1. Oxides can be acidic, basic, neutral or amphoteric depending on whether they react with acids and bases to form salts and water. Acidic oxides form salts with bases while basic oxides form salts with acids. 2. Acids are compounds that form hydrogen ions in aqueous solution and have characteristic properties like turning litmus red. Strong acids fully ionize while weak acids partially ionize. 3. Bases are compounds that form hydroxide ions in aqueous solution and have properties like turning litmus blue. Alkalis are soluble bases that produce hydroxide ions in water.

1. 1
Oxides, Acids, Bases and Salts
Oxides – When elements burn in air, they combine with oxygen to form oxides
1. (a) Acidic Oxides are oxides of non-metals which neutralise bases to form a salt and water only.
CO2 (g) + NaOH(aq) → Na2CO3 (aq) + H2O (l)
Acidic oxide + base → salt + water
e.g. CO2, SO2, SO3, NO2, SiO2, P2O3, P2O5
(b) Acid Anhydrides are acidic oxides which dissolve in water to form acidic
solutions.
CO2 (g) + H2O (l) → H2CO3 (aq) (carbonic acid)
SO3 (g) + H2O (l) →H2SO4 (aq) (sulphuric acid)
P2O5 (s) + 3H2O (l) → 2H3PO4 (aq) (phosphoric acid)
NB: Nitrogen dioxide (NO2) is a mixed acid anhydride because it dissolves in water to form
two acids. 2NO2(g) + H2O(ℓ) → HNO2(aq) + HNO3(aq)
(Nitrous acid) (Nitricacid)
2. Basic Oxides are oxides of metal which neutralise acids to form salt and water only.
MgO (s) + H2SO4 (aq) → MgSO4 (aq) + H2O (l)
Basic oxides + acid → salt + water
Most basic oxides are insoluble in water. The soluble basic oxides (oxides of group 1 and
group II metals) dissolve to form alkalis.
Very K2O (s) + H2O (l) → 2KOH(aq) (potassium hydroxide)
Soluble Na2O (s) + H2O (l) →2NaOH(aq) (sodium hydroxide)
Slightly soluble → CaO (s) + H2O (l) → Ca (OH) 2 (aq) (calcium hydroxide)
Sparingly Soluble → MgO(s) + H2O (l) → Mg (OH) 2 (aq) (magnesium hydroxide)
Alkaline oxides: basic oxides that dissolves in water to form an alkali- K2O, Na2O, CaO, MgO
3. Neutral Oxides are the few oxides of non-metals that neither neutralise acids nor bases
e.g. Carbon monoxide (CO), Nitric oxide (NO), Water (H2O)
4. Amphoteric Oxides are the few oxides of metals which neutralises both acids and bases.
e.g. Al2O3, ZnO, PbO
ZnO(s) + H2SO4 (aq) → ZnSO4 (aq) + H2O (l)
(Acid)
ZnO(s) + 2NaOH(aq) → Na2ZnO2 (aq) + H2O (l)
(Base) (Sodium zincate)
Acids
An acid is a compound which forms hydrogen ions (H+) as the only positively charged ions in aqueous
solution.
NB: . The acidic behaviour of acid is due to the presence of hydrogen ions. The acids will not show
its acidic behaviour in the absence of water, this is because acids do not dissociate to produce H+ (aq)
ions in the absence of water.
Acid in everyday life
Citric acid – citrus fruits Lactic acid – `sour` milk
Ethanoic acid – vinegar Tartaric acid – grapes, baking powder
Carbonic acid – aerated (soft) drinks Methanoic acid – ants `sting`
Mineral acids
Hydrochloric acid (HCl (aq)), Sulphuric acid (H2SO4 (aq)), Nitric acid (HNO3 (aq))
State symbols
aq = aqueous
g = gas
l = liquid
s = solid

2. 2
Properties of acids
1. -have a sharp, sour taste
2. -turn damp blue litmus red
3. -neutralise bases to form a salt and water only
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(ℓ)
Acid + base → salt + water
4. -liberate CO2 from carbonates (CO3
2-) and hydrogen carbonates (HCO3
-) with effervescence
[effervescence is the rapid bubbling of gas out of solution]
CaCO3 (s) + 2HCl (aq) → CaCl2 (aq) + CO2 (g) + H2O (l)
2NaHCO3 (s) + H2SO4 (aq) → Na2SO4 (aq) + 2CO2 (g) + 2H2O (l)
5. -The more reactive metals (Mg, Zn, and Fe) liberate hydrogen from dilute acids with
effervescence
Zn (s) + H2SO4 (aq) → ZnSO4 (aq) + H2 (g)
6. -liberate SO2 from sulphites (SO3
2-)
Na2SO3 (s) + 2HCl (aq) → 2NaCl(aq) + SO2 (g) + H2O (l)
NB: The mineral acids (HCl, H2SO4, HNO3) are corrosive (burn) when concentrated
The basicity (proticity) of an acid is the number of moles of hydrogen (H+) ions produced from one
mole of the acid in aqueous solution.
