Popular exam questions chemistry:

Electrolysis:
• Give an example of inert electrode:
• Carbon.
• Platinum
• Graphite.

• Define the Term electrolysis:

• The breakdown of an ionic compound (molten or in aqueous) using
electricity.

• State the function of electrodes:

• To conduct electricity.

• State the use of heat in electrolysis of molten electrolyte:

• To melt the electrolyte.

• Platinum is a good conductor of electricity, State one other property of

platinum which makes it suitable for use as electrodes:
• Its inert.

• Suggest why the electrolyte is a paste:

• To allow ions to flow.

• Name the type of particles responsible for the conduction of electricity

during electrolysis in the metal wires and electrolyte respectively:
• Metal wires: electrons.
• Electrolyte: ions.

• Copper (II) sulfate solution conducts electricity but solid copper (II)
sulfate does not. Explain this statement:
• Ions are free to move in the solution.
• Ions are in lattice form in solid or ions are in fixed positions in solid.

• Describe the electrolysis of concentrated aqueous potassium bromide.

Include an ionic half equation for the reaction at the cathode, the name of the
product at the anode, and the name of the potassium compound formed.
• 2H+ +2e H2
• Bromine gas was formed at the anode.
• Potassium hydroxide.

• When molten potassium bromide is electrolysed, the product at the

cathode is different. Name the product at the Cathode when molten potassium
bromide is electrolysed:
• Potassium
 • What difference would the student observe at the positive electrode if
the aqueous copper (II) sulfate was replaced by concentrated aqueous copper (II)
chloride:
• Yellow gas.

• Name the products of the electrolysis of concentrated aqueous sodium

chloride:
• Hydrogen (cathode)
• Chlorine gas (anode)
• Sodium hydroxide

• The masses of copper electrodes changed during electrolysis. State

how and explain why the masses of the two copper electrodes changed. Use ionic
half-equations to help you:
• The mass of the anode decreases.
• Copper is removed from anode and becomes aqueous. Cu Cu2+ + 2e-
• The mass of cathode increases.
• Copper solid deposited on cathode. Cu2+ + 2e- Cu

• Explain why, during the electrolysis, the color of copper (II) sulfate
solution does not change:
• The rate of formation copper (II) ions is the same rate as the removal of
copper (II) ions.

• Give reasons why steel is electroplated with metal such as chromium:

• To resist corrosion/rusting.
• To improve the appearance.

• Name the ore of aluminum:

• Bauxite.

• Describe how aluminum is extracted from bauxite. Include an ionic

half-equation for the reaction at each electrode:
• Bauxite is dissolved in molten cryolite.
• Cryolite lowers the melting point and improves conductivity.
• Molten aluminum is formed.
• Cathode reaction: Al3+ + 3e- Al3+
• Anode reaction: 2O2- O2 + 4e

• Explain why the anodes have to be replaced regularly:

• Carbon anode will react with oxygen to produce carbon dioxide which can
lead to corrosion.

• Why is aluminum not extracted from its ore by reduction with carbon:
• Aluminum is more reactive than carbon.
 • Give 2 reasons why cryolite is used:
• To lower the melting point.
• To increase conductivity.

• During the extraction of aluminum by electrolysis, carbon dioxide is

found at the anode. Explain how carbon dioxide is formed at the anode.
• Anode is made of carbon.
• Oxygen gas is produced at the anode.
• Carbon anode reacts with oxygen gas to produce carbon dioxide.
• Is electrolysis reaction endothermic of exothermic? Explain your answer:
• Endothermic because electrical energy is supplied.

• Is the reaction in simple cell endothermic or exothermic? Explain

your answer.
• Exothermic because electrical energy is released.

Sulfur and carbonate

• Give a use of Sulphur dioxide:
• Bleaching
• Food preservative
• Killing bacteria.

• Name a source of sulfur:

• Fossil fuels.

• Describe how sulfur is converted into sulfur dioxide.

• Heated/roasted/burned in the air.

• What is added to sulfur dioxide to convert it into oleum:

• Concentrated sulfuric acid.

• What is added to oleum to convert it into sulfuric acid:

• Water

• Give one use of sulfuric acid:

• Detergents
• Car batteries
• Dyes
• Sulfur is burned by spraying droplets of molten sulfur into the air. Suggest
and explain an advantage of using this method:
• Larger surface area.
• Faster reaction rate.

• Sulfur dioxide is more expensive than air. What is the advantage of

 using excess of air?
• Move equilibrium to right.
• Increase the yield sulfur trioxide.

