Contents
1.
1 Notes
1.
1.1 Metals in the Periodic Table
2.
1.2 Properties
1.
1.2.1 Extremes in Metals
2.
1.2.2 Metallic properties
3.
1.3 Alloys
4.
1.4 Reactivity series of metals
5.
1.5 Stability of metal compounds
6.
1.6 Displacement power of metals
1.
1.6.1 Displacement of oxides
2.
1.6.2 Displacement from solutions
3.
1.6.3 Reaction of Metal Oxides with Carbon
4.
1.6.4 Reaction of Metal Oxides with Hydrogen
5.
1.6.5 Decomposition of Metal Carbonates
7.
1.7 Extraction of Metals
1.
1.7.1 Uses of Metals
2.
1.7.2 Recycling Metals
8.
1.8 Iron
9.
1.9 Steel
10. 1.10 Rusting
2.
2 MCQ Questions
1.
2.1 Answers
2.
2.2 Structured Questions and Worked Solutions
Notes
Metals in the Periodic Table


Mainly in Group I, Group II, and the Transition Block
those near the staircase line
Properties
1. High density, melting point and boiling point




due to close packing of the atoms in metals
strong forces between these atoms
high density except sodium
high melting and boiling points except mercury and sodium
2. Malleable and ductile



when a force is applied to a metal, the atoms can slide over one another
malleable: can be bent and beaten into different shapes
ductile: can be stretched to form wires
3. Thermal conductivity


heat energy can be transferred from one atom to another by vibration as the atoms are very close
together
the outermost electrons also help to conduct heat
4. Electrical conductivity


when a metal is connected to a circuit, the free outermost electrons move towards the positive
terminal
to replace them, more electrons are fed into the metal from the negative terminal
Metals always from positive ions
Extremes in Metals
Lightest: Lithium
Heaviest: Osmium
Most brittle: Manganese and chromium
Lowest melting point: Mercury
Highest melting point: Tungsten
Most expensive: Platinum
Rarest: Rhodium
Most abundant: Aluminium
Metallic properties
Metals have high:
- density: high mass per unit volume
- tensile strength: high strength of the metal under stress
- durability: resistant to corrosion
- malleability: ability to be made into sheets
- ductility: ability to be made into wires
- thermal conductivity: ability to conduct heat
- electrical conductivity: ability to conduct electricity
- sonority: ability to produce sound when struck
Alloys





a mixture of metallic elements or metallic with non-metallic.
Pure metals are weak as the layers of atoms slide over each other easily. in alloy of 2 metals,
they have different sizes of atoms so this distrupts the orderly layer of atoms making it difficult for
atoms to slide over.
Eg of alloys
 Steel: iron and carbon
 bronze: copper and tin
 brass: copper and zinc
 duralumin: aluminium, copper, magnesium
 Uses of duralumin: it is light but strong and durable so used for aircraft parts,
greenhouse frames, overhead cables, curtain walling in high-rise buildings etc.
 pewter: tin and lead
Uses of solder: mixture of tin and lead, has a much lower melting point than either of its
components so more easily fusible --- suitable for welding electrical wire together
Uses of stainless steel: is an alloy of iron containing chromium or nickel. Is the most expensive
way
 applications for:
 cutleries
 medical instruments
 kitchen sinks
 steel objects in chemical factories and oil refineries
Reactivity series of metals





Reactive metals tend to form positive ions easily, by losing electrons and forming compounds
unreactive metals prefer to remain in uncombined form, as the element itself
the order of reactivity is worked out from the metal's reaction (if any) with water or steam and
acids
if there is a reaction, the metal displaces hydrogen
Metal + hydrogen ion ---> metal ion + hydrogen gas
Metal
potassium
sodium
calcium
Metal with water/steam
react with cold water
M(s) + 2H2O(l) --> MOH(aq) + H2(g)
Metal + Water --> Metal Hydroxide +
Hydrogen
Metal with acid
violent reaction with dilute acids
M(s) + 2HCl(aq) --> MCl2(aq) + H2(g)
Metal + Acid --> Metal Chloride + Hydrogen
magnesium
aluminium
zinc
iron
react with steam
M(s) + 2H2O(g) --> MO(s) + H2(g)
Metal + Water --> Metal Oxide +
Hydrogen
react with dilute acids with decreasing ease
M(s) + 2HCl(aq) --> MCl2(aq) + H2(g)
Metal + Acid --> Metal Chloride + Hydrogen
lead
hydrogen
do not react with water or steam
react with dilute acids with decreasing ease
copper
mercury
silver
platinum
do not react with water or steam
react only with concentrated acids







