2. Acids
An acid is a compound which when
dissolved in water produces
hydrogen ions as the only positively
charged ions e.g. sulphuric acid
(H2SO4), hydrochloric acid(HCl),
nitric acid(HNO3), phosphoric acid
(H3PO4) and carbonic acid (H2CO3).
3. Basicity of an acid
Basicity of an acid is the number
of hydrogen ions produced by one
molecule of an acid when dissolved
in water.
• Monobasic acids: acids that
produces only one hydrogen ions
when dissolved in water e.g.
nitric acid and hydrochloric acid.
4. • Dibasic acids: acids that
produces two hydrogen ions when
dissolved in water e.g. sulphuric
acid, carbonic acid, oxalic acid.
• Tribasic acids: acids that
produces three hydrogen ions
when dissolved in water e.g.
phosphoric acid.
5. Types of acids
• Strong acids: are acids which
ionize completely when dissolved
in water e.g. hydrochloric acid,
nitric acid and sulphuric acid.
HCl(aq) H+(aq) + Cl-(aq)
HNO3(aq) H+(aq) + NO3
-(aq)
H2SO4(aq) 2H+(aq) + SO4
2-(aq)
6. • Weak acids: are acids which
ionize partially when dissolved in
water e.g. carbonic acid,
ethanoic acid, sulphurous acid.
7. Remember: Solution containing a
small amount of acid dissolved in
water is dilute whereas solution
containing large amount of acid
dissolved in water is concentrated.
8. Properties of acids
• Acids have a sour taste.
• Acids are soluble in water.
• Acids are corrosive.
• Acids turn blue litmus paper red/
pink.
• Acids have pH less than 7.
9. • Acids react with highly
electropositive metals to form
salts and hydrogen gas.
Mg(s) + H2SO4(aq) MgSO4(aq) + H2(g)
Zn(s) + HNO3(aq) Zn(NO3)2(aq) + H2(g)
OR
Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)
10. • Acids react with oxides and
hydroxides to form salts and
water only.
CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l)
NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)
11. • Acids react with carbonates and
hydrogencarbonates to form
salts, water and carbon dioxide.
CuCO3(s) + 2HNO3(aq)
Cu(NO3)2(aq) + H2O(l) + CO2(g)
Observation: Bubbles of a colourless
gas observed.
12. • Acids react with sulphites and
hydrogensulphite to form salts,
water and sulphur dioxide.
2H+(aq) + SO3
2-(aq) H2O(l) + SO2(g)
H+(aq) + HSO3
-(aq) H2O(l) + SO2(g)
Observation: Bubbles of a colourless
gas observed.
13. Bases and alkalis
A base is an oxide or hydroxide of
a metal that reacts with an acid to
form salt and water only.
An alkali is a compound which when
dissolved in water forms hydroxide
ions, OH- as the only negatively
charged ions e.g. sodium oxide,
potassium oxide, sodium hydroxide,
14. Potassium hydroxide, ammonium
hydroxide and calcium hydroxide.
Alkalis are soluble bases.
There are two types of alkalis:
• Strong alkalis: are alkalis which
completely ionized in solution e.g.
sodium hydroxide and potassium
hydroxide.
15. • Weak alkalis: are alkalis which
partially ionizes in solution e.g.
ammonium hydroxide.
Remember: Bases are proton
acceptors whereas acids are proton
donators.
16. Properties of alkalis
• Alkalis have bitter taste.
• They turn red litmus paper blue.
• Alkalis reacts with acids
(neutralization reaction) to form
salts and water only.
NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)
• Alkalis liberates ammonia gas
when heated with ammonium salts
17. NaOH(aq) + NH4Cl(s)
NaCl(aq) + NH3(g) + H2O(l)
• Alkalis react with Ca2+, Mg2+,
Zn2+, Al3+, Pb2+ ions to form
white precipitate, Cu2+ forms blue
precipitate, Fe2+ forms dirty
green precipitate and Fe3+ forms
brown precipitate.
