are the products of electrolysis of aqueous sodium chloride using inert electrodes?

Hydrogen at the cathode and chlorine at the anode, with sodium hydroxide remaining in solution.

5.

What are the products of electrolysis of aqueous copper sulfate using inert electrodes?

Copper at the cathode and oxygen at the anode.

6.

Why can water compete with the ions from an ionic compound during electrolysis of an aqueous solution?

Water contains ions that can also be discharged at the electrodes.

C3.4d Describe electrolysis in terms of the ions present and reactions at the electrodes

1.

What happens to positive ions at the cathode?

They gain electrons.

2.

What happens to negative ions at the anode?

They lose electrons.

3.

What type of reaction occurs at the cathode?

Reduction.

4.

What type of reaction occurs at the anode?

Oxidation.

5.

Write the half equation for the formation of hydrogen from hydrogen ions.

2H⁺ + 2e⁻ → H₂.

6.

Write the half equation for the formation of chlorine from chloride ions.

2Cl⁻ → Cl₂ + 2e⁻.

C3.4e Describe the technique of electrolysis using inert and non-inert electrodes

1.

What is an inert electrode?

An electrode that does not react during electrolysis.

2.

Give one example of an inert electrode.

Platinum.

3.

Why do inert electrodes generally not react during electrolysis?

They are chemically unreactive under the conditions.

4.

What is a non-inert electrode?

An electrode that can react during electrolysis.

5.

How can a non-inert electrode take part in electrolysis?

It can be oxidised or reduced and become part of the reaction.

6.

What should be connected to the electrodes to provide the electrical energy needed for electrolysis?

A direct current power supply.

C4.1a Recall the simple properties of Groups 1, 7 and 0

1.

What are the elements in Group 1 called?

Alkali metals.

2.

What are the elements in Group 7 called?

Halogens.

3.

What are the elements in Group 0 called?

Noble gases.

4.

What is one physical property of Group 1 metals?

They are relatively soft and have low densities.

5.

What is one physical property of Group 7 elements?

They are non-metals.

6.

What is one chemical property of Group 0 elements?

They are very unreactive.

C4.1b Explain how observed simple properties of Groups 1, 7 and 0 depend on the outer shell of electrons of the atoms and predict properties from given trends down the groups

1.

How many electrons are in the outer shell of Group 1 atoms?

One.

2.

How many electrons are in the outer shell of Group 7 atoms?

Seven.

3.

Why do Group 1 elements tend to lose one electron when they react?

Losing one electron gives them a full outer shell.

4.

Why do Group 7 elements tend to gain one electron when they react?

Gaining one electron gives them a full outer shell.

5.

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

It increases down the group.

6.

How does the reactivity of Group 7 elements change down the group?

It decreases down the group.

C4.1c Predict possible reactions and probable reactivity of elements from their positions in the Periodic Table

1.

What information about an element's position can be used to predict its reactivity?

Its group and position within the group.

2.

Which Group 1 element is more reactive: lithium or potassium?

Potassium.

3.

Which Group 7 element is more reactive: fluorine or iodine?

Fluorine.

4.

What type of reaction occurs when a more reactive halogen reacts with a halide of a less reactive halogen?

Displacement.

5.

Will chlorine displace bromine from potassium bromide?

Yes.

6.

Will iodine displace chlorine from potassium chloride?

No.

C4.1d Explain how the reactivity of metals with water or dilute acids is related to the tendency of the metal to form its positive ion

1.

What happens to a metal atom when it forms a positive ion?

It loses electrons.

2.

Why do metals that readily lose electrons tend to be more reactive?

Losing electrons allows the metal to form a positive ion.

3.

What gas is produced when a reactive metal reacts with water?

Hydrogen.

4.

What gas is produced when a reactive metal reacts with a dilute acid?

Hydrogen.

5.

What two products are formed when a metal reacts with dilute hydrochloric acid?

A metal salt and hydrogen.

6.

Why does a more reactive metal react more readily with water or dilute acid?

It loses electrons more readily and forms positive ions more easily.

C4.1e Deduce an order of reactivity of metals based on experimental results

1.

How can reactions with water be used to compare the reactivity of metals?

More reactive metals react more readily or vigorously with water.

2.

