Topic 1: Atomic Structure
Specification points
Atoms, elements and compounds
· Use the names and symbols of the first 20 elements in the periodic table and the elements in Groups 1 and 7.
· Name compounds of these elements from given formulae or symbol equations.
Mixtures
· Describe, explain and give examples of the specified processes of separation.
· Suggest suitable separation and purification techniques for mixtures when given appropriate information.
The development of the model of the atom
· Describe why the new evidence from the scattering experiment led to a change in the atomic model.
· Describe the difference between the plum pudding model of the atom and the nuclear model of the atom.
Relative electrical charges of subatomic particles
· Use the nuclear model to describe atoms (relative charges and atomic number)
Size and mass of atoms
· Calculate the number of protons, neutrons and electrons in an atom or ion, given its atomic number and mass number.
· Relate size and scale of atoms to objects in the physical world.
Relative atomic mass
· Calculate the relative atomic mass of an element given the percentage abundance of its isotopes.
Electronic structure
· Draw a dot-and-cross diagram of the electronic structures for the first 20 elements and their ions.
· Work out the electronic configuration of the first 20 elements and their ions.
Keywords
Definitions
Atom
An atom is the smallest part of an element that can exist
Compound
Two or more different atoms chemically joined together
Element
Made up of one type of atom
Molecule
Two or more atoms chemically joined together
Mixture
Consists of two or more elements or compounds not chemically joined together
Atomic number/proton number
Number of protons an atom of an element contains
Relative atomic mass
Average mass of an element, taking into account the abundances of the isotopes of the element, compared to 1/12th of 12-C atom
Topic 2: The Periodic Table
Specification points
The Periodic Table
· 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.
· Predict possible reactions and probable reactivity of elements from their positions in the periodic table.
Development of the Periodic Table
· Describe the steps in the development of the Periodic Table.
Metals and non-metals
· Explain the differences between metals and non-metals on the basis of their characteristic physical and chemical properties. This links to Group 0, Group 1, Group 7 and bonding, structure and the properties of matter.
· Explain how the atomic structure of metals and non-metals relates to their position in the periodic table.
· Explain how the reactions of elements are related to the arrangement of electrons in their atoms and hence to their atomic number.
Group 0, Group 1 and Group 7
· Explain how properties of the elements in Group 0, Group 1 and Group 7 depend in the outer shell of electrons of the atoms.
· Predict properties from given trends down the group.
Properties of transition metals
· Describe the difference compared with Group 1 in melting points, densities, strength, hardness and reactivity with oxygen, water and halogens.
· Exemplify these general properties by reference to Cr, Mn, Fe, Co, Ni, Cu.
Topic 3: Bonding & Structure
Specification points
Chemical bonds
· Explain chemical bonding in terms of electrostatic forces and the transfer or sharing of electrons.
Ionic bonding
· Draw dot-and-cross diagrams for ionic compounds formed by metals in Groups 1 and 2 with non-metals in Groups 6 and 7.
· Work out the charge on the ions of metals and non-metals from the group number of the element, limited to the metals in Groups 1 and 2, and non-metals in Groups 6 and 7.
Ionic compounds
· Deduce that a compound is ionic from a diagram of its structure in one of the specified forms.
· Describe the limitations of using dot-and-cross, ball and stick, two and three-dimensional diagrams to represent a giant ionic structure.
· Work out the empirical formula of an ionic compound from a given model of diagram that shows the ions in the structure.
· Be familiar with the structure of sodium chloride.
· Describe and explain the properties of ionic compounds.
Covalent bonding
· Draw dot-and-cross diagrams for the molecules of hydrogen, chlorine, oxygen, nitrogen, hydrogen chloride, water, ammonia and methane.
· Represent the covalent bonds in small molecules, in the repeating units of polymers and in part of giant covalent structures, using a line to represent a single bond.
· Describe the limitations of using dot-and-cross, ball and stick, two and three-dimensional diagrams to represent molecules or giant structures.
· Deduce the molecular formula of a substance from a given model or diagram in these forms showing the atoms and bonds in the molecule.
Properties of small molecules
· Use the idea of intermolecular forces are weak compared with covalent bonds to explain the bulk properties of molecular substances.
