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Acids, Bases & Salts
• Characteristic Properties of Acids and
Bases
• The pH scale, pH indicators and
measurement of pH
• Classification of Oxides
• Salts: Uses and Preparation
ACIDS
An acid is a substance which when dissolved in water produces
hydrogen ions (H+) as the only positive ion in solution.
Types of acids
Organic acids
(from plant or animal material)
Mineral acids or Inorganic acids
(from mineral elements or inorganic material)
 acetic acid (vinegar or ethanoic acid)
hydrochloric acid
formic acid (from ants)
sulfuric acid
ascorbic acid (vitamin C)
nitric acid
citric acid
phosphoric acids
tartaric acid
carbonic acid
amino acids (subunits of proteins)
lactic acid (fermented milk)
Pure acids
Pure acids do not have water added to them.
ie. HCl(l), HCl(g) H2SO4(l), HNO3(l)
Pure acids exist as covalent molecules.
They only ionise when they dissolve in water.
As they do not possess hydrogen ions, they do
not exhibit acidic properties.
No water => no hydrogen ions
Water must be present for an acid to exhibit its
acidic properties.
Strong and Weak acids
 A strong acid is an acid which undergoes
complete ionization in water (ie. all the
molecules ionize to produce H+ ions. There are
no molecules left in solution.)
 H2SO4(aq)
2H+(aq) + SO42-(aq)
 eg. H2SO4,HCl, HNO3
 A weak acid is an acid which undergoes only
partial ionization in water. (ie, most of the acid
molecules remain uncharged in water, only a
few molecules ionize to release H+ ions)
 CH3COOH(aq)
CH3COO-(aq) + H+(aq)
 H2SO3(aq)
2H+(aq) + SO32-(aq)
 H2CO3(aq)
2H+(aq) + CO32-(aq)
H+
SO42-
SO42H+
H+ SO42-
CH3COOH
H+ CH COO3
CH3COOH
Basicity of an Acid
• The basicity of an acid refers to the number of hydrogen
ions that can be produced per molecule of acid when
dissolved in water.
• Example: Sulfuric acid is a dibasic acid as it produces
two hydrogen ions per molecule of acid when dissolved
in water.
• Basicity and strength of some common acids
Basicity
Strong acid
Weak acid
Monobasic
, HCl,
Dibasic
HNO3
H2SO4
CH3COOH
H2SO3,
H2CO3
Tribasic
H3PO4
Properties of acids
1. Dilute acids have a sour taste.
2. Acids are soluble in water and give solutions with
a pH below 7. They give charatcteristic colours
with indicators.They turn litmus red.
3. The concentrated forms of strong acids like
hydrochloric acid and sulphuric acid are
corrosive.
4. Aqueous solutions of acids are able to conduct
electricity due to the presence of mobile ions
Reactions of Acids
1 Reaction with metals
 Dilute acids reacts with metals above hydrogen in the
electrochemical series to liberate hydrogen.
 Nitric acid (HNO3) doesn't usually form hydrogen with a
metal, instead you get brown fumes of nitrogen dioxide!
but you still get the metal nitrate salt.
 exceptions: Al + H2SO4
oxide layer)




X
(Al is unreactive due to
Ca(s) + H2SO4(aq)
CaSO4(s) + H2(g)
Ba(s) + H2SO4(aq)
BaSO4(s) + H2(g)
Pb(s) + H2SO4(aq)
PbSO4(s) + H2(g)
Salt formed is insoluble and coats metal preventing further
reaction.
Reactions of Acids
2 Reaction with carbonates
 carbonate + acid
salt + carbon dioxide + water
 exception:
 Aluminium carbonate does not exist.
 CaCO3(s) + H2SO4(aq)
CaSO4(s) + H2O(l) + CO2(g)
Poor yield - similarly with PbCO3 and BaCO3.