How can reactions with dilute acids be used to compare the reactivity of metals?

More reactive metals generally react more quickly or vigorously.

3.

What does a faster reaction with dilute acid suggest about a metal's reactivity?

The metal is more reactive.

4.

If metal A displaces metal B from its compound, which metal is more reactive?

Metal A.

5.

If magnesium reacts with hydrochloric acid but copper does not, which metal is more reactive?

Magnesium.

6.

How can experimental results be used to place metals in order of reactivity?

Compare the reactions and arrange metals from most to least reactive.

C5.1a Suggest practical methods for determining the rate of a given reaction

1.

What is meant by the rate of a chemical reaction?

How quickly reactants are changed into products.

2.

How can the volume of gas produced be used to measure the rate of a reaction?

Measure the volume of gas produced over time.

3.

How can the loss of mass be used to measure the rate of a reaction?

Measure how quickly the mass decreases.

4.

How can the time taken for a visible change to occur be used to compare reaction rates?

A shorter time means a faster reaction.

5.

In a disappearing cross experiment, what is measured to compare the reaction rates?

The time taken for the cross to disappear.

6.

What calculation can be used to compare rates when the same reaction endpoint is measured by time?

Rate ∝ 1/time.

C5.1b Interpret rate of reaction graphs

1.

What does the gradient of a volume-of-gas-against-time graph represent?

The rate of reaction.

2.

What does a steeper gradient on a rate graph indicate?

A faster reaction.

3.

What does a horizontal section on a volume-of-gas-against-time graph show?

The reaction has stopped producing gas.

4.

What does the final volume of gas on a graph show?

The total amount of gas produced.

5.

If two reactions reach the same final volume but one reaches it sooner, which reaction has the faster rate?

The reaction that reaches the final volume sooner.

6.

If a reaction takes 20 seconds to reach a given endpoint, what is its relative rate using 1/t?

1/20 = 0.05 s⁻¹.

C5.1c Describe the effect of changes in temperature, concentration, pressure, and surface area on rate of reaction

1.

What happens to the rate of a reaction when temperature increases?

It increases.

2.

What happens to the rate of a reaction when the concentration of a reactant increases?

It increases.

3.

What happens to the rate of a reaction involving gases when pressure increases?

It increases.

4.

What happens to the rate of a reaction when the surface area of a solid reactant increases?

It increases.

5.

Which reacts faster with hydrochloric acid: large marble chips or powdered marble?

Powdered marble.

6.

Which reacts faster: a gas reaction at high pressure or the same reaction at low pressure?

The reaction at high pressure.

C5.1d Explain the effects on rates of reaction of changes in temperature, concentration and pressure in terms of frequency and energy of collision between particles

1.

What must happen when particles collide for a chemical reaction to occur?

They must collide with enough energy and in the correct orientation.

2.

How does increasing temperature affect the frequency of successful collisions?

It increases the number of successful collisions.

3.

How does increasing temperature affect the energy of particles?

They have more kinetic energy.

4.

How does increasing concentration affect the frequency of collisions?

Collisions happen more frequently.

5.

How does increasing pressure affect the frequency of collisions between gas particles?

Particles collide more frequently.

6.

Why does increasing temperature usually increase the rate of a chemical reaction?

Particles have more energy and collide more frequently, increasing successful collisions.

C5.1e Explain the effects on rates of reaction of changes in the size of the pieces of a reacting solid in terms of surface area to volume

1.

What happens to the surface area of a solid when it is broken into smaller pieces?

It increases.

2.

What happens to the surface area-to-volume ratio when a solid is broken into smaller pieces?

It increases.

3.

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

More particles are exposed and available to collide.

4.

Which has the greater surface area-to-volume ratio: a large lump or the same mass broken into small pieces?

The small pieces.

5.

Why do powdered solids generally react faster than large pieces of the same solid?

More surface is exposed for collisions.

6.

Why does increasing surface area increase the number of collisions between reactant particles?

More particles are exposed at the surface.

C5.1f Describe the characteristics of catalysts and their effect on rates of reaction

1.

What is a catalyst?

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

2.

What does a catalyst do to the rate of a chemical reaction?

It increases the rate.

3.

Is a catalyst used up during a reaction?

No.