Giant covalent structures
· Recognise giant covalent structures from diagrams showing their bonding and structure.
· Explain the properties of giant covalent structures for diamond, graphite and silicon dioxide (silica) in terms of their structure and bonding.
· Know that graphite is similar to metals in that it has delocalised electrons.
· Recognise graphene and fullerenes from diagrams and descriptions of their bonding and structure.
· Give examples of the uses of fullerenes, including carbon nanotubes.
Polymers
· Recognise polymers from diagrams showing their bonding and structure.
· Describe the bonding in polymers and relate their bonding to their solid state at room temperature.
Metallic bonding
· Draw a diagram to represent metallic bonding.
Properties of metals and alloys
· Explain why alloys are harder than pure metals in terms of distortion of the layers of atoms in the structure of a pure metal.
· Explain why metals are good conductors of electricity.
The three states of matter
· Predict the states of substances at different temperatures given appropriate data.
· Explain the different temperatures at which changes of state occur in terms of energy transfers and types of bonding.
· Recognise that atoms themselves do not have the bulk properties of materials.
· Explain the limitations of the particle theory in relation to changes of state when particles are represented by solid inelastic spheres which have no forces between them.
Nanoparticles
· Compare ‘nano’ dimensions to typical dimensions of atoms and molecules.
· Given appropriate information, evaluate the use of nanoparticles for a specified purpose.
· Explain that there are possible risks associated with the use of nanoparticles.
Keywords
Definitions
Ionic bonding
Occurs between metals and non-metals. Electrostatic attraction between oppositely charged ions, formed by the transfer of electrons.
Covalent bonding
Occurs between non-metals. Sharing of a pair of electrons between two positive nuclei.
Metallic bonding
Occurs in metals. Electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons
Nanoparticle
1-100 nm in size, of the order of a few hundred atoms
Topic 4: Chemical calculations
Specification points
Conservation of mass and balanced chemical equations
· Understand the use of multipliers in equations in normal script before a formula and in subscript within a formula.
· Write word equations for chemical reactions.
· Write formulae and balanced chemical equations, including appropriate state symbols.
· Explain any observed changes in mass in non-enclosed systems during a chemical reaction given the balanced symbol equation for the reaction and explain these changes in terms of the particle model.
Relative formula mass
· Calculate the relative formula mass of given formulae.
Chemical measurements
· Represent the distribution of results and make estimations of uncertainty.
· Use the range of a set of measurements about the mean as a measure of uncertainty.
Moles
· Understand that the measurement of amounts in moles can apply to atoms, molecules, ions, electrons, formulae and equations.
· Use the relative formula mass of a substance to calculate the number of moles in a given mass of that substance and vice versa.
Reacting masses
· Calculate the masses of substances shown in a balanced symbol equation.
· Calculate the masses of reactants and products from the balanced symbol equation and the mass of a given reactant or product.
· Balance an equation given the masses of reactants and products.
· Change the subject of a mathematical equation.
Limiting reactants
· Explain the effect of a limiting quantity of a reactant on the amount of products it is possible to obtain in terms of amounts in moles and masses in grams.
Concentration of solutions
· Calculate the mass of a solute in a given volume of solution of known concentration in terms of mass per given volume of solution.
· Explain how the mass of a solute and the volume of a solution is related to the concentration of the solution.
· Explain how the concentration of a solution in mol/dm3 is related to the concentration of a solution in g/dm3.
Volume of gases
· Calculate the volume of a gas at room temperature and pressure from its mass and relative formula mass.
· Calculate volumes of gaseous reactants and products from a balanced equation and a given volume of a gaseous reactant or product.
Percentage yield
· Calculate the percentage yield of a product from the actual yield of a reaction.
· Calculate the theoretical mass of a product from a give mass of reactant and the balanced equation for the reaction.
Atom economy
· Calculate the atom economy of a reaction to form a desired product from the balanced equation.
· Explain why a particular reaction pathway is chosen to produce a specified product given appropriate data such as atom economy, yield, rate, equilibrium position and usefulness of by-products.