• Describe how concentrated sulfuric acid is made from sulfur trioxide:

• React sulfur trioxide with concentrated sulfuric acid.
• Then add water.

• Sulfur dioxide is a reductant. Describe what you would see when

aqueous sulfur dioxide is added to acidified potassium manganate (VII)
• Purple to colorless.

Air and water

• Cobalt (II) chloride test for water:
• Blue to pink in the presence of water.

• Anhydrous copper (II) sulfate test for water:

• White to blue.

• State two industrial uses of water:

• Manufacture of ethanol.
• Manufacture of hydrogen or ammonia for the Haber process.

• State two uses of water in the home:

• Cooking.
• Drinking
• Washing.

• Rain water is collected in reservoirs. How is it treated before entering

the water supply?
• Filtration.
• Chlorination.

• In many regions, drinking water is obtained by the distillation of sea-

water. Explain how distillation separates the water from sea-water.
• Boiling or turning into steam.
• Condensation.

• Describe how high temperature in vehicles cause acid rain:

• Oxygen and nitrogen from air react.
• Oxides of nitrogen or nitrogen oxides are formed.
• Nitrogen oxides formed react with water to form acids (nitrous oxide.)

• One cause of acidity in soil is acid rain. Explain how acid ran is
formed?
• Burning of fossil fuels that contain sulfur.
 • Formation of sulfur dioxide in the air.
• React with rain water to form sulfuric acid, H2SO4.

• Name two bases which are used to increase the pH of acidic soils:
• Calcium oxide
• Calcium hydroxide.
• Calcium carbonate.

• Why does the combustion of fuel rather than coal reduce the amount
of pollution?
• Wood burns to produce less carbon dioxide.
• When the trees grow, they take in carbon dioxide for photosynthesis.
• Therefore, wood is a carbon neutral.

• Explain why the uncoated piece started to rust:

• Steel or iron is exposed to oxygen and water.

• Explain in terms of electron transfer why steel doesn’t rust eve if the
layer of zinc is scratched so that the steel is exposed to salt and water:
• Zinc is more reactive than iron.
• It loses electrons more readily than iron.
• Iron gains electrons.
• Therefore, iron doesn’t lose electrons.

• How are waste gases formed?

• Carbon monoxide
• From incomplete combustion of carbon containing fuels.
• Sulfur dioxide
• From burning fossil fuels.
• Oxides of nitrogen.
• Nitrogen reacting with oxygen in car engines/lightning.
• Methane
• From anaerobic respiration.

• Describe the Haber process giving conditions and a balanced

equation:
• Pressure of 200atm.
• Temperature of 450 degrees Celsius.
• Iron catalyst.
• N2 + 3H2 ⇌ 2NH3

• Name the elements essential for plant growth commonly found in

fertilizers.
• Nitrogen
• Phosphorus
• Potassium.
Atoms, elements and compounds.
• What is ionic bonding:
• Electrostatic forces of attraction.
• Between oppositely charged ions.

• Describe what happens in terms of electron loss and gain when

potassium reacts with iodine:
• Moment of electron from potassium to iodine.
• One electron is transferred.

• Why isotopes have the same chemical properties:

• Same number of electrons.
• At the outermost electron shells.

• Describe the structure of solid potassium iodide:

• Giant lattice of alternating.
• Positive potassium ions and negative iodine ions.

• Explain why potassium has a high melting point:

• Strong electrostatic forces between oppositely charged ions.
• Bonds are strong which require lots of heat energy to overcome the forces.

• Explain why silicon (IV) oxide has a very high melting point:
• All bonds are strong.
• A lot of heat energy is needed to overcome the bonds.

• Describe the structure of silicon (IV) oxide.

• Tetrahedral.
• Each oxygen is joined to 2 silicon atoms.
• Each silicon is joined to 4 oxygen atoms

• State 3 properties which silicon (IV) oxide and diamond have in

common.
• High melting and boiling point.
• Hard
• Strong Poor conductor of electricity.

• Explain why the physics properties of carbon dioxide are different

from those of diamond and silicon (IV) oxide.
• Carbon dioxide has a simple molecular structure.

Metals
• Give a use of mild steel:
• Car bodies
• Chains
 • Nails.

• Give a use of high carbon steel:

• Knives
• Drills
• Cutting tools.

• Brass is an alloy which contains zinc, name the other metal in brass:
• Copper.