19. Remember: Zn(OH)2 is soluble in both
excess ammonia solution and sodium
hydroxide solution whereas Al(OH)3
and Pb(OH)2 are soluble in excess
sodium hydroxide solution only. A
colourless solution is formed.
Cu(OH)2 is soluble in excess ammonia
forming a deep blue solution.
21. • Sodium hydroxide and potassium
hydroxide are deliquescent. Both
absorbs water from the
atmosphere to form solutions and
thus absorbs carbon dioxide from
the atmosphere to form sodium
carbonate and potassium
carbonate (both are white
crystalline solids).
23. pH of solutions
The pH of a solution is a measure of
the number of hydrogen ions in that
solution.
The lower the pH number, the more
hydrogen ions there are, and thus the
more acidic the solution.
Strong acids have low pH values
whereas strong bases have high pH
values and neutral solutions show
pH = 7.
24. Indicators
Indicators are substances which
changes colour according to the pH
of the medium in which it is added.
Common indicators are:
Indicator Colour in acid Colour in alkaline
Phenolphthalein Colourless Pink
Methyl orange Red Yellow
Litmus paper Red Blue
25. Salts
A salt is a compound formed when
either all or part of the ionisable
hydrogen ions of an acid is replaced
by a metallic ion or ammonium ion.
There are two types of salts:
• Normal salts: are salts formed
when all the ionisable hydrogen ions
of an acid is replaced by a metallic
ion or ammonium ion e.g. sodium
26. chloride, calcium nitrate, zinc
sulphate, ammonium chloride etc.
• Acid salts: are salts formed when
part of the ionisable hydrogen ions
of an acid is replaced by a metallic
ion or ammonium ion e.g. sodium
hydrogensulphate, magnesium
hydrogencarbonate etc.
27. Remember: Acid salts are formed
by dibasic acids like H2SO4,
H2CO3, H2SO3 or by tribasic acids
like H3PO4.
28. Nature of salts in water
• Acid salts dissolve in water to form
acidic solutions because hydrogen
ion is produced.
NaHCO3(aq) Na+(aq) + H+(aq) + CO3
2-(aq)
• Normal salts dissolve in water to
form neutral solutions. However,
ammonium chloride/ ammonium
nitrate dissolve in water to form
acidic solution. This is because
29. ammonium chloride reacts with water
to form ammonium hydroxide and
hydrochloric acid.
NH4Cl(s) + H2O(l) NH4OH(aq) + HCl(aq)
In aqueous solutions, both ammonium
hydroxide and hydrochloric acid
undergo dissociation.
30. NH4OH(aq) NH4
+(aq) + OH-(aq)
HCl(aq) H+(aq) + Cl-(aq)
Ammonium hydroxide being a weak
base undergoes partial dissociation
unlike hydrochloric acid which
shows complete dissociation since it
is a strong acid.
31. This results into a solution of
ammonium chloride having more
hydrogen ions compared to the
hydroxide ions formed by the weak
base. So the solution is overall
acidic.
32. Also sodium carbonate dissolves in
water to form an alkaline solution.
This is because sodium carbonate
reacts with water to form sodium
hydroxide and carbonic acid.
Na2CO3(s) + 2H2O(l) 2NaOH(aq) + H2CO3(aq)
In aqueous solutions, both sodium
hydroxide and carbonic acid undergo
dissociation.
33. NaOH(aq) Na+(aq) + OH-(aq)
H2CO3(aq) 2H+(aq) + CO3
2-(aq)
Carbonic acid being a weak acid
undergoes partial dissociation unlike
sodium hydroxide which shows
complete dissociation since it is a
strong base.
34. This results into a solution of
sodium carbonate having more
hydroxide ions compared to the
hydrogen ions formed by the weak
acid. So the solution is overall
alkaline.
35. Water of crystallisation
Water of crystallization is the
definite amount of water with
which some substances combine
chemically to form crystals from
their solutions.