4.

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

It lowers the activation energy.

5.

Does a catalyst change the overall energy change of a reaction?

No.

6.

What happens to the reaction rate when a suitable catalyst is added?

It increases.

C5.1g Identify catalysts in reactions

1.

What catalyst is commonly used to catalyse the decomposition of hydrogen peroxide?

Manganese dioxide.

2.

What is the role of manganese dioxide in the decomposition of hydrogen peroxide?

It catalyses the decomposition of hydrogen peroxide.

3.

What happens to the rate of hydrogen peroxide decomposition when a suitable catalyst is added?

It increases.

4.

How can you identify whether a substance is acting as a catalyst in an experiment?

Compare the reaction rate with and without the substance while checking that the substance is not used up.

5.

What happens to a catalyst at the end of a reaction?

It remains chemically unchanged.

6.

Why can the same catalyst be used in small amounts?

It is not consumed by the reaction.

C5.1h Explain catalytic action in terms of activation energy reaction profiles

1.

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

It decreases.

2.

How is the activation energy of a catalysed reaction shown on a reaction profile?

With a lower peak.

3.

Why does lowering activation energy increase the rate of a reaction?

More particles have enough energy to react.

4.

What happens to the number of successful collisions when activation energy is lowered?

It increases.

5.

How would the reaction profile of a catalysed reaction differ from an uncatalysed reaction?

The catalysed profile has a lower activation-energy peak.

6.

Does a catalyst change the energy of the reactants or products?

No; it changes the activation energy, not the overall energy change.

C5.1i Recall that enzymes act as catalysts in biological systems

1.

What are enzymes?

Biological catalysts, usually proteins.

2.

What role do enzymes play in biological systems?

They speed up biological reactions.

3.

Are enzymes catalysts?

Yes.

4.

What do enzymes do to the activation energy of biological reactions?

They lower the activation energy.

5.

Why are enzymes important for reactions in living organisms?

They allow important reactions to occur quickly at suitable temperatures.

6.

Give one example of a biological process that requires enzymes.

Digestion.

C5.2a Recall that some reactions may be reversed by altering the reaction conditions

1.

What is a reversible reaction?

A reaction that can proceed in both directions.

2.

What symbol is used to represent a reversible reaction?

⇌.

3.

What happens in the forward reaction?

Reactants form products.

4.

What happens in the reverse reaction?

Products form reactants.

5.

How can changing reaction conditions affect a reversible reaction?

They can change the position of equilibrium.

6.

What happens when a reversible reaction can produce the original reactants again?

The products can react to reform the original reactants.

C5.2b Recall that dynamic equilibrium occurs in a closed system when the rates of forward and reverse reactions are equal

1.

What is dynamic equilibrium?

A state in which the forward and reverse reactions occur at equal rates.

2.

In what type of system can dynamic equilibrium be established?

A closed system.

3.

What happens to the rate of the forward reaction at equilibrium?

It equals the rate of the reverse reaction.

4.

What happens to the rate of the reverse reaction at equilibrium?

It equals the rate of the forward reaction.

5.

Are the concentrations of reactants and products constant at dynamic equilibrium?

Yes.

6.

Do the forward and reverse reactions stop at dynamic equilibrium?

No; both reactions continue.

C5.2c Predict the effect of changing reaction conditions on equilibrium position and suggest appropriate conditions to produce as much of a particular product as possible

1.

What happens to the equilibrium position when the concentration of a reactant is increased?

It shifts towards products.

2.

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

It shifts towards ammonia.

3.

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

It shifts towards the reactants.

4.

The forward Haber process is exothermic. What happens to the equilibrium position when the temperature is increased?

It shifts towards the reactants.

5.

For the reaction 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), what happens to the equilibrium position when the pressure is increased?

It shifts towards SO₃.

6.

For an exothermic reversible reaction, what temperature would favour the greatest amount of product at equilibrium?

A lower temperature.

C6.1a Explain, using the position of carbon in the reactivity series, the principles of industrial processes used to extract metals, including extraction of a non-ferrous metal

1.

Where is carbon positioned in the reactivity series compared with metals such as iron and copper?

Carbon is less reactive than iron but more reactive than copper.

2.

Why can carbon be used to extract iron from its ore?