(a) Examples of monobasic (monoprotic) acids
HCl (aq) → H+
(aq) + Cl−
(aq)
HNO3 (aq) → H+
(aq) + NO3
−
(aq)
CH3COOH (aq) (Ethanoic acid) ⇌ H+
(aq) + CH3COO-
(aq)
(b) Examples of dibasic (diprotic) acids
H2SO4 (aq) → 2H+
(aq) + SO4
2−
(aq)
H2CO3 (aq) ⇌ 2H+
(aq) + CO3
2−
(aq)
(c) Example of a tribasic (triprotic) acid
H3PO4 (aq) ⇌ 3H+
(aq) + PO4
3−
(aq)
Strong and weak acids
NB: The term `strength` relates to the degree of ionisation of the acid or alkali in aqueous solution.
The acidity of a solution is proportional to the number of H+ ions present.
The term `concentrated` relates to amount of solute (acid or alkali) used to make up a given
volume of solution.
.A Strong Acid is one which is completely ionised in aqueous solution
e.g. HCl (aq) → H+
(aq) + Cl−
(aq)
HNO3 (aq) → H+
(aq) + NO3
−
(aq)
H2SO4 (aq) → 2H+
(aq) + SO4
2−
(aq)
A Weak Acid is one which is partially ionised in aqueous solution. It consists mainly of acid
`molecules` and its ionisation is reversible.
e.g. H2CO3 (aq) 2H+
(aq) + CO3
2−
(aq)
CH3COOH (aq) H+
(aq) + CH3COO−
(aq)
Bases

3. 3
A Base is a compound which contains oxide (O2-) or hydroxide (OH-) ions and which neutralises acids
to form a salt and water only (A base is a metal oxide or metal hydroxide).
Base
(Metal oxide or metal hydroxide)
Insoluble base soluble base (alkalis)
An Alkali is a compound which produces hydroxide ions (OH−) as the only negatively charged ions in
aqueous solution (an alkali is a soluble base).
Neutralisation is the reaction between the hydrogen (H+) ions of an acid and the oxide (O2-) or
hydroxide (OH-) ions of a base to form water. A salt is also formed
H+
(aq) + OH−
(aq) → H2O (l)
2H+
(aq) + O2−
(aq) → H2O (l)
Common Alkalis
Chemical Name Formula Common Name Solubility In Water Type Of Alkali
Potassium Hydroxide KOH Caustic Potash Very soluble Strong
Sodium Hydroxide NaOH Caustic Soda Very soluble Strong
Calcium Hydroxide Ca(OH)2 Lime Water Slightly Soluble
Aqueous Ammonia NH3 (aq) --- --- Weak
Magnesium Hydroxide Mg(OH)2 Milk of Magnesia Slightly Soluble
NB: Zn (OH) 2, Pb (OH) 2, Al (OH) 3, are Amphoteric Hydroxides as they neutralize both acids and bases
to form a salt and water only.
Properties of Alkalis
1. -Have a soapy feeling and a bitter taste.
2. -Turn red litmus blue
3. -Neutralize acids to form salt and water only.
MgO (s) + H2SO4 → MgSO4 (aq) + H2O (l)
4. -Liberate ammonia (NH3) from ammonium salts
NH4Cl (s) + NaOH (aq) → NaCl (aq) + NH3 (g) + H2O (l)
5. -Precipitate many hydroxides from solutions of their salts.