In the reactivity series, metals at the top, like potassium and sodium, react violently with cold
water. Hence, they are stored under the surface of oil to prevent water vapour in the atmosphere
from reacting with them
 eg. 2Na + 2H2O ---> 2NaOH + H2
Down the series, the reactivity of the metal decreases.
Magnesium will react only with steam, and for metals below iron there is no reaction with either
cold water or steam.
 eg. Mg + H2O ---> MgO + H2
With dilute hydrochloric acid, the metals at the top of the series react very violently. As we go
down, the metals react less vigorously.
Aluminium, although above iron and zinc, reacts more slowly because of a protective oxide coat
on its surface.
 eg. Fe + 2HCl ---> FeCl2 + H2
Below lead, there is no reaction with steam or with dilute acids and so hydrogen is never
displaced. Hence its position in the series.
The metals below hydrogen will react only with concentrated acids which are capable of oxidising
the metal first to its oxide. Such acids are concentrated nitric or sulphuric acids
 eg. Cu + 4HNO3 ---> Cu(NO3)2 + 2NO2 +2H2O
Stability of metal compounds

Compounds of metals high up in the reactivity series are stable and not easily decomposed by
heating.




Compounds of metals low down in the series are unstable, and are often decomposed by
heating, or are easily reduced.
The oxides of metals above zinc in the series can only be reduced to the metal by using
electrolysis.
At cathode, reduction occurs
+
 Al3 + 3e ---> Al
The oxides below can be reduced with reducing agents like carbon or hydrogen, except zinc
oxide which cannot be reduced by action of hydrogen
 ZnO + C --> Zn + CO
 CuO + H2 --> Cu + H2O
Metal
potassium
sodium
Oxide
electrolytic
reduction
calcium
magnesium
aluminium
electrolytic
reduction
zinc
iron
lead
copper
reduced by
heating with
carbon
mercury
silver
platinum
reduced by
heating alone






Hydroxide
stable to heat
Carbonate
stable to heat
Nitrate
decompose to
nitrite and
oxygen
decompose to decompose to
decompose to
metal oxide
metal oxide and metal oxide,
and steam on carbon dioxide nitrogen dioxide
heating
gas on heating and oxygen on
heating
decompose
decompose to
decompose to
to metal oxide metal oxide and metal oxide,
and steam on carbon dioxide nitrogen dioxide
heating
gas on heating and oxygen on
heating
unstable, do
unstable, do
decompose to
not exist
not exist
metal, oxygen
and nitrogen
dioxide gas on
heating
Down the series, reduction becomes easier because the metals prefer to exist as atoms, as
opposed to ions
For metal oxides like mercury(II) oxide, no reducing agent is needed - just heating alone
 2HgO --> 2Hg + O2
Hydroxides of the metals calcium and below decompose to their corresponding oxide and give off
steam, on heating. This can be confirmed by using anhydrous copper(II) sulphate which turns
white to blue with steam
 Ca(OH)2 ---> CaO + H2O
Similarly, most carbonates, except sodium and potassium carbonates, undergo thermal
decomposition again to a metal oxide, but this time giving off carbon dioxide gas. This can be
confirmed by bubbling the gas through limewater, which turns milky with carbon dioxide
 PbCO3 ---> PbO + CO2
Nitrates also decompose on heating, but the stable ones at the top of the series only decompose
as far as the nitrite (nitrite(III)), giving off oxygen gas. This can be identified by the gas relighting a
glowing splinter
 2KNO3 ---> 2KNO2 + O2
The majority of nitrates decompose to the metal oxide, giving off brown fumes of nitrogen dioxide
as well as oxygen gas.
 2Mg(NO3)2 ---> 2MgO + 4NO2 + O2