A salt without water of
crystallization is called anhydrous
salt whereas a hydrated salt is a
36. salt that contains the water of
crystallization e.g.
sodium carbonate – 10 – water,
Na2CO3.10H2O;
copper(II) sulphate – 5 – water,
CuSO4.5H2O;
iron(II) sulphate – 7 – water,
FeSO4.7H2O
37. Exposure of salts to air
Salts behave differently when
exposed to air. Some lose their
water of crystallization to the
atmosphere and others absorb
water vapour from the atmosphere.
• A deliquescent substance absorbs
water from the atmosphere and
dissolves in it to form a solution
38. e.g. iron(III) chloride, potassium
hydroxide, sodium hydroxide, zinc
chloride, calcium chloride.
• A hygroscopic substance absorbs
water from the water e.g. all
deliquescent substances, calcium
oxide, concentrated sulphuric
acid.
40. • Efflorescent substance loses
water to the atmosphere e.g.
sodium carbonate-10-water,
copper(II) sulphate-5-water,
sodium sulphate-10-water,
iron(II) sulphate-7-water.
41. Preparation of soluble salts
Soluble salts of copper, lead, iron,
magnesium, aluminium and zinc are
prepared by the general steps:
Add the metal(Mg, Zn, Fe), metal
oxide, metal carbonate or metal
hydroxide to an appropriate acid
until when the solid is in excess.
Heat if necessary.
42. Filter off the excess solid.
Saturate the filtrate by heating to
evaporate off most of the water.
Cool the saturated solution until
when the crystals form.
Filter off the crystals.
Wash the crystals with distilled
43. water to remove any soluble
impurities.
Dry the crystals using filter
papers.
44. Example (UNEB 2003/P2/14(a))
Describe how a dry sample of
copper(II) sulphate may be
prepared from copper(II) oxide.
Possible Answer
To warm dilute sulphuric acid in a
beaker, copper(II) oxide is added
little at a time while stirring the
45. mixture until in excess.
Copper(II) sulphate is produced
according to the equation:
CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l)
The excess copper(II) oxide is
filtered off and the filtrate is
saturated by heating.
46. The saturated solution is left to
cool until when the crystals form.
The crystals are filtered off,
washed with distilled water to
remove any soluble impurities and
dried by pressing between filter
papers.
47. Preparation of insoluble salts
Insoluble salts are prepared by
precipitation/ double decomposition.
Precipitation involves reacting two
soluble salts to form soluble salt
and an insoluble salt.
The insoluble salt is obtained as a
residue after filtering.
The residue is washed with distilled
48. water to remove any soluble
impurities and dried by pressing
between filter papers.
Example (UNEB 2006/P2/12(c))
Briefly describe how a pure sample
of calcium carbonate can be
prepared.(Diagrams not required)
49. Possible Answer
To a solution of calcium chloride in
a beaker, aqueous sodium
carbonate is added a little at a
time while stirring the mixture until
precipitation stops.
Calcium carbonate is formed
according to the equation:
50. Ca2+(aq) + CO3
2-(aq) CaCO3(s)
Filter out the calcium carbonate
from the mixture.
Wash the residue with distilled
water and dry it using filter paper.
51. Preparing salts by neutralization
The salts of potassium, sodium and
ammonium are prepared by
neutralization.
Neutralization is where an acid is
reacted with an alkali to form a
salt and water only.
52. Example 1(UNEB 1992/P2/13(c))
Describe how a dry sample of
sodium chloride can be prepared in
the laboratory.
Possible Answer
Titrate a known volume(V1cm3) of a
standard sodium hydroxide solution
with a standard hydrochloric acid
53. using phenolphthalein indicator.
Sodium chloride is formed according
to the equation:
NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)
Note the volume(V2cm3) of the acid
at the end point.
To a fresh volume(V1cm3) of the
sodium hydroxide, add a volume of
hydrochloric acid equal to V2cm3 and
stir.
54. Heat the mixture to saturation
point.
Allow the saturated solution to cool
in order for the crystals to form.
Filter, wash with distilled water
and dry them between filter
papers.