3 Reaction with bases(metal oxides or hydroxides)
 Dilute acids neutralise bases to form a salt and water only.
 eg. 2NaOH(s) + H2SO4(aq)
Na2SO4(aq) + 2H2O(l)
Comparison of the properties of
strong and weak acids
Property
Strong Acid
Weak Acid
effect on litmus
red
pink
action on magnesium
rapid effervescence
of hydrogen gas
slow effervescence
of hydrogen gas
action on Na2CO3
rapid effervescence
of carbon dioxide
slow effervescence
of carbon dioxide
conductivity
conducts electricity
poor conductor
Uses of acids
1 Preserving food
 eg. ethanoic acid, benzoic acid, citric acid
2 Manufacture of products like paint, detergent and
fertilizers.
 eg. sulfuric acid, hydrochloric acid and nitric acid.
3 Batteries for vehicles
 eg. sulfuric acid
4 removing rust from iron or steel
 eg. sulfuric acid
5 making rubber from latex
 eg methanoic acid
6 Pure fruit juice
 eg. citric acid and tartaric acid
Bases and Alkalis
 A base is a substance which can react with acids to form
salt and water only. eg. CuO, MgO, CaO, KOH
 Bases are generally metal oxides or hydroxides.
 NOTE:
All basic oxides are ‘soluble’ in acids (react).
But only a few are soluble in water.
Basic oxides which are soluble in water form
hydroxides called alkalis.
 An alkali is a basic oxide which is soluble in water.
 eg. sodium hydroxide, potassium hydroxide, calcium
hydroxide, barium hydroxide & aqueous ammonia.
 Neutralization is the process whereby an acid reacts
completely with an appropriate amount of alkali to form a
salt and water only.
 H+ (aq) + OH- (aq)
H2O (l)
Strong and weak alkalis
 A strong alkali is one that undergoes complete
ionization in water.
OH- Na+
 NaOH(aq) → Na+(aq) + OH-(aq)
Na+ OH KOH(aq)
→
K+(aq) + OH-(aq)
 A weak alkali is one that undergoes only partial
ionization when dissolved in water.
NH4+ NH3
 NH3(g) + H2O
NH3(aq) + H2O
 NH3(aq) + H2O
NH4+(aq) + OH-(aq)
NH3 OH-
Note: Calcium hydroxide is considered as a weak alkali as it is only slightly
soluble in water.
Properties of alkalis
1
2
3

4

Alkalis have a bitter taste.
Alkalis are soapy to the touch.
Alkalis have a pH above 7
They turn litmus blue.
Concentrated forms of potassium hydroxide
and sodium hydroxide are known to be
corrosive.
Aqueous solutions of alkalis are able to conduct
electricity due to the presence of mobile ions.
Reactions of bases / alkalis
Action on acids
A All bases react with acids to form a salt and
water only.
 Base + Acid
 MgO + 2HNO3
 KOH + HCl
Salt + water
Mg(NO3)2 + H2O
KCl + H2O
B All bases react with acidic gases to form a
salt ( + water for alkalis)
 CaO + CO2
 CuO + SO2
 2NaOH + CO2
CaCO3
CuSO3
Na2CO3 + H2O
Reactions of bases / alkalis
Alkalis react with ammonium salts in the
presence of heat to produce ammonia (NH3)
gas.
 alkali + ammonium salt
ammonia + salt + water
 NaOH + NH4NO3
NH3 + NaNO3 + H2O
 Ca(OH)2 + 2NH4Cl
2NH3 + 2H2O + CaCl2
 Alkalis precipitate many insoluble hydroxides from their
salts
 Alkali + salt solution
precipitation reaction
Point toNote:
 Alkalis dissolve in water
to produce hydroxide
ions.
 It is the hydroxide ions
that give alkalis their
properties
Uses of alkalis
1 Neutralise acids
eg. tooth paste, antacids
2 Dissolve dirt and grease
eg. soap, detergents
pH scale
 A scale to measure the acidity or alkalinity of a
solution.
 The scale ranges from 0 - 14.
 pH less than 7 => an acidic solution
 pH = 7 => a neutral solution
 pH more than 7 => an alkaline solution
 Strong acids : eg. HCl, H2SO4
pH = 1 or 2
 Weak acids : eg. H2CO3 ,H2SO3 pH = 5 or 6
 Weak alkalis:eg. NH3(aq)
pH = 9 or 10
 Strong alkalis : eg. NaOH, KOH pH = 13 or 14
pH Indicators
Indicators are compounds which change colour in
accordance to the pH of the medium.