The unstable nitrates at the bottom of the reactivity series decompose all the way to the metal
itself
 2AgNO3 ---> 2Ag + 2NO2 + O2
Displacement power of metals
Displacement reaction is the displacement of ions of metal from compounds of metals lower in reactivity
series by metals higher in reactivity series.
E.g. Magnesium displaces copper(II) chloride
Mg(s) + CuCl2(aq) -> MgCl2(aq) + Cu(s)


For observation, we’ll see silver magnesium metal coated with brown copper metal
Displacement is due to Mg atoms transfer electrons to Cu2+ ions forming Cu atoms.
2+
-
Mg(s) → Mg (aq) + 2e
2+
Cu (aq) + 2e → Cu(s)


Loss of electrons is due to it’s less reactive as less reactive metal has higher chance of losing
electrons.
That’s why when Mg is placed in KCl, no reaction occurs.
Mg(s) + KCl2(aq) --> No reaction
Displacement of oxides





The Thermit Reaction is an example of displacement of oxides.
Iron(III) oxide and aluminium powder are heated in a crucible, with a magnesium fuse to start the
reaction.
The aluminium is more reactive, and takes the oxygen from the iron oxide, leaving molten iron at
the bottom of the crucible.
 Fe2O3 + 2Al ---> Al2O3 + 2Fe
This reaction is called the Thermit Reaction as it produces large quantities of heat.
It has been used to weld railway lines in remote areas where normal welding techniques are not
possible.
Displacement from solutions




In general, the more reactive metal goes into solution displacing the less reactive.
For eg, if iron filings were slowly added, with stirring, to a blue solution of copper(II) sulphate, the
blue color would fade and become faintly greenish.
This is because the copper has been pushed out, and is left as pink copper metal, while the iron
has gone into solution as green iron(II) sulphate
 CuSO4 + Fe ---> FeSO4 + Cu
In displacement reactions, the powder form is used as powders have a greater surface area and
so will react more quickly.

Using ionic equations to show displacements:
2+
2+
 Cu + Fe ---> Fe + Cu
Reaction of Metal Oxides with Carbon




The lower the position of metal in reactivity series, the easier for carbon to remove oxygen from
metal oxide by heating. At higher position, stronger heat is needed.
E.g. CuO reacts with C can be reduced by bunsen burner flame temperature
CuO(s) + C(s) --> Cu(s) + CO2(g)
For iron oxide to be reduced, it needs very high temperature.
Reaction of Metal Oxides with Hydrogen



The lower position of metal in reactivity series, the easier hydrogen remove oxygen from metal
oxide by heating. At higher position, stronger heat is needed.
E.g. PbO reacts with H2 can be reduced by bunsen burner flame temperature
PbO(s) + H2(g) --> Pb(s) + H2O(l)
Decomposition of Metal Carbonates



The lower position of metal in reactivity series, the easier hydrogen remove oxygen from metal
oxide by heating. At higher position, stronger heat is needed.
E.g. CuCO3 reacts decomposes by heat of bunsen burner flame temperature
CuCO3(s) --> Cu(s) + CO2(g)
Extraction of Metals
Metals from Rocks


Minerals – elements/compounds that make up rocks
Metal ore – rock containing metal
Extracting these metals



Metal ores are removed from ground.
The ores contain useful and unwanted materials. Unwanted materials are separated to obtain
concentrated mineral.
Metal is extracted from the mineral.
Occurrence of Metals

Metal ores are compounds, usually as:
 Metal oxides – metal + oxygen, eg: Al2O3
 Metal sulphides – metal + sulphur, eg: HgS
 Metal carbonates – metal + carbon + oxygen, eg: MgCO3