55. Example 2(UNEB 2008/P2/13(b))
Outline how a pure dry sample of
sodium hydrogensulphate can be
prepared in the laboratory (No
equation or diagram required).
Possible Answer
Titrate a known volume(V1cm3) of a
standard sodium hydroxide solution
with a standard sulphuric acid
56. using phenolphthalein indicator.
Note the volume(V2cm3) of the acid
at the end point.
To a fresh volume(V1cm3) of the
sodium hydroxide, add a volume of
sulphuric acid equal to V2cm3 and
stir.
57. Heat the mixture to saturation
point.
Allow the saturated solution to cool
in order for the crystals to form.
Filter, wash with distilled water
and dry them between filter
papers.
58. Direct synthesis
Direct synthesis can be used to
prepare both soluble and insoluble
salts e.g.
• Iron(III) chloride is formed when
dry chlorine is passed over
heated iron metal.
2Fe(s) + 3Cl2(g) 2FeCl3(s)
60. • Zinc sulphide, copper(II) sulphide
and iron(II) sulphide are formed
when a mixture of the metal and
sulphur is heated.
Zn(s) + S(s) ZnS(s)
Cu(s) + S(s) CuS(s)
Fe(s) + S(s) FeS(s)
61. Effect of heat on carbonates
• Ammonium carbonate decomposes
on heating to form ammonia gas,
carbon dioxide and water.
(NH4)2CO3(s) NH3(g) + CO2(g) + H2O(l)
• Sodium carbonate and potassium
carbonate do not decompose on
heating.
62. • Magnesium carbonate and calcium
carbonate decomposes to form
respective metal oxide and
carbon dioxide.
MgCO3(s) MgO(s) + CO2(g)
CaCO3(s) CaO(s) + CO2(g)
63. Observation: White solid forms white
residue.
• Zinc carbonate decomposes to form
zinc oxide and carbon dioxide.
ZnCO3(s) ZnO(s) + CO2(g)
Observation: White powder turns to
yellow residue when hot and white on
cooling.
64. • Lead(II) carbonate decomposes
to form lead(II) oxide and
carbon dioxide.
PbCO3(s) PbO(s) + CO2(g)
Observation: White solid turned to
reddish-brown residue when hot
and yellow on cooling.
65. • Copper(II) carbonate and iron(II)
carbonate decomposes to form
respective metal oxide and
carbon dioxide.
CuCO3(s) CuO(s) + CO2(g)
FeCO3(s) FeO(s) + CO2(g)
Observation: Green solid turned
into black residue.
66. Effect of heat on nitrates
• Ammonium nitrate decomposes to
form dinitrogen oxide and water.
NH4NO3(s) N2O(g) + 2H2O(l)
• Sodium nitrate and potassium
nitrate decomposes to form
respective metal nitrite and
oxygen gas.
67. 2NaNO3(s) 2NaNO2(s) + O2(g)
2KNO3(s) 2KNO2(s) + O2(g)
• Calcium nitrate, magnesium
nitrate and aluminium nitrate
decomposes to form respective
metal oxide, nitrogen dioxide gas
and oxygen gas.
68. 2Ca(NO3)2(s) 2CaO(s) + 4NO2(g) + O2(g)
2Mg(NO3)2(s) 2MgO(s) + 4NO2(g) + O2(g)
4Al(NO3)2(s) 2Al2O3(s) + 12NO2(g) + 3O2(g)
Observation: White solid forms
white residue and brown gas given
off.
• Lead(II) nitrate decomposes with
cracking sound to form lead(II)
oxide, nitrogen dioxide
69. and oxygen gas.
2Pb(NO3)2(s) 2PbO(s) + 4NO2(g) + O2(g)
Observation: White solid turned to
reddish-brown residue when hot
and yellow on cooling and brown gas
given off.
• Zinc nitrate decomposes on
heating to form zinc oxide,
70. nitrogen dioxide and oxygen gas.
2Zn(NO3)2(s) 2ZnO(s) + 4NO2(g) + O2(g)
Observation: White solid turned to
yellow residue when hot and white
on cooling and brown gas given off.