Indicators are commonly used for determining the pH of
colourless liquids.
Although they are not very accurate or sensitive , they give
quicker results.
Each indicator has an acidic and alkaline colour and a pH at
which it changes colour .
Indicator
Colour in medium which is
strongly acidic strongly alkaline
water
pH at which
colour changes
methyl orange
pink
yellow
yellow
4 (orange)
litmus
red
blue
red
8 (purple)
phenolphthalein colourless
pink
colourless
screened
methyl orange
green
green
red
10 (colourless)
4 (grey)
Colour Changes of some Indicators
methyl
orange
pH 3.7
screened
methyl
orange
pH 3.7
pH 9.3
phenolphthalein
pH 7
litmus
Universal
indicator
pH 1 2 3
4
5
6
7
8
9 10 11 12 13 14
Universal Indicator
 It is made up of a mixture of indicators working at
different pH ranges.
 The universal indicator is also known as pH
indicator.
 The pH of a solution can be measured by dipping a
piece of universal indicator paper (pH paper) in the
solution and comparing the colour obtained with a
standard colour chart. Alternatively a pH meter can
be used
pH range 1 2 3 4 5 6 7 8 9 10 11 12 13 14
colour red
orange yellow green blue
purple/violet
OXIDES
Acidic
• Oxides of
non metals
• turn moist
litmus red
• reacts
with alkalis
to form
salt &
water
• e.g. SO2,
CO2, NO2
Basic
• Oxides of
metals
(Group 1 &
II)
• reacts with
acids to
form salt &
water
• e.g. Na2O,
MgO, CuO
Amphoteric
Neutral
• Oxides of
metals
• Oxides of
non metals
• reacts with
acids and
alkalis to form
salt & water
• no effect
on litmus
• e.g. Al2O3,
ZnO, PbO
• e.g. H2O,
NO, CO
Salts
 A salt is a substance formed when any of the replacable hydrogen ions
in an acid have been partly or completely replaced by an equivalent
number of metal or ammonium ions.
 Normal salts are formed when all the replaceble hydrogen ions in the
acid has been completely replaced by metallic or ammonium ions
 HCl + NaOH
NaCl + H2O
 H2SO4 + ZnO
ZnSO4 + H2O
 Acid salts are formed when the replaceble hydrogen ions in the acid
have been partially replaced by metallic or ammonium ions.
 H2SO4 + KOH
KHSO4 + H2O
 H2CO3 + NaOH
NaHCO3 + H2O
Normal salts and acid salts
Acid
Acid Salts
HCl
Normal Salts
NaCl, CaCl2, NH4Cl
H2SO4
KHSO4, Mg(HSO4)2
K2SO4, CuSO4,
NH4(SO4)2
H3PO4
NaH2PO4, Na2HPO4
Na3PO4,, (NH4)2PO4
HNO3
H2CO3
NaHCO3, Mg(HCO3)2
NaNO3 Mg(NO3)2,
Al(NO3)3, NH4NO3
K2CO3, CaCO3,
NH4(CO3)2
Salt Preparation
Soluble salts
Na+, K+
NH4+ salts
Titration
Soluble salts
that are
NOT Na+,
K+ NH4+
salts
acid +
soluble
base
/oxide/
carbonate
Excess solid
(metal, base,
carbonate) +
acid
Insoluble salts
Precipitation
Points to note
 The procedure used in preparing salts are such that the salts obtained
are of a high degree of purity.
 eg. A + B
salt X + D
 The method used in making salt X must be such that salt X obtained is
free off excess reagent A and B and by-product D.
 The method used depends on whether the salts are soluble or insoluble
in water.