Least Reactive – easiest to extract; extracted by physical methods
Less Rective – harder to extract than least reactive; by blast furnace; usually occur as
compounds of oxides or sulphides.
Most Reactive – hardest to extract – strong bonds in compounds; by electrolysis – decomposing
compounds with electricity.
Uses of Metals
The choice of metals over another depends on 3 factors:



Physical properties (e.g. melting point, strength, density, conductivity)
Chemical properties (e.g. resists corrosion)
Cost
Recycling Metals



There are many iron on the surface but copper and tin are seriously reducing.
High temperatures and pressures and greater depth increases hazards that prevent mining up to
the lower part of crust, although there are more metals further down
Ways to conserve metals
 Use alternative materials to replace the use of iron (e.g. use of plastic pipes instead of
iron, use of glass bottles for soft drinks instead of aluminium)
 Recycle unused metals by melting them to produce new blocks of clean metal
Advantages of recycling metals







Recyling helps conserving metals, especially valuables such as gold and platinum.
 E.g. used computer parts processed to extract gold used for electrical contacts of
processors and memory chips
Recycling saves the cost of extracting new metals
Recycling benefits environment, e.g. if there is a car wasteland, it causes eyesore
Recycling metals can damage the environment by smelting process which sends a lot of fumes
into the air
Cost to separate metals from waste is high. E.g. separat metals in alloys is hard
Transport costs for collecting scrap metal is high, e.g. trucks should be used
People are not interested in depositing their used materials in recycling bins
Iron


Steel
Iron is extracted from the iron ore haematite, Fe2O3
Iron is extracted from the oxide in a blast furnace


Iron made from blast furnace is not good as:
 it contains impurities which makes it brittle (can break easily)
 it cannot be bent or stretched
Most iron is converted into steel which is an alloy of iron and carbon with small amounts of other
elements. Advantages of steel:
 it is strong and tough
 it can be bent and stretched without shattering
Making Steel:


Impurities of iron is removed by blowing oxygen into molten iron to change the impurities into
oxides. They are then combined with CaO and removed as slag.
Carbon and other metals are added in certain amount to make steel.
Different Types of Steel:



Mild steel – is a low carbon steel with 0.25% carbon
 It is strong and quite malleable. It is used for car bodies, ships, railway lines and steel
rods to reinforce concrete
Hard steel – is a high-carbon steel with about 1% carbon
 It is harder than mild steel and less malleable. It is used to make tools
Stainless steel – is iron with large amounts of chromium and nickel
 It is hard, shiny and doesn’t rust. It is used to make cutleries, medical instrument and
pipes in chemical industries.
Rusting



Rusting – corrosion of iron and steel
Rust – brown solid product formed during rusting
Rust is hydrated iron(III) oxide Fe2O3.xH2O where water molecules varies.
Conditions for rusting
Tubes A B C



After a few days, only nail in tube A rust.
This shows that air and water is needed for rust.
In boiled water, the nail doesn’t rust in B as boiled water removes dissolved air while in C, CaCl
keeps air dry so there’s no water.
Preventing Rusting