71. • Copper(II) nitrate decomposes on
heating to form copper(II) oxide,
nitrogen dioxide and oxygen gas.
2Cu(NO3)2(s) 2CuO(s) + 4NO2(g) + O2(g)
Observation: Blue solid turned to
black residue and brown gas given
off.
72. • Iron(II) nitrate decomposes on
heating to form iron(II) oxide,
nitrogen dioxide and oxygen gas.
2Fe(NO3)2(s) 2FeO(s) + 4NO2(g) + O2(g)
Observation: Green solid turned to
black residue and brown gas given
off.
73. • Mercury(II) nitrate and silver
nitrate decomposes to form
respective metal, nitrogen dioxide
gas and oxygen gas.
2AgNO3(s) 2Ag(s) + 2NO2(g) + O2(g)
Hg(NO3)2(s) Hg(s) + 2NO2(g) + O2(g)
74. Effect of heat on sulphates
• Ammonium sulphate decomposes
to form ammonia gas and
sulphuric acid
(NH4)2SO4(s) 2NH3(g) + H2SO4(aq)
• Potassium sulphate, sodium
sulphate and magnesium sulphate
are stable and do not decompose
on heating.
75. • Copper(II) sulphate, zinc sulphate
and lead(II) sulphate decomposes
to form metal oxides and sulphur
trioxide.
CuSO4(s) CuO(s) + SO3(g)
ZnSO4(s) ZnO(s) + SO3(g)
PbSO4(s) PbO(s) + SO3(g)
76. • Iron(II) sulphate decomposes on
strong heating to form iron(III)
oxide, sulphur dioxide and sulphur
trioxide.
2FeSO4(s) Fe2O3(s) + SO3(g) + SO2(g)
Observation: White solid turned to
reddish-brown residue.
77. Remember: When hydrated iron(II)
sulphate (green solid) is gently
heated, anhydrous iron(II) sulphate
(white powder) is formed and on
strong heating, iron(III) oxide,
sulphur dioxide and sulphur trioxide
are formed whereas
78. hydrated copper(II) sulphate (blue
solid) forms anhydrous copper(II)
sulphate (white powder) on gentle
heating, then copper(II) oxide and
sulphur trioxide on strong heating.
79. Effect of heat on hydroxides
• Sodium hydroxide and potassium
hydroxide are stable and do not
decompose on heating.
• When zinc hydroxide is strongly
heated, the white solid turns
yellow when hot and white on
cooling.
Zn(OH)2(s) ZnO(s) + H2O(l)
80. • When lead(II) hydroxide is
strongly heated, the white solid
turns reddish-brown on heating
and yellow on cooling.
Pb(OH)2(s) PbO(s) + H2O(l)
81. • When iron(II) hydroxide is
strongly heated, the green solid
turns into black solid.
Fe(OH)2(s) FeO(s) + H2O(l)
• When copper(II) hydroxide is
strongly heated, the blue solid
turns into black solid.
Cu(OH)2(s) CuO(s) + H2O(l)
82. Solubility of salts
Solubility of a salt is the mass of
salt(solute) required to saturate 100g
of water(solvent) at a particular
temperature.
Solubility of most salts increase with
increase in temperature.
• Solute is a substance that dissolves
in a solvent to form a solution.
83. • A solution is a uniform mixture of
a solute and a solvent.
• A saturated solution is one that
contains the maximum amount of
a solute in a given amount of a
solvent at a particular
temperature.
84. • Supersaturated solution is one
which has dissolved more solute
than it can normally hold in the
presence of undissolved solute at
a particular temperature.
85. Determination of solubility of a salt
Prepare a saturated solution, measure
and record its temperature.
Weigh an empty dish and transfer
the saturated solution into the dish.
Reweigh the dish to determine the
mass of the dish and the saturated
solution.
Evaporate the solution to dryness and
weigh the dish again to determine the
mass of the salt and the dish.