CuSO4 (aq) + 2NaOH (aq) → Cu (OH) 2 (s) + Na2SO4 (aq)
NB: Alkaline solutions will absorb CO2 from the atmosphere resulting in a white crust being formed
on the inside of the bottle. This often makes it difficult to remove glass stoppers from the
bottles.
Strong Alkalis – completely ionized in aqueous solution.
NaOH(aq) → Na+
(aq) + OH−
(aq)
KOH(aq) → K+
(aq) + OH−
(aq)
Weak alkali – slightly ionized in aqueous solution
NH3(aq) + H2O(ℓ) ⇌ NH4
+
(aq) + OH−
(aq)

4. 4
The pH scale
The pH scale is a measure of the acidity (concentration of H+ ions) or alkalinity (concentration of
OH- ions) of a solution. The pH of a solution can be determined using universal indicator or a pH
meter.
Universal indicator is a mixture of indicators which changes colour depending on the acidity or
alkalinity of a solution.
HCl
(aq)
acid
H
2
SO
4
(aq)
acid
Lime
juice
(citric
acid)
Ethanoic
acid
(vinegar)
Tomato
juice
Carbonic
acid
(sodas)
Normal
rain
Milk
pure
water
Blood
Toothpaste
Baking
soda
Milk
of
magnesia
Aqueous
ammonia
Lime
water
(Ca(OH)
2
NaOH
(aq)
KOH
(aq)
Red Orange Yellow Lime green Green Blue- green Blue Violet(purple)
1 2 3 4 5 6 7 8 9 10 11 12 13 14
Acid Rain
Acidity increasing Alkalinity increasing
Neutral
pH Range
Neutral = 7
Acidic = < 7
Alkalinity = > 7
TEST FOR COMMON GASES
Test for Oxygen this gas relights a glowing splint
Test for Hydrogen this gas puts out a lighted splint with a ‘pop’
Test for Carbon dioxide this gas turn lime water milky
Test for Sulphur dioxide this gas turns acidified potassium permanganate from purple to colourless
Test for Ammonia this gas has a pungent smell and turn damp red litmus blue
SALTS
There are two types of salts,
 Acid salt
 Normal salt
The type formed depends on the quantity of acid used.
An acid salt is one formed when the hydrogen (H+) ions of an acid are partially replaced by a metal ion
or the ammonium (NH4
+) ion. Acid salts, therefore, contain some H+ ions from the original acid; hence,
their solutions turn blue litmus red.
NB Only dibasic and tribasic acids can form acid salts.
e.g. H2SO4 (aq) + NaOH (aq) → NaHSO4 (aq) + H2O (l)
Sodium hydrogensulphate (acid salt)
H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2H2O (l)
Sodium sulphate (normal salt)
Colour changes of common indicators
Indicator Acid Neutral Alkali
Litmus Red Purple Blue
Screened methyl orange Red Colourless Green
Methyl orange Pink Orange Yellow
Phenolphthalein Colourless Colourless Pink

5. 5
A normal salt is one formed when the hydrogen (H+) ions of an acid are completely replaced by a metal
ion or ammonium ion
2HCl (aq) + CaO(s) → CaCl2(aq) + H2O (l)
NB The solution of normal salts does not turn blue litmus red.