Carbon is more reactive than iron and can remove oxygen from iron oxide.

3.

What happens to a metal compound when it is heated with carbon?

The metal is released from its compound.

4.

What is the general word equation for extracting a metal from its oxide using carbon?

Metal oxide + carbon → metal + carbon dioxide.

5.

Why can carbon not be used to extract metals that are more reactive than carbon?

Carbon cannot remove oxygen from compounds of metals more reactive than carbon.

6.

How can copper be extracted by heating copper oxide with carbon?

Copper oxide + carbon → copper + carbon dioxide.

C6.1b Explain why and how electrolysis is used to extract some metals from their ores

1.

Why is electrolysis needed to extract metals that are more reactive than carbon?

They are more reactive than carbon and cannot be extracted by carbon reduction.

2.

What happens to metal ions at the cathode during electrolysis?

They gain electrons and form metal atoms.

3.

What happens to negative ions at the anode during electrolysis?

They lose electrons.

4.

Why must an ionic compound be molten or dissolved in water for electrolysis to occur?

The ions must be free to move.

5.

Why is electrolysis an expensive method of extracting metals?

It requires a large amount of electrical energy.

6.

Why is electrolysis used to extract aluminium rather than carbon reduction?

Aluminium is more reactive than carbon.

C6.1c Evaluate alternative biological methods of metal extraction

1.

What is bacterial extraction of metals?

Using bacteria to help extract metals from ores.

2.

How can bacteria help extract metals from ores?

They can react with minerals and release metal ions into solution.

3.

What is phytoextraction?

Using plants to absorb metals from soil.

4.

How can plants be used to extract metals from contaminated soil?

Plants absorb metal ions and concentrate them in their tissues.

5.

What is one advantage of using biological methods to extract metals?

They can use less energy and produce less environmental damage.

6.

What is one disadvantage of using biological methods to extract metals?

They can be slow and may only work for certain metals or concentrations.

C6.1d Describe the basic principles in carrying out a life-cycle assessment of a material or product

1.

What is a life-cycle assessment?

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

2.

What is considered when assessing the production of a material or product?

Raw materials, manufacture, transport, use and disposal.

3.

Why is the transport of a product considered in a life-cycle assessment?

Transport uses energy and can produce emissions.

4.

Why is the use of a product considered in a life-cycle assessment?

Its energy use and environmental effects are considered.

5.

Why is disposal of a product considered in a life-cycle assessment?

Disposal can produce waste or pollution.

6.

What types of environmental impacts can be considered during a life-cycle assessment?

Energy use, resource use, emissions, waste and pollution.

C6.1e Interpret data from a life-cycle assessment of a material or product

1.

What can data from a life-cycle assessment be used to compare?

The environmental impacts of different products or materials.

2.

What does a higher environmental impact score indicate?

Greater environmental impact.

3.

How can a life-cycle assessment help identify the stage with the greatest environmental impact?

Compare the impacts of each stage.

4.

Why should data from all stages of a product's life cycle be considered?

Impacts can occur at any stage.

5.

How could life-cycle assessment data be used to compare two different materials?

Compare the total impacts across their life cycles.

6.

Why can a material with a low impact during production still have a high overall life-cycle impact?

Its use, transport or disposal may have a large environmental impact.

C6.1f Describe a process where a material or product is recycled for a different use, and explain why this is viable

1.

What is recycling?

Processing used materials so they can be used again.

2.

What is meant by recycling a material for a different use?

Turning a used material into a different useful product.

3.

Give one example of a material being recycled for a different use.

Used plastic bottles being made into polyester clothing.

4.

Why does recycling reduce the need for new raw materials?

Less new material needs to be extracted.

5.

Why can recycling reduce the amount of waste sent to landfill?

Less waste is sent to landfill.

6.

What factors determine whether recycling a material for a different use is viable?

Cost, energy use, available technology, environmental benefits and demand for the recycled material.

C6.1g Evaluate factors that affect decisions on recycling

1.

Why does the availability of recycling facilities affect recycling decisions?

Recycling is difficult without somewhere to collect and process materials.

2.

How does the cost of recycling affect whether a material is recycled?

High costs can make recycling less viable.

3.

Why does the energy required for recycling need to be considered?