• Suggest 2 reasons why an alloy is preferred to either of its constituent

metals:
• Harder
• Stronger
• Resistance to corrosion
• Better appearance.

• Explain why metals are malleable:

• Metal ions
• Arranged in a lattice.
• The layers or rows slide over one another.

• Explain why alloys are less malleable than metals:

• New atoms are different sizes.
• They distort the regular arrangement of ions and make it more difficult for
the layers to slide over each other.

• Describe the extraction of iron with equations [5 marks]

• One redox equation:
• Fe2O3 + 3CO 2Fe + 3CO2

• One acid/base equation:

• Cao + SiO2 CaSiO3

• Three comments from:

• Carbon reacts to form carbon dioxide.
• The reaction is exothermic and produces heat.
• Carbon dioxide is reduced to carbon monoxide.
• Carbon monoxide reduces hematite to iron
• Carbon reduces hematite to iron.
• Limestone removes silica to form slag
• Limestone decomposes.

• State the function of the coke used in the blast furnace:

• As a reducing agent,
 • Source of heat.

• Why is there no immediate reaction or observation when aluminum

metal is added to HCL:
• Protective oxide layer.

• Explain how the silica impurity is removed and separated from the
molten iron.
• Silica reacts with limestone or calcium oxide.
• To form slag.
• Molten slag floats on molten iron.

• Explain how the amount of carbon in the iron can be decreased:

• Oxygen blown in
• Carbon dioxide forms
• C + O2 CO2

• Describe how zinc oxide is made from zinc blende:

• Zinc blende is roasted in air.
• Zinc sulfide + Oxygen Zinc oxide + Sulfur dioxide

• Write a word equation for the reduction of zinc oxide by coke:

• Zinc oxide + carbon Zinc + Sulfur dioxide

• Describe the extraction of zinc from its ore, zinc blende. Your answer
should include details of how the heat is produced and equations for all the reactions
you describe:
• Zinc sulfide is roasted in the air.
• Zinc oxide is formed.
• Zinc oxide is reduced.
• By coke or carbon
• Heat produced by carbon burning in air.
• 2ZnS + 3O2 2ZnO + 2SO2.

The Periodic table

• What is the relationship between group numbers and number of
valency electrons?
• No. of valency electrons is the same as the group number

• Describe how the type of oxide changes across period.

• Basic to Amphoteric to acidic.

• Predict 2 physical properties of Group 1 metals that are different from

transition metals:
• Has lower melting and boiling point.
 • Soft and can be cut easily.
• Has lower density.

• Predict 3 chemical properties of transition metals that are different

from Group 1 metals:
• More than 1 oxidation state.
• Form colored compounds.
• Can act as a catalyst.

• Observations when group 1 metals are added to cold water.

• Bubbles/effervescence of hydrogen
• Explosion
• Dissolves and forms a solution.

• Explain why noble gases are chemically inert.

• Full outer shell of electrons.

Chemical energetics
• Based on the energy diagram, explain why the reaction is exothermic:
• Chemical energy of reactant is higher than the chemical energy of
products.

• Based on the energy diagram, explain why the reaction is

endothermic:
• Chemical energy of reactant is lower than the chemical energy of products.

• Define the term activation energy:

• The minimum energy needed by reactants to start reaction.

• What is the advantage and disadvantage of using hydrogen fuel cell:

• Advantage: No pollutant produced.
• Disadvantage: Hard to store.

Chemical reactions
• Suggest a way of increasing the reaction using the same amounts of
reactants:
• Reactant should be powdered.
• Higher conc. of solution used.
• Increase temperature.
• Use a catalyst.

• What would be the effect of volume of products if the mass of catalyst

 was increased:
• No change.

• Explain why an increased concentration increases the rate of reaction:

• More particles are closer together.
• Increases the rate of collisions.

• Explain why an increase in temperature increases the rate of reaction:

• Particles gain more kinetic energy.
• Particles move faster.
• Increasing rate of collisions.
• Higher proportions of particles have sufficient energy to react.

• Define equilibrium:
• Rate of forward reaction is the same as rate of backward reaction.
• Conc. of reactants and products remain constant.

• Define in terms of electron transfer the term oxidation:

• Loss of electron.

• Explain why positive ions are oxidizing agents:

• They can accept electrons.

Organic chemistry.
• Fraction X is used a jet fuel, name fraction x:
• Kerosene.

• What happens to petroleum before it is separate into useful

substances.
• It is heated.