Naming of salts
Acid Salt formed Anion present Example
Hydrochloric acid Chlorides Cl- NaCl
Ethanoic acid Ethanoates CH3COO- CH3COONa
Nitric acid Nitrates NO3
- NaNO3
Sulphuric acid Hydrogen sulphate (acid salts),
Sulphates (normal salt)
HSO4
-
SO4
2-
NaHSO4
Na2SO4
Carbonic acid Hydrogen carbonate (acid salts),
Carbonates(normal salt)
HCO3
-
CO3
2-
NaHCO3
Na2CO3
Phosphoric acid Dihydrogen phosphates (acid salts),
Hydrogen phosphates (acid salts),
Phosphates (normal salt)
H2PO4
-
HPO4
2-
PO4
3-
NaH2PO4
Na2HPO4
Na3PO4
Soluble salts and insoluble salts
Salts Soluble Insoluble
Chlorides Most are soluble Silver chloride- AgCl,
Lead (II)chloride-PbCl2 (soluble in hot water)
Sulphates Most are soluble Barium, silver, and lead (II) sulphates,
Calcium sulphate is slightly soluble
BaSO4, CaSO4, PbSO4 and Ag2SO4
Nitrates All are soluble None
Carbonates Sodium and potassium
carbonates, Na2CO3 and
K2CO3 and (NH4)2CO3
Most are insoluble
Ethanoates All are soluble None
Sodium, potassium
and ammonium salts
All are soluble None
Salts used in everyday life
Salts
Colour and other
characteristics Uses
Ammonium chloride White crystals Dry cells (batteries), fertilizers
Ammonium sulphate (sulphate of
ammonia)
White crystals fertilizers
Calcium carbonate (marble; chalk;
limestone)
White but can be
coloured
Decorative stones, manufacture of
cement and lime
Calcium sulphate (plaster of paris,
gypsum)
White crystals Plastering walls; making casts, etc.
Magnesium sulphate (Epsom salts) White crystals Purgative (laxative)
Copper (II) sulphate Blue crystals Fungicides
Sodium carbonate (washing soda) White crystals or
powder
In cleaning, in laundry as a water
softener, in the manufacture of glass
Sodium hydrogen carbonate
(baking soda)
White crystals Baking
Sodium chloride (common salt) White crystals Seasoning
Potassium nitrate (salt petre) White crystals Curing meats eg. hams

6. 6
Acid rain
Caused by: Oxides of sulphur and nitrogen
The pH of natural rain is 6.0 – 6.5. Acid rain is rainwater of a lower pH. This can reach values of pH4–5.
The reactions that lead to the formation of acid rain are not fully understood. However some reactions are
thought to occur as follows:
2SO2(s) + 2H2O (g) + O2 (g) → 2H2SO4 (aq)
2NO2 (g) + H2O (g) → HNO2 (aq) + HNO3 (aq)
Some effects acid rain can have on the environment are:
 Death of aquatic organisms, which can affect food supplies.
 Destruction of coral reefs, which can affect tourism.
 Increase in the solubility of heavy metals, which can lead to their accumulation in waterways and
the consequent death of organisms.
 Damage of leaves, bodies and shoots of trees.
 Washing away of nutrients from the soil.
 Corrosion of monuments; buildings, statues, metals e.g. iron.
CaCO3(s) + H2SO4(aq) → CaSO4(aq) + CO2(g) + H2O(ℓ)
Fe(s) + HNO3(aq) → Fe(NO3)2(aq) + H2(g)
Pollutant Major sources Major Effects
Oxides of sulphur
SO2, SO3
Combustion of fossil fuels e.g. coal, smelting of sulphide
ores to produce Zinc, lead and copper.
Acid rain
Oxides of nitrogen
NO, NO2
Produced at high temperatures in the internal combustion
engine
smog, acid rain
Acid soils
Soils become acidic because of (i) acid rain or (ii) treatment with ammonium fertilizers e.g. Ammonium
sulphate. Soils treated with (NH4)2SO4 become acidic with sulphuric acid overtime. However, plants
thrive best in soils with pH between 6.3 -7. Liming the soil by adding slaked lime (solid calcium
hydroxide), quick-lime (CaO) or powdered limestone (CaCO3) neutralises the soil.
e.g., H2SO4 (aq) + Ca (OH)2 (s) → Ca SO4 (s) + 2 H2O (l)
acid in soil lime
Liming also improves drainage as it causes soil particles to clump together. This is a physical change.
NB: Lime and ammonium fertilizers should not be added to the soil at the same time as the nitrogen
needed by the plants to make protein will be lost to the atmosphere as ammonia (alkalis liberate
NH3 from ammonium salts)
e.g., (NH4)2SO4 (aq) + Ca (OH) 2 (s) → Ca SO4 (s) + 2NH3 (g) + H2O (l)
NB: Bee sting is acidic; hence treat with baking soda (alkaline)
Wasp sting is alkaline; hence treat with vinegar (ethanoic acid)
WORK SMART (BE PRODUCTIVE) TO ACHIEVE MORE!