Surface protection: covers metal with a layer of substance
 Paint
 Grease or oil (also help to lubricate)
 Plastic
 Metal Plating – covering metal with thin layer of another metal (e.g. tin,
chromium, silver)
 Advantage – These methods are cheap (except metal plating)
 Disadvantage – If the layer is broken, air and water an reach metal to rust
Sacrificial protection
 to sacrifice more reactive metal to corrode with water and air by layering it over less
reactive metal (e.g. iron covered by magnesium).
 If layer is broken, water & air reach underneath layer, overlying metal still protect it.
 Applications:
 Galvanised Iron – is steel coated with zinc, usually used on roofs.
 Protecting ships – blocks of zinc are attached to hulls to corrode instead of steel
which is the ship metal.
 Underground steel pipes – these are attached to magnesium block using
insulated copper cables. Magnesium corrodes first than steel.
Use of stainless steel
MCQ Questions
1. Caesium is a metal that is more reactive than aluminium. Which reaction would produce
caesium?
a. electrolysing aqueous caesium chloride
b. electrolysing molten caesium chloride
c. heating caesium carbonate
d. heating caesium oxide with carbon
2. Which of the following processes does not result in the formation of both carbon dioxide and
water?
a. addition of a dilute acid to a carbonate
b. burning ethanol
c. burning methane
d. heating crystals of hydrated sodium carbonate
3. Which element is always present with iron in mild steel?
a. aluminium
b. carbon
c. chromium
d. nickel
4. Hydrogen gas will reduce
a. calcium oxide
b. silver oxide
c. magnesium oxide
d. potassium oxide
5. Which oxide can be reduced to the metal using carbon
a. calcium oxide
b. magnesium oxide
c. sodium oxide
d. zinc oxide
6. Which substance removes impurities from iron ore in the blast furnace?
a. carbon
b. limestone
c. sand
d. slag
2+
2+
7. An excess of iron filings is added to a solution containing a mixture of the ions Mg , Ca ,
2+
+
Cu and Ag . Which 2 metals will be displaced from this solution?
a. calcium and copper
b. calcium and magnesium
c. copper and silver
d. magnesium and silver
8. What reacts with hydrochloric acid to give hydrogen?
a. ammonia
b. iron
c. silver
d. sodium hydroxide
9. Why does the color of aqueous potassium bromide change when chlorine gas is bubbled into
it?
a. a compound is formed between chlorine and bromine
b. a solution of potassium chloride is formed
c. the chlorine oxidises bromide ions to bromine
d. the potassium bromide is reduced
10. Which carbonate decomposes on heating to give a black solid and a colourless gas?
a. calcium carbonate
b. copper(II) carbonate
c. sodium carbonate
d. zinc carbonate
11. Which substance is not an essential raw material in the extraction of iron in a blast furnace?
a. air
b. coke
c. limestone
d. sand
12. Which element reacts with oxygen to form a compound that is a gas at room temperature?
a. magnesium
b. hydrogen
c. copper
d. carbon
13. Caesium is a metal that is more reactive than aluminium. Which reaction would produce
caesium?
a. electrolysing aqueous caesium chloride
b. electrolysing molten caesium chloride
c. heating caesium carbonate
d. heating caesium oxide with carbon
14. What is a disadvantage of recycling metals?
a. collection and transportation costs money
b. metal ores are a finite resource
c. most metals corrode slowly in the environment
d. scrap metal melts when heated
15. Which of the following processes does not result in the formation of both carbon dioxide and
water?
a. addition of a dilute acid to a carbonate
b. burning ethanol
c. burning methane
d. heating crystals of hydrated sodium carbonate
16. Hydrogen gas will reduce
a. calcium oxide
b. silver oxide
c. magnesium oxide
d. potassium oxide
17. Which element is always present with iron in mild steel?
a. aluminium
b carbon
c. chromium
d. nickel
18. Which oxide can be reduced to the metal using carbon?
a. calcium oxide
b. magnesium oxide
c. sodium oxide
d. zinc oxide
19. Which substance removes impurities from iron ore in the blast furnace?
a. carbon
b. limestone
c. sand
d. slag
20. What reacts with hydrochloric acid to give hydrogen?