86. Results
Mass of empty = xg
Mass of dish and saturated solution = yg
Mass of dish and salt = zg
Mass of salt = (z – x)g
Mass of water = (y – z)g
87. Treatment of results
(y – z)g of H2O dissolve (z – x)g of salt
100g of H2O dissolve
(z − x)
(y − z)
× 100g of salt
Solubility of the salt is
(z − x)
(y − z)
× 100g per
100g of water.
88. Example 1 (UNEB 1987/P2/11(b))
75g of a saturated solution
contains 30g of salt.
Calculate
(i) the solubility of the salt.
(ii) the percentage of the salt in
the saturated solution.
89. Possible Answers
(i) Mass of water = 75 – 30 = 45g
45g of H2O dissolve 30g of KNO3
100g of H2O dissolve
30×100
45
g of KNO3
66.7g of KNO3
Solubility of potassium nitrate is 66.7g
per 100g of water.
90. (ii) Percentage of KNO3
=
Mass of KNO3×100
Mass of saturated solution
=
30×100
75
= 40%
91. Example 2 (UNEB 1997/P1/6)
The solubility of salt W is 35g per
100cm3 of water at 20oC. The
mass of W in 40cm3 of water at
the same temperature is
A. 7.0g B. 14.0g
C. 87.5g D. 114.3g
92. Possible Answers
100cm3 of H2O dissolve 35g of W
40cm3 of H2O dissolve
35 × 40
100
g of W
= 14g of W
Correct option is B
93. Example 3
200g of water at 55oC are saturated
by 150g of sodium chloride.
(a) Calculate the solubility of sodium
chloride at 55oC.
(b) If the solubility of sodium
chloride at 30oC is 30g per 100g
of water. Calculate the mass of
94. mass of sodium chloride crystals
formed when the saturated solution
at 55oC is cooled to 30oC.
Possible Answers
(a) 200g of H2O dissolve 150g of NaCl
100g of H2O dissolve
150 × 100
200
g of NaCl
= 75g of NaCl
95. (b) 100g of H2O dissolve 30g of NaCl
200g of H2O dissolve
30 × 200
100
g of NaCl
= 60g of NaCl
Mass of crystals formed = 150 – 60
= 90g
96. Solubility curves
Solubility curve is a graph that
shows how the solubility of a salt
varies with temperature.
It is used to determine
• solubility of a salt at various
temperature,
• temperature at which a certain
mass of salt dissolved in water,
97. • the mass of salt obtained by
cooling the solution from higher
(T2) to lower(T1) temperature i.e.
Mass of crystals = solubility at T2 – solubility at T1
Remember: The marking points for
a graph are:
Axes, scales, shape and sharpness.
98. Example (UNEB 2006/P2/14)
(b) The solubilities of potassium chloride
and potassium nitrate at certain
temperatures are shown in the table
below.
Temperature(oC) 0 11 15 30 40 50 57
Solubility of KCl/
100g of water
27.9 31.0 32.0 36.6 40.0 43.0 45.0
Solubility of KNO3 /
100g of water
14 21.5 25 43 63 84 102
99. (i) Plot on the same axes, a graph
of solubility against temperature
for the solubilities of potassium
chloride and potassium nitrate.
(ii) State which one of the two
salts has a solubility which
increases less rapidly with
increase in temperature.
100. (iii) Determine the temperature at
which the solubilities of the
two salts are equal.
(c) A saturated solution of
potassium nitrate at 30oC was
cooled to 5oC.
Calculate the number of moles of
potassium nitrate crystal formed.
101. Possible Answers
(i)
0
20
40
60
80
100
120
0 10 20 30 40 50 60
Solubility(g/100g
of
water)
Temperature(oC)
Graphs of solubilities of potassium chloride and
potassium nitrate against temperature
Scale
Vertical axis 1cm:5g
Horizontal axis 1cm:5oC
102. (ii) Potassium chloride
(iii) 24oC (intersection)
(c) Solubility at 30oC = 43
Solubility at 5oC = 16
Mass of KNO3 crystals formed
= 43 – 16
= 27
RFM of KNO3 = 39 + 14 + 48 = 101
Moles of KNO3 formed =
27
101
= 0.267
104. Qn.2(UNEB 1999/P1/4)
The solubility of copper(II) sulphate
at 30oC is 25g per 100g of water.