Titrations
· Describe how to carry out titrations using strong acids (sulphuric, hydrochloric and nitric acids only) and strong alkalis only to find the reacting volumes accurately.
· Calculate the chemical quantities in titrations involving concentrations in mol/dm3 and in g/dm3.
Keywords
Definitions
Conservation of mass
The law of the conservation of mass states that no atoms are lost or made during a chemical reaction so that the mass of the products equals the mass of the reactants.
Relative formula mass (Mr)
Sum of the relative atomic masses of the atoms in the numbers shown in the formula
Mole
Unit of measurement for atoms, molecules or ions.
Avogadro’s constant
Number of atoms, molecules or ions in a mole of a given substance. 6.02 x 1023
Key equations/formulae to learn
mass = moles x Mr or mass = moles x Ar
Percentage yield = Experimental mass x 100
Theoretical mass
Atom economy = Mr of desired product x 100
Sum Mr of all reactants
Moles = concentration x volume (dm3)
Gas volume in dm3 = moles x 24 dm3
Topic 5: Chemical changes
Specification points
The Reactivity Series
· Recall and describe the reactions, if any, of potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper with water or dilute acids and where appropriate, to place these metals in order of reactivity.
· Explain how the reactivity of metals with water or dilute acids is related to the tendency of the metal to form its positive ion.
· Deduce an order of reactivity of metals based on experimental results.
Extraction of metals
· Describe how metals less reactive than carbon can be extracted from their oxides by reduction with carbon.
· Interpret or evaluate specific metal extraction processes when given appropriate information.
· Identify the substances which are oxidised or reduced in terms of gain or loss of oxygen.
REDOX
· Explain reduction and oxidation in terms of loss or gain of oxygen and in terms of electrons.
· Write ionic equations for displacement reactions.
· Identify in a given reaction symbol equation or half equation which species are oxidised and which are reduced.
Reaction of acids
· Know acids react with some metals to produce salts and hydrogen.
· Know acids react with alkalis and bases to produce salts and water.
· Know acids react with metal carbonates to produce salts, water and carbon dioxide.
· Predict products from given reactants.
· Use the formulae of common ions to deduce the formulae of salts.
Soluble salts
· Describe how to make pure, dry samples of named soluble salts.
The pH scale and neutralisation
· Describe the use of universal indicator to measure the approximate pH of a solution.
· Use the pH scale to identify acidic or alkaline solutions.
· Recall the equation for a neutralisation reaction.
Titrations
· Describe how to carry out titrations using strong acids (sulfuric, hydrochloric and nitric acid only) and strong alkalis only to find the reacting volumes accurately.
· Calculate the chemical quantities in titrations involving concentrations in mol/dm3 in g/dm3.
Strong and weak acids
· Use and explain the terms dilute and concentrated (in terms of amount of substance) and weak and strong (in terms of degree of ionisation) in relation to acids.
· Describe neutrality and relative acidity in terms of the effect on hydrogen ion concentration and the numerical value of pH (whole numbers only).
Keywords
Definitions
Oxidation
Gain of oxygen or loss of electrons
Reduction
Loss of oxygen or gain of electrons
Key equations/formulae to learn
Acid + Metal Salt + Hydrogen
Acid + Alkali (e.g. metal hydroxide) Salt + Water
Acid + Metal carbonate Salt + Water + Carbon dioxide
H+ (aq) + OH- (aq) H2O (l)
Moles = concentration x volume (dm3)
Required practicals
Summary of method including reactions and results (6 marks)
1A: Making a soluble salt from a metal oxide
About 40cm3 of sulfuric acid was gently heated until almost boiling. Copper oxide was added in small amounts until it was in excess. The excess copper oxide was filtered off. The remaining copper sulfate solution was gently heated until crystals formed. It was left in a cool place to crystallise. Blue crystals of hydrated copper sulfate were formed.
Copper oxide + Sulfuric acid Copper sulphate + Water
CuO (s) + H2SO4 (aq) CuSO4 (aq) + H2O (l)
Copper oxide + Sulfuric acid Hydrated copper sulfate
CuO (s) + H2SO4 (aq) + 4H2O (l) CuSO4.5H2O (s)
Anhydrous copper sulfate + Water Hydrated copper sulphate
CuSO4 (s) + 5 H2O (l) CuSO4.5H2O (s)
1B: Making a soluble salt from a metal carbonate
Same as above but with magnesium carbonate and magnesium sulfate crystals form.