a. ammonia
b. iron
c. silver
d. sodium hydroxide
21. Which carbonate decomposes on heating to give a black solid and a colorless gas?
a. calcium carbonate
b. copper(II) carbonate
c. sodium carbonate
d. zinc carbonate
22. Which substance is not an essential raw material in the extraction of iron in a blast furnace?
a. air
b. coke
c. limestone
d. sand
23. Which element reacts with oxygen to form a compound that is a gas at room temperature?
a. magnesium
b. hydrogen
c. copper
d. carbon
24. A sample of air is slowly passed through aqueous sodium hydroxide and then over heated
copper. Which gases are removed by this process?
a. carbon dioxide and water vapour
b. carbon dioxide and oxygen
c. nitrogen and oxygen
d. nitrogen and water vapour
25. When heated, solid X gives off a gas which turns limewater milky. The residue reacts with
dilute acid and also with aqueous alkali. What is X?
a. copper(II) carbonate
b. magnesium carbonate
c. sodium carbonate
d. zinc carbonate
26. An element is burned in an excess of oxygen. Which statement about the oxide formed is
always correct?
a. it is a crystalline solid
b. it is greater in mass than the element
c. it is soluble in water
d. it is white in color
27. Which substance can be reduced by carbon?
a. aluminium oxide
b. calcium carbonate
c. iron(III) oxide
d. magnesium oxide
28. Which of the following is a typical property of transition metals?
a. they form colored compounds
b. they have low densities
c. they have low melting points
d. they react with cold water to give hydrogen
29. What happens when zinc is placed in aqueous copper(II) sulphate?
a. copper atoms are oxidised
b. zinc atoms are oxidised
c. copper ions are oxidised
d. zinc ions are oxidised
30. Which substance does not need air as a raw material for its manufacture?
a. ammonia
b. iron
c. sodium
d. sulphuric acid
31. Which of the following is not a use of silicon or its compounds?
a. making fire-resistant plastics
b. making glass
c. making polishes
d. making smokeless fuel
32. Compound X reacts with some metals to liberate hydrogen and is used to make fertilisers. It
gives a white precipitate when added to aqueous barium nitrate. What is X?
a. ammonium sulphate
b. hydrochloric acid
c. potassium nitrate
d. sulphuric acid
33. Which industrial process uses iron as a catalyst?
a. making ammonia from nitrogen and hydrogen
b. making ethanol from ethene and steam
c. making steel
d. making sulphur trioxide from sulphur dioxide and oxygen
34. Which pair of elements will combine to form an ionic compound?
a. carbon and chlorine
b. fluorine and sodium
c. hydrogen and oxygen
d. oxygen and carbon
35. How does the mass of a sample of copper(II) oxide change when it is heated in hydrogen and
in oxygen?
a
b
c
d
mass after heating in
hydrogen
decreases
decreases
unchanged
unchanged
mass after heating in
oxygen
decreases
unchanged
decreases
unchanged
36. Sodium is a metal. Using only this information, what can be deduced about sodium?
a. it has a low melting point
b. it is a conductor of electricity
c. it is less dense than water
d. it is very reactive
37. Which substance reacts with water to form a soluble compound and an insoluble gas?
a. ammonium sulfate
b. caesium
c. calcium carbonate
copper
38. Which compound does not give off a gas when heated?
a. hydrated copper(II) sulfate
b. hydrate sodium carbonate
c. magnesium carbonate
d. sodium carbonate
39. Which metal should be used in the sacrificial protection of the hull of a boat made from iron?
a. calcium
b. copper
c. lead
d. zinc
40. A coil of clean copper wire is suspended in a beaker of aqueous silver nitrate. Crystals of
silver are deposited on the copper wire. Which statement is not correct?
a. the copper is oxidised
b. the solution turns blue
c. the total mass of the crystals of silver increases gradually
d. the total number of positive ions in the solution is unchanged
41. In the manufacture of iron by the blast furnace, which are the main gases that escape from the
top of the blast furnace?
a. carbon monoxide, carbon dioxide, hydrogen
b. nitrogen, carbon dioxide, carbon monoxide
c. nitrogen, oxygen, steam
d. oxygen, carbon dioxide, sulfur dioxide