The mass of copper(II) sulphate that
would crystallise if a solution
containing 50g of copper(II) sulphate
in 100g of water at 60oC is cooled to
30oC is
A. 12.5g B. 25.0g
C. 50.0g D. 75.0g
105. Qn.3(UNEB 2009/P1/32)
Which one of the following is true
about bases?
A. are soluble in water.
B. are hydroxides.
C. neutralise acids.
D. are oxides.
106. Qn.4(UNEB 1994/P1/14)
Which one of the following acids
can react with a base to produce
an acid salt?
A. Nitric acid.
B. Ethanoic acid.
C. Sulphuric acid.
D. Hydrochloric acid.
108. Qn.6(UNEB 2005/P1/14)
Which one of the following nitrates
will produce nitrogen dioxide when
strongly heated?
A. Potassium nitrate.
B. Sodium nitrate.
C. Zinc nitrate.
D. Ammonium nitrate.
109. Qn.7(UNEB 2006/P1/12)
When heated strongly, lead (II)
nitrate leaves a solid residue whose
colour is
A. reddish brown (hot), grey (cold).
B. yellow (hot), white cold.
C. reddish brown (hot), yellow (cold)
D. reddish brown (hot), white (cold)
110. Qn.8(UNEB 2006/P1/36)
The substance which does not produce
carbon dioxide when heated strongly
is
A. Calcium carbonate
B. Sodium carbonate
C. Potassium hydrogencarbonate
D. sodium hydrogencarbonate
111. Qn.9(UNEB 1987/P1/15)
10g of a saturated sodium chloride
solution was evaporated and 6g of
solid sodium chloride was left. The
solubility of sodium chloride is
A.
6 100
10
B.
6 100
4
C.
6 100
16
D.
10 100
16
113. The mass of potassium nitrate
which would dissolve in 25g of
water at 30oC is
A. 0.6g
B. 1.2g
C. 6.0g
D. 12.0g
114. Qn.11(UNEB 2002/P2/2)
(a) Define the terms:
(i) a normal salt
(ii) an acid salt
(b) Give one example of
(i) a normal salt
(ii) an acid salt
115. Qn.12(UNEB 1995/P2/13(a))
Sulphuric acid is a strong dibasic
acid.
(i) Explain the terms strong acid
and basicity.
(ii) Write an equation to show how
sulphuric acid ionizes in water.
116. Qn.13(UNEB 1991/P2/12)
(a) Describe how you would prepare
pure crystals of lead(II) nitrate in
the laboratory starting from
lead(II) oxide. Write an equation
for the reaction that takes place.
(b)State what happens when lead(II)
nitrate is strongly heated.
117. (c) State what is observed if
ammonia solution in gradually
added to a solution of lead(II)
nitrate until the alkali is in
excess. Write an equation for
the reaction that takes place.
118. Qn.14(UNEB 2004/P2/3)
A mixture containing copper(II) sulphate
and copper(II) carbonate was shaken
with excess water and filtered.
(a) Identify the residue.
(b) The dry residue was heated strongly.
(i) State what was observed.
(ii) Write an equation for the
reaction.
119. (c) (i) Name a reagent that can be
used to identify the anion in
the filtrate.
(ii) Write an ionic equation for
the anion and the reagent
you have named in (c)(i).
120. Qn.15(UNEB 1994/P2/11)
(a) Copper(II) carbonate was heated
strongly until there was no further
change.
(i) State what was observed.
(ii) Write an equation for the
reaction.
(iii) Name one reagent which can be
used to identify the gaseous
product.
121. (b) Excess dilute sulphuric acid
was added to the residue
in (a) and the mixture
warmed.
(i) State what was observed.