• State 3 differences between bitumen and diesel:

• Bitumen has higher m.p. and b.p.
• Bitumen has larger molecular size.
• Bitumen is more viscous.
• Bitumen is darker in color.

• Give 3 characteristics of a homologous series:

• Same general formula.
• Similar chemical properties.
• Contain the same functional group.
• Consecutive members differ by CH2.

• Why are the following isomers?

• Have the same molecular formula.
 • Different structure.

• Define the term hydrocarbon.

• A compound only containing hydrogen and carbon.

• What is the empirical formula of hexene, C6H12?

• CH2

• Give uses of ethanol.

• Fuel
• Making esters
• Perfumes
• Name the 2 products made during the complete combustion of alcohol:
• Carbon dioxide
• Water.

• How to differentiate between alkane and alkene.

• Test with bromine water.
• Alkane – solutions remain the same.
• Alkene – solution turns from brown to colorless.

• What is the source of alkane and alkene respectively?

• Alkane – fractional distillation of petroleum.
• Alkene – cracking of an alkane.

• What is the difference between alkane and alkene?

• Alkane - saturated, contains only single bonds.
• Alkene – unsaturated, contains double bonds.

• Explain the difference between the 2 types of polymerization:

• Addition – Polymer is the only product.
• Condensation – Polymer and a by-product (water) formed.

• Describe pollution caused by non-biodegradable plastics.

• Ingestion can be fatal to animals.
• Animals can be caught in plastics.
• Combustion of plastics can release toxins.

• What reagent can be used to oxidize an alcohol?

• Acidified potassium manganate (VII) solution.

• Describe the color change you would observe when an alcohol is

oxidized with acidified potassium manganate (VII) solution.
• Purple to colorless.

• Explain why X is an unsaturated hydrocarbon.

• Contains a (carbon-carbon) double bond.
 • Contains only C and H only.

• Describe how ethanol is formed from ethene (hydration).

• C2H4 + H2O C2H5OH
• 300 degrees Celsius.
• 6O atm.
• Phosphoric catalyst.

• Describe the difference in the structures of nylon and protein.

• Proteins are made from many monomers.
• Nylon is made from only 1 or 2 monomers.

• Name the reaction that changes amino acids into protein.

• Condensation polymerization.

• Name the reaction that changes protein into amino acid, state the
condition needed.
• Hydrolysis.
• Acid and heat.

• What is the name of the reaction that converts protein into amino
acids?
• Hydrolysis.

• What are the 2 different conditions needed to carry out hydrolysis of

polymer?
• Enzyme
• Acid and heat.

• List the reactions of:

• Alkanes:
• Complete combustion: CH4 + 2O2 CO2 + H2O
• Incomplete combustion: CH4 + 2O2 CO + H2O
• Substitution: CH4 + CL2 CH3CL + HCL (in the presence of UV light
rays)
• Cracking: C2H6 C2H4 + H2 (at high temperature)

• Alkenes:
• Halogenation: C2H4 + BR2 C2H4BR2.
• Hydration: C2H4 + H2O C2H5OH (alcohol is produced)
(300d degrees, 60 atm, phosphoric acid catalyst)
• Hydrogenation: C2H4 + H2 C2H6. (alkane is produced).
• Polymerization.

• Alcohols:
• Manufactured by hydration of ethene or fermentation:
C6H12O6 2C2H5OH + 2CO2 (exothermic)
 • Combustion: complete and incomplete
• Oxidation: C2H5OH CH3COOH
• Oxidizing agents are acidified potassium manganate and potassium
dichromate (VII)
• Dehydration: ethanol dehydrated to produce ethene
C2H5OH C2H4 + H2O (catalyst: concentrated sulphuric acid)
• Esterification: alcohols and carboxylic acids react to form ester.
CH3COOH + C2H5OH CH3COOC2H5 + H2O

• Carboxylic acids:
• Esterification
• Weak acid reactions.

• [Fermentation of glucose by yeast] why does the rate of reaction

increase?
• The temperature increases.
• More yeast because they multiply.

• [Fermentation of glucose by yeast] Why does the solution become

cloudy?
• More yeast.

• [Fermentation of glucose by yeast] Why does the fermentation stop?

• All glucose used up.
• Yeast die

• What is the advantage of producing ethanol via fermentation

compared to addition reaction of ethene?
• Renewable resources.

• What is the disadvantage of producing ethanol via fermentation

compared to addition reaction of ethene?
• Slower rate of reaction.