Recycling can require significant energy.

4.

How does the environmental impact of recycling affect decisions?

The environmental benefits must be weighed against environmental costs.

5.

Why does the demand for recycled materials affect whether recycling is viable?

Recycling is more useful when there is demand for the recycled material.

6.

Why might a material not be recycled even when recycling is technically possible?

It may be too expensive, energy-intensive or have little demand.

C6.1h Describe the separation of crude oil by fractional distillation

1.

What is crude oil?

A mixture of hydrocarbons.

2.

What process is used to separate crude oil into fractions?

Fractional distillation.

3.

What happens to crude oil before fractional distillation?

It is heated so that most of it vaporises.

4.

What are fractions in crude oil?

Mixtures of hydrocarbons with similar boiling points.

5.

Where do the fractions with lower boiling points collect in the fractionating column?

Near the top of the column.

6.

Name four fractions obtained from crude oil by fractional distillation.

Refinery gases, petrol, kerosene and diesel.

C6.1i Explain the separation of crude oil by fractional distillation

1.

Why can the compounds in crude oil be separated by fractional distillation?

They have different boiling points.

2.

How does molecular size affect the boiling point of hydrocarbons?

Larger molecules generally have higher boiling points.

3.

How do intermolecular forces change as hydrocarbon molecules become larger?

They become stronger.

4.

Why do larger hydrocarbon molecules generally have higher boiling points?

More energy is needed to separate larger molecules.

5.

Why do hydrocarbons with similar boiling points collect in the same fraction?

They condense at similar temperatures.

6.

Why do hydrocarbons condense at different heights in a fractionating column?

The temperature changes with height, so different hydrocarbons condense at different levels.

C6.1j Describe the fractions as largely a mixture of compounds of formula CₙH₂ₙ₊₂ which are members of the alkane homologous series

1.

What is the general formula of the alkanes?

CₙH₂ₙ₊₂.

2.

What is a homologous series?

A family of compounds with similar chemical properties and the same general formula.

3.

What is the molecular formula of methane?

CH₄.

4.

What is the molecular formula of propane?

C₃H₈.

5.

How does the formula of one alkane differ from the next member of the series?

Each successive alkane has one more carbon and two more hydrogen atoms.

6.

Why is each crude oil fraction a mixture rather than a single pure substance?

Each fraction contains several hydrocarbons with similar boiling points.

C6.1k Recall that crude oil is a main source of hydrocarbons and is a feedstock for the petrochemical industry

1.

What is crude oil a major source of?

Hydrocarbons.

2.

What is a hydrocarbon?

Compounds containing only carbon and hydrogen.

3.

What is meant by a feedstock?

A raw material used to make other products.

4.

Why is crude oil used as a feedstock in the petrochemical industry?

It provides hydrocarbons used to make many chemicals and materials.

5.

What types of products can be made from hydrocarbons obtained from crude oil?

Fuels, plastics, solvents and other chemicals.

6.

Why are crude oil fractions important raw materials for the chemical industry?

They provide the raw materials needed to manufacture many useful products.

C6.1l Explain how modern life is crucially dependent upon hydrocarbons and recognise that crude oil is a finite resource

1.

Why are hydrocarbons important to modern life?

They provide fuels and raw materials for many products.

2.

Give two uses of hydrocarbons as fuels.

Petrol and diesel.

3.

Give two useful products made from hydrocarbons that are not used as fuels.

Plastics and solvents.

4.

What does it mean if a resource is finite?

It is available in a limited amount and cannot be replaced quickly.

5.

Why is crude oil described as a finite resource?

It takes millions of years to form and is being used much faster than it forms.

6.

Why is dependence on crude oil a long-term problem?

Supplies will eventually become limited or exhausted.

C6.1m Describe the production of materials that are more useful by cracking

1.

What is cracking?

Breaking long-chain hydrocarbons into shorter molecules.

2.

Why is cracking carried out on some long-chain hydrocarbons?

To produce more useful shorter-chain hydrocarbons.

3.

What conditions are needed for cracking?

High temperature and a catalyst, or high temperature and steam depending on the process.

4.

What types of products are produced by cracking?

Shorter-chain alkanes and alkenes.

5.

Why are alkenes produced by cracking useful?