(ii) Write an equation for the
reaction.
122. (c) To the product in (b) was
added dilute sodium
hydroxide solution dropwise
until in excess.
(i) State what was observed.
(ii) Write a equation for the
reaction.
123. Qn.16(UNEB 1991/P2/9)
Copper(II) sulphate-5-water decomposes
when heated.
(a) State what would be observed when
copper(II) sulphate-5-water is
strongly heated.
(b) Write an equation for the reaction.
(c) Name one reagent that can be used
to convert the residue back to
copper(II) sulphate.
124. Qn.17(UNEB 2018/P2/9)
(a) When a sample of copper(II)
nitrate was strongly heated, a
reddish brown gas was evolved.
(i) Identify the gas.
(ii) Write the formula of the
residue.
125. (b) A sample of copper(II) nitrate
contaminated with zinc nitrate was
dissolved in water and the solution
was treated with excess sodium
hydroxide solution and then
filtered. Identify the cation in the
(i) Filtrate
(ii) Residue
126. (c) The residue from (b) was
strongly heated.
(i) State what was observed.
(ii) Write equation for the
reaction that took place.
127. Qn.18(UNEB 2014/P2/7)
(a) When a nitrate of a metal Y was
heated strongly, brown fumes were
observed together with a solid
residue which was reddish brown
when hot and yellow when cooled.
(i) Identify Y.
(ii) Write equation for the reaction
that took place.
128. (b) The residue from (a) was
heated with dilute nitric acid.
Write equation for the reaction
that took place.
(c) To the product in (b), dilute
sodium hydroxide was added drop
wise until there was not further
change. State what was
observed.
129. Qn.19(UNEB 1997/P2/12)
(a) Describe briefly how copper (II)
sulphate crystals can be prepared
from copper(II) oxide.
(b) What would be observed if
(i) sodium hydroxide solution was
gradually added to a solution of
copper(II) sulphate until the
alkali was in excess?
130. Write the equation for the
reaction that took place.
(ii) hydrated crystals of copper
sulphate were heated strongly?
131. Qn.20(UNEB 1991/P2/2)
2.5g of zinc carbonate was heated
strongly until there was no further
change.
(a) State what was observed.
(b) Write an equation for the
reaction.
132. Qn.21(UNEB 1994/P2/10)
The table below shows the
solubilities of a salt P in water at
different temperatures.
Temperature (oC) 10 20 30 40 50 60
Solubility (g/100g of
water)
18 20 24 30 38 50
133. (a) Plot a graph of solubility of P
against temperature.
(b) Use your graph to determine
the solubility of P at 25oC.
(c) Calculate the mass of P that
would dissolve in 45g of water
at 25oC.
134. Qn.22(UNEB 1995/P2/12)
(a) Explain what is meant by the
term saturated solution.
(b) Describe how the solubility of
potassium chloride can be
determined in the laboratory.
135. (c) The table below shows the
solubilities of potassium chloride
and potassium nitrate at various
temperatures.
Temperature (oC) 0 20 40 60
Solubility of potassium
chloride(g)
28.2 33.5 38.8 44.7
Solubility of potassium
chloride(g)
12.9 31.8 61.2 108.2
136. (i) On the same axes, plot graphs
of solubilities of potassium
chloride and potassium against
temperature.
(ii) Determine the temperature at
which the concentrations of the
two salts are equal.
137. (iii)Which of the two salts
dissolves more rapidly with
increase in temperature?
(iv)State what would happen if a
saturated solution of
potassium chloride at 40oC
was cooled to 30oC.
138. Qn.23(UNEB 1999/P2/6)
The solubility of hydrated Copper(II)
sulphate, CuSO4.5H2O in moles per litre
at various temperatures is shown in
figure 2 below.
139. (a) Determine the solubility of
hydrated copper(II) sulphate at
80oC.
(b) Calculate the solubility of
hydrated copper(II) sulphate in
g/100g of water at 80oC.
(H = 1, 0 = 16, S = 32, Cu = 64)