42. When heated, solid X gives off a gas. When this gas is bubbled through limewater, a white
precipitate is formed. The residue after heating solid X reacts with dilute acid and also with
aqueous alkali. What is X?
a. copper(II) carbonate
b. magnesium carbonate
c. sodium carbonate
d. zinc carbonate
43. The information below concerns 3 elements X, Y, and Z.
X: Its oxide is decomposed by heat to the element.
Y: Its carbonate is not decomposed by heat.
Z: Its oxide is not decomposed by heat but its carbonate decomposes.
In order of decreasing reactivity, the 3 elements should be arranged as:
a. Y Z X
b. X Y Z
c. Y X Z
d. X Z Y
Answers
1. b
2. d
3. b
4. b
5. d
6. b
7. c
8. b
9. c
10. b
11. d
12. d
13. b
14. a
15. d
16. b
17. b
18. d
19. b
20. b
21. b
22. d
23. d
24. b
25. d
26. b
27. c
28. a
29. b
30. c (sodium is obtained through electrolysis)
31. d
32. d
33. a
34. b
35. b
36. b
37. b
38. c
39. d
40. b
41. b
42. d
43. a
Structured Questions and Worked Solutions
1a. Under what conditions does water react with
i. sodium
ii. magnesium
In each case, name the products formed.
b. Water supplies are obtained from rivers, boreholes and reservoirs. The water must be treated
before use. Describe and explain the two main processes in the purification of water supplies.
c. Water supplies that have passed through iron pipes contain iron(II) ions, Fe
3+
Fe .
2+
and iron(III) ions,
In the presence of air, iron(II) ions are slowly changes to iron(III) ions.
+
Construct the equation for the reaction between iron(II) ions, hydrogen ions, H , and oxygen to
form iron(III) ions and water.
Solution
1ai. In cold water.
Products: sodium hydroxide and hydrogen
1aii. Heated with steam.
Products: magnesium oxide and hydrogen
1bi. filtration: solid particles are removed.
1bii. chlorination: germs and bacteria are killed by sterilising water with chlorine.
1c. 4Fe
2+
+
3+
(aq) + O2 (g) + 4H (aq) ---> 4Fe (aq) + 2H2O (l)
2. Calcium oxide is produced by heating a mixture of limestone and coke in a lime kiln.
CaCO3 <---> CaO + CO2
ai. Explain the meaning of the symbol <--->
aii. In the lime kiln, the carbon dioxide is allowed to escape. Why does this increase the yield of
calcium oxide?
b. The calcium oxide reacts with water to form slaked lime.
i. Give the equation for this reaction
ii. State a use of slaked lime
Solution
2ai. It shows that the decomposition of calcium oxide in the lime kiln is a reversible reaction.
2aii. The decrease in carbon dioxide concentration causes the equilibrium to shift to the right to produce
more carbon dioxide to replace those that escaped. Therefore, more calcium carbonate decomposes to
give calcium oxide.
2bi. CaO + H2O --> Ca(OH)2
2bii. It is used to treat acidic soils. It reacts with acid to produce salt and water.
3. Choose from the following metals to answer the questions below.
aluminium
calcium
copper iron
magnesium
potassium sodium zinc
Each metal can be used once, more than once, or not at all.
Name a metal which
a. is manufactured by the electrolysis of its molten oxide
b. has a variable valency
c. is used to galvanise iron
d. has a carbonate which is coloured
e. is alloyed with zinc to make brass
Solution
3a. aluminim
3b. copper/iron
3c. zinc
3d. copper
3e. copper
4. In separate experiments, powdered samples of metal X and metal Y reacted with solutions of
nickel(II) sulphate and of iron(II) sulphate. The following table shows how the colours of the
solutions changed.
metal X
metal Y
nickel(II) sulphate
Solution goes from
green to colourless
Solution goes from
green to colourless
iron(II) sukphate
Solution stays pale
green
Solution goes from
pale green to
colourless
a. predict the order of reactivity for the four metals X, Y, nickel, and iron.
b. Metal Y was placed in aqueous copper(II) sulphate.
i. What colour change was seen?
ii. Give one other observation
c. Write the ionic equation, with state symbols, for the reaction between iron and aqueous
nickel(II) sulphate.
Solution
a. from least reactive to most reactive: nickel, metal X, iron, metal Y
bi. blue copper(II) sulphate solution decolourises
bii. a reddish brown deposit is formed
2+
c. Fe (s) + Ni
(aq) ----> Fe
2+
(aq) + Ni (s