They are used to make polymers and other useful chemicals.

6.

Why is cracking useful when the demand for shorter-chain hydrocarbons is greater than the supply from fractional distillation?

It produces more of the shorter-chain hydrocarbons that are in greater demand.

C6.2a Interpret evidence for how it is thought the atmosphere was originally formed

1.

What gases are thought to have been released by volcanic activity on the early Earth?

Carbon dioxide, water vapour, nitrogen and small amounts of other gases.

2.

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

Volcanic gases were released into the atmosphere.

3.

Which gas was present in very small amounts or absent from the early atmosphere?

Oxygen.

4.

How has the composition of Earth's atmosphere changed over time?

Carbon dioxide decreased and oxygen increased over time.

5.

What evidence can scientists use to work out how Earth's atmosphere changed over time?

Evidence from rocks, fossils and scientific measurements.

6.

Why is the composition of Earth's early atmosphere uncertain?

There is limited direct evidence from billions of years ago.

C6.2b Describe how it is thought an oxygen-rich atmosphere developed over time

1.

Which organisms first produced significant amounts of oxygen?

Photosynthetic microorganisms such as cyanobacteria.

2.

What process carried out by these organisms released oxygen?

Photosynthesis.

3.

How did increasing oxygen levels affect the atmosphere?

Oxygen levels increased.

4.

Why did oxygen levels increase as photosynthetic organisms became more widespread?

More photosynthesis released more oxygen.

5.

What happened to carbon dioxide as photosynthesis increased?

Carbon dioxide levels decreased.

6.

How did the development of plants contribute to an oxygen-rich atmosphere?

Plants carried out photosynthesis, releasing oxygen and removing carbon dioxide.

C6.2c Describe the greenhouse effect in terms of the interaction of radiation with matter within the atmosphere

1.

What is the greenhouse effect?

The warming of Earth's surface and lower atmosphere because greenhouse gases absorb and re-emit infrared radiation.

2.

What type of radiation from the Sun reaches Earth's surface?

Mainly visible and ultraviolet radiation.

3.

What happens to the energy absorbed by Earth's surface?

It warms the Earth's surface.

4.

What type of radiation is emitted by Earth's surface?

Infrared radiation.

5.

How do greenhouse gases interact with infrared radiation?

They absorb and re-emit infrared radiation.

6.

Why does the greenhouse effect warm Earth's surface?

This reduces the rate at which thermal energy escapes into space.

C6.2d Evaluate the evidence for additional anthropogenic (human activity) causes of climate change and describe the uncertainties in the evidence base

1.

What does the term anthropogenic mean?

Caused by human activity.

2.

How does burning fossil fuels affect the concentration of carbon dioxide in the atmosphere?

It increases atmospheric carbon dioxide.

3.

What relationship has been observed between atmospheric carbon dioxide concentration and fossil fuel consumption?

Both have increased over time.

4.

What evidence supports the idea that human activities contribute to climate change?

Rising greenhouse gas concentrations, temperature records and evidence from climate observations and models.

5.

Why can correlation between two changes not by itself prove causation?

Other factors may cause one or both changes.

6.

Why are there uncertainties when predicting the future effects of climate change?

Future climate depends on uncertain future emissions and complex interactions in the climate system.

C6.2e Describe the potential effects of increased levels of carbon dioxide and methane on the Earth’s climate and how these effects may be mitigated

1.

Why are carbon dioxide and methane called greenhouse gases?

They absorb infrared radiation.

2.

What happens to the greenhouse effect when atmospheric carbon dioxide increases?

The greenhouse effect becomes stronger.

3.

What happens to Earth's average temperature if the greenhouse effect becomes stronger?

It increases.

4.

What are two possible effects of increased global temperatures?

Rising sea levels and changes in weather patterns.

5.

How could reducing the burning of fossil fuels help reduce climate change?

It reduces carbon dioxide emissions.

6.

How could increasing the use of renewable energy help mitigate climate change?

Renewable energy produces little or no greenhouse gas emissions during electricity generation.

C6.2f Describe the major sources of carbon monoxide, sulfur dioxide, oxides of nitrogen and particulates in the atmosphere and explain the problems caused by increased amounts of these substances

1.