Magnesium carbonate + Sulfuric acid Magnesium sulphate + Carbon dioxide
MgCO3 (s) + H2SO4 (aq) MgSO4.7H2O (s) + CO2 (g)
Magnesium Carbonate + Sulfuric acid + water Hydrated Magnesium sulphate + CO2
MgCO3 (s) + H2SO4 (aq) + 6H2O (l) MgSO4.7H2O(s) + CO2 (g)
1C: Making an insoluble salt from soluble salts
5cm3 of lead nitrate solution was mixed with 10cm3 of sodium chloride solution. The solution was filtered. The precipitate was washed. The precipitate was heated in suspension with water until the suspension dissolved. It was allowed to cool.
White precipitate of lead chloride was formed.
Lead Nitrate + Sodium chloride Lead Chloride + Sodium nitrate
Pb(NO3)2 (aq) + 2NaCl (aq) PbCl2 (s) +2NaNO3 (aq)
Pb2+ (aq) +2Cl- (aq) PbCl2 (s)
2: Neutralisation
Dilute sulfuric acid was gradually added to 25 cm3 of sodium hydroxide with methyl orange indicator. Sulfuric acid was continued to be added until the indicator changed from yellow to orange. This was repeated several times until concordant results were obtained.
The methyl orange indicator turned from yellow to orange.
H+ + OH- H2O
Topic 6: Electrolysis
Specification points
The process of electrolysis
· Define electrolysis and electrolyte.
Electrolysis of molten ionic compounds
· Predict the products of the electrolysis of ionic compounds in the molten state.
Using electrolysis to extract metals
· Explain why a mixture of aluminium oxide and cryolite is used as the electrolyte.
· Explain why the positive electrode must be continually replaced.
· Interpret or evaluate specific metal extraction processes when given appropriate information.
Electrolysis of aqueous solutions
· Predict the products of the electrolysis of aqueous solutions containing a single ionic compound.
Half equations
· Write balanced half equations for the reactions occurring at the electrodes during electrolysis.
Keywords
Definitions
Electrolysis
The chemical breakdown of a substance (usually ionic) by passing through an electric current
Electrolyte
Liquids and solutions able to conduct electricity which are broken down by passing an electric current through.
Cation
Positive ion
Anion
Negative ion
Cathode
Negative electrode
Anode
Positive electrode
Required practicals
Summary of method including reactions and results (6 marks)
3: Electrolysis of aqueous solution
Pour electrolyte (copper (II) chloride) solution into a beaker. Insert carbon electrodes and connect to DC voltage supply. Turn voltage on and observe.
Electrolysis of copper chloride produces copper at the cathode and chlorine gas at the anode. Blue litmus paper bleaches.
Electrolysis of copper sulfate produces copper at the cathode and oxygen gas at the anode. Blue litmus paper stays blue.
Electrolysis of sodium chloride produces hydrogen gas at the cathode and chlorine gas at the anode. Blue litmus paper bleaches.
Electrolysis of sodium sulfate produces hydrogen gas at the cathode and oxygen gas at the anode. Blue litmus paper stays blue.
Topic 7: Energy Changes
Specification points
Energy transfer during exothermic and endothermic reactions
· Distinguish between exothermic and endothermic reactions on the basis of the temperature change of the surroundings.
· Evaluate uses and applications of exothermic and endothermic reactions given appropriate information.
Reaction profiles
· Draw simple reaction profiles/energy level diagrams for exothermic and endothermic reactions showing the relative energies of reactants and products, the activation energy and the overall energy change, with a curved line to show the energy as the reaction proceeds.
· Use reaction profiles to identify reactions as exothermic or endothermic.
· Explain that the activation energy is the energy needed for a reaction to occur.
Energy change of reactions
· Calculate the energy transferred in chemical reactions using the bond energies supplied.
Chemical cells and fuel cells
· Interpret data for relative reactivity of different metals.