What is the main source of carbon monoxide from human activity?

Incomplete combustion of fossil fuels.

2.

Why is carbon monoxide harmful to humans?

It reduces the blood's ability to carry oxygen.

3.

What is a major source of sulfur dioxide in the atmosphere?

Burning sulfur-containing fossil fuels.

4.

What environmental problem can sulfur dioxide cause?

Acid rain.

5.

What are major sources of oxides of nitrogen?

High-temperature combustion in engines and power stations.

6.

What health and environmental problems can particulates cause?

They can damage the lungs and contribute to respiratory and cardiovascular problems.

C6.2g Describe the principal methods for increasing the availability of potable water in terms of the separation techniques used

1.

What is potable water?

Water that is safe to drink.

2.

How can filtration be used to make water potable?

It removes insoluble solid particles.

3.

Why is distillation used to obtain pure water from seawater?

To remove dissolved salts and produce pure water.

4.

How does distillation separate water from dissolved salts?

Water is evaporated and then condensed, leaving dissolved salts behind.

5.

What is one advantage of treating water from waste or ground water rather than using seawater?

It generally requires less energy than desalinating seawater.

6.

What factors affect the choice of method used to produce potable water?

Availability of water, cost, energy requirements, local conditions and the level of contamination.

Chemistry

Paper 1

C1 – Particles

C1.1a Describe the main features of the particle model in terms of states of matter and change of state

C1.1b Explain in terms of the particle model the distinction between physical changes and chemical changes

C1.1c Explain the limitations of the particle model in relation to changes of state when particles are represented by inelastic spheres (e.g. like bowling balls)

C2 – Elements, Compounds and Mixtures

C2.1a Explain what is meant by the purity of a substance, distinguishing between the scientific and everyday use of the term ‘pure’

C2.1b Use melting point data to distinguish pure from impure substances

C2.1c Calculate relative formula masses of species separately and in a balanced chemical equation

C2.1d Deduce the empirical formula of a compound from the relative numbers of atoms present or from a model or diagram and vice versa

C2.1e Explain that many useful materials are formulations of mixtures

C2.1f Describe, explain and exemplify the processes of filtration, crystallisation, simple distillation, and fractional distillation

C2.1g Describe the techniques of paper and thin layer chromatography using aqueous and non-aqueous solvents and locating agents

C2.1h Recall that chromatography involves a stationary and a mobile phase and that separation depends on the distribution between the phases

C2.1i Interpret chromatograms, including measuring Rf values

C2.1j Suggest suitable purification techniques given information about the substances involved

C2.1k Suggest chromatographic methods for distinguishing pure from impure substances

C2.2a Describe metals and non-metals and explain the differences between them on the basis of their characteristic physical and chemical properties

C2.2b Explain how the atomic structure of metals and non-metals relates to their position in the Periodic Table

C2.2c Explain how the position of an element in the Periodic Table is related to the arrangement of electrons in its atoms and hence to its atomic number, group number and period number

C2.2d Describe and compare the nature and arrangement of chemical bonds in: i. ionic compounds ii. simple molecules iii. giant covalent structures iv. polymers v. metals

C2.2e Explain chemical bonding in terms of electrostatic forces and the transfer or sharing of electrons

C2.2f Construct dot and cross diagrams for simple covalent and binary ionic substances

C2.2g Describe the limitations of particular representations and models

C2.2h Explain how the reactions of elements are related to the arrangement of electrons in their atoms and hence to their atomic number

C2.2i Explain in terms of atomic number how Mendeleev’s arrangement was refined into the modern Periodic Table

C2.3a Recall that carbon can form four covalent bonds

C2.3b Explain that the vast array of natural and synthetic organic compounds occur due to the ability of carbon to form families of similar compounds, chains and rings

C2.3c Explain the properties of diamond, graphite, fullerenes and graphene in terms of their structures and bonding

C2.3d Use ideas about energy transfers and the relative strength of chemical bonds and intermolecular forces to explain the different temperatures at which changes of state occur

C2.3e Use data to predict states of substances under given conditions

C2.3f Explain how the bulk properties of materials (ionic compounds; simple molecules; giant covalent structures; polymers and metals) are related to the different types of bonds they contain, their bond strengths in relation to intermolecular forces and the ways in which their bonds are arranged

C3 – Chemical Reactions

C3.1a Use chemical symbols to write the formulae of elements and simple covalent and ionic compounds

C3.1b Use the names and symbols of common elements and compounds and the principle of conservation of mass to write formulae and balanced chemical equations and half equations

C3.1c Use the names and symbols of common elements from a supplied Periodic Table to write formulae and balanced chemical equations where appropriate

C3.1d Use the formula of common ions to deduce the formula of a compound

C3.1e Construct balanced ionic equations

C3.1f Describe the physical states of products and reactants using state symbols (s, l, g and aq)

C3.1g Describe tests to identify selected gases

C3.1h Recall and use the definitions of the Avogadro constant (in standard form) and of the mole

C3.1i Explain how the mass of a given substance is related to the amount of that substance in moles and vice versa

C3.1j Explain how the mass of a solute and the volume of the solution is related to the concentration of the solution

C3.1k Recall and use the law of conservation of mass

C3.1l Explain any observed changes in mass in non-enclosed systems during a chemical reaction and explain them using the particle model

C3.1m Deduce the stoichiometry of an equation from the masses of reactants and products and explain the effect of a limiting quantity of a reactant

C3.1n Use a balanced equation to calculate masses of reactants or product

C3.2a Distinguish between endothermic and exothermic reactions on the basis of the temperature change of the surroundings

C3.2b Draw and label a reaction profile for an exothermic and an endothermic reaction

C3.2c Explain activation energy as the energy needed for a reaction to occur

C3.2d Calculate energy changes in a chemical reaction by considering bond making and bond breaking energies

C3.3a Explain reduction and oxidation in terms of loss or gain of oxygen, identifying which species are oxidised and which are reduced

C3.3b Explain reduction and oxidation in terms of gain or loss of electrons, identifying which species are oxidised and which are reduced

C3.3c Recall that acids form hydrogen ions when they dissolve in water and solutions of alkalis contain hydroxide ions

C3.3d Describe neutralisation as acid reacting with alkali or a base to form a salt plus water

C3.3e Recognise that aqueous neutralisation reactions can be generalised to hydrogen ions reacting with hydroxide ions to form water

C3.3f Recall that carbonates and some metals react with acids and write balanced equations predicting products from given reactants

C3.3g Use and explain the terms dilute and concentrated (amount of substance) and weak and strong (degree of ionisation) in relation to acids

C3.3h Recall that relative acidity and alkalinity are measured by pH

C3.3i Describe neutrality and relative acidity and alkalinity in terms of the effect of the concentration of hydrogen ions on the numerical value of pH

C3.3j Use the idea that as hydrogen ion concentration increases by a factor of ten, the pH value of a solution decreases by one

C3.3k Describe techniques and apparatus used to measure pH

C3.4a Recall that metals (or hydrogen) are formed at the cathode and non-metals are formed at the anode in electrolysis using inert electrodes

C3.4b Predict the products of electrolysis of binary ionic compounds in the molten state

C3.4c Describe competing reactions in the electrolysis of aqueous solutions of ionic compounds in terms of the different species present

C3.4d Describe electrolysis in terms of the ions present and reactions at the electrodes

C3.4e Describe the technique of electrolysis using inert and non-inert electrodes

C4 – Predicting and Identifying Chemical Products

C4.1a Recall the simple properties of Groups 1, 7 and 0

C4.1b Explain how observed simple properties of Groups 1, 7 and 0 depend on the outer shell of electrons of the atoms and predict properties from given trends down the groups

C4.1c Predict possible reactions and probable reactivity of elements from their positions in the Periodic Table

C4.1d Explain how the reactivity of metals with water or dilute acids is related to the tendency of the metal to form its positive ion

C4.1e Deduce an order of reactivity of metals based on experimental results

Paper 2

C5 – Monitoring and Controlling Chemical Reactions

C5.1a Suggest practical methods for determining the rate of a given reaction

C5.1b Interpret rate of reaction graphs

C5.1c Describe the effect of changes in temperature, concentration, pressure, and surface area on rate of reaction

C5.1d Explain the effects on rates of reaction of changes in temperature, concentration and pressure in terms of frequency and energy of collision between particles