· Evaluate the use of hydrogen fuel cells in comparison with rechargeable cells and batteries.
· Write half equations for the electrode reactions in the hydrogen fuel cell.
Keywords
Definitions
Exothermic
Energy release so surrounding temperature increases
Endothermic
Energy taken in so surrounding temperature decreases
Bond energy
Energy required to break/make a bond
Activation energy
The minimum amount of energy needed for a reaction to occur
Key equations/formulae to learn
Overall energy change = Energy to break bonds – Energy released to make bonds
Required practicals
Summary of method including reactions and results (6 marks)
4: Temperature Changes
Put 30 cm3 dilute hydrochloric acid into a polystyrene cup and measure the temperature. Put 5 cm3 of sodium hydroxide into the cup, stir and measure the temperature. Repeat until a total of 40 cm3 of sodium hydroxide is added.
Sodium hydroxide + Hydrochloric acid Sodium Chloride + Water
NaOH (aq) + HCl (aq) NaCl (aq) + H2O (l)
Topic 8: Rate of reactions, reversible reactions and dynamic equilibrium
Specification points
Calculating rate of reactions
· Calculate the mean rate of reaction from given information about the quantity of a reactant used or the quantity of a product formed and the time taken.
· Draw and interpret graphs showing the quantity of product formed or quantity of reactant used up against time.
· Draw tangents to the curves on these graphs and use the slope of the tangent as a measure of the rate of reaction.
· Calculate the gradient of a tangent to the curve on these graphs as a measure of rate of reaction at a specific time.
Factors affecting the rates of chemical reactions
· Recall how changing the concentration, pressure, surface area, temperature and the presence of catalysts affect the rate of chemical reactions.
Collision theory
· Predict and explain using collision theory the effects of changing conditions of concentration, pressure and temperature on the rate of reaction.
· Predict and explain the effects of changes in the size of pieces of a reacting solid in terms of surface area to volume ratio.
· Use simple ideas about proportionality when using collision theory to explain the effect of a factor on the rate of a reaction.
Catalysts
· Identify catalysts in reactions from their effect on the rate of reaction.
· Explain catalytic action in terms of activation energy.
Reversible reactions
· Define and identify reversible reactions.
· Know the symbol for reversible reactions.
Equilibrium
· Define dynamic equilibrium.
· Make qualitative predictions about the effects of changes on systems at equilibrium when given appropriate information.
· Interpret appropriate given data to predict the effect of a change in concentration of a reactant or product, change in temperature and change in pressure, on given reactions at equilibrium.
Keywords
Definitions
Exothermic
Energy release so surrounding temperature increases
Endothermic
Energy taken in so surrounding temperature decreases
Bond energy
Energy required to break/make a bond
Activation energy
The minimum amount of energy needed for a reaction to occur
Catalyst
Increases the rate of reaction by providing an alternative pathway with a lower activation energy, without getting used up in the reaction
Closed system
No reactants or products can enter or leave the reaction vessel
Dynamic equilibrium
In a closed system, the forward and backward reaction occur at the same rate and the net concentration of reactants and products remains constant
Key equations/formulae to learn
Mean rate of reaction = Quantity of reactant used
Time taken
Mean rate of reaction = Quantity of product formed
Time taken
Required practicals
Summary of method including reactions and results (6 marks) Include diagrams
5a: Rates of reaction (observing colour change)
Put 10 cm3 of sodium thiosulfate solution into a conical flask. Add 40 cm3 of water to dilute the sodium thiosulfate solution to a concentration of 8 g/dm3. Put the conical flask on the black cross. Add 10 cm3 of dilute hydrochloric acid into the flask, gently swirl and start the stopclock. Stop the clock when you can no longer see the cross. Repeat for other concentrations of sodium thiosulfate.
Sodium thiosulfate + Hydrochloric acid Sodium chloride + Sulfur + Sulfur dioxide + Water
Na2S2O3 (aq) + HCl (aq) NaCl (aq) + S (s) + SO2 (g) + H2O (l)
5b: Rates of reaction (measuring volume of gas produced)
Set up apparatus as shown
Add magnesium ribbon to hydrochloric acid in conical flask. Record volume of gas produced at suitable intervals.