Preparation of soluble salts
1 Action of an acid on a metal.
acid + metal
salt + hydrogen gas
2 Action of an acid on an insoluble base.
acid + base
salt + water
3 Action of an acid on an insoluble carbonate.
acid + carbonate
salt + water + carbon dioxide
4 Action of an acid on a soluble base or carbonate
(ie.Neutralization of an acid by a base - titration)
acid + base
salt + water
acid + carbonate
salt + water + carbon dioxide
Preparation of soluble salts that do
not contain Na+, K+ or NH4+
Adding Excess Solid to a fixed volume of Acid
• Reacting an acid with a metal or with an insoluble
base or carbonate.
• This method is suitable as one of the reagents
used is in the solid state and when added in
excess can be easily removed by filtration
1.
2.
3.
4.
5.
The required volume of acid is measured out into the beaker
with a measuring cylinder. The metal, oxide, hydroxide or
carbonate is weighed out and added in small portions to the
acid in the beaker with stirring.
The mixture maybe heated to speed up the reaction. When
no more of the solid dissolves it means the acid is
neutralised.
The hot solution is filtered to remove the excess solid
The filtrate is retained and heated in an evaporating dish to
saturate it. The solution is left to cool and crystallise.
The crystals are filtered out and rinsed quickly with cold
distilled water & dried between pieces of dry filter paper.
Example: Preparation of Copper (II) Sulfate
•
•
•
Reactants: Copper (II) oxide & Sulfuric acid
Reaction:
CuO(s) + H2SO4(aq)
CuSO4 (aq) + H2O(l)
• Procedure:
1. Add copper (II) oxide powder until excess to a fixed volume
of warm dilute sulfuric acid i.e. till the copper (II) oxide
forms a residue at the base of the vessel.
2. Filter the mixture to remove the excess copper (II) oxide.
3. Heat the filtrate, which is copper(II) sulfate solution until it is
saturated and allow it to cool. Crystals of copper (II) sulfate
will form.
4. Filter off the crystals and dry them between two pieces of
filter paper.
Preparation of Na+, K+ , NH4+ salts
1. A known volume of alkali is pipetted into a conical flask and screened methyl
orange indicator added.
2. The acid is titrated with the alkali in the burette until the indicator turns from
green to grey.
3. The volume of acid needed for neutralisation is then noted, this is called the
endpoint.
4. Steps (1-3) are repeated with both known volumes mixed together BUT
without the contaminating indicator. i.e. no indicator is added
5. The solution is transferred to an evaporating dish and heated to saturate it.
6. The solution is left to cool to complete the crystallisation.
7. The residual liquid can be filtered away and the crystals can be carefully
collected and rinsed & dried between 2 pieces of filter paper.
• In this method, both the reactants used are in
the aqueous state thus excess reagent cannot
be easily removed from the product which is in
the same state. Thus exact quantities of
reactants must be used to ensure that product is
not contaminated by excess reagent.
• The two main steps involved are:
1. Finding the exact quantity of reactant required
for reaction by titration
2. Using these quantities to prepare the salt.
• This method of neutralising an acid with a
soluble base (e.g. sodium hydroxide) is called a
titration.
Example: Preparation of Sodium Chloride
•
•
•
1.
2.
3.
4.
5.
Reactants: sodium hydroxide hydrochloric acid
Reaction: NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
Procedure:
Pipette a fixed volume of sodium hydroxide into a
conical flask. Add a few drops of methyl orange
indicator. Slowly run the hydrochloric acid from a
burette till one drop of acid added turns the indicator
from yellow to orange.
Note the volume of acid added.
Repeat the whole procedure without adding the
indicator.
Evaporate the solution obtained till it is saturated and
allow it to cool. Crystals of sodium chloride will form.
Filter off the crystals and rinse quickly with cold distilled
water & dry them between pieces of dry filter paper.
Preparation of insoluble salts
 Precipitation reactions
 These reactions involve the addition of two soluble compounds; one
containing the metallic ion and the other ion of the desired salt.
 Exchange of ions between the two compounds occurs.
 Formation of desired salt as a precipitate and the by-product as an
aqueous solution.
 The desired salt is obtained by filtration.
 The salt is washed thoroughly to remove any excess reagent or by
product, then dried by pressing between pieces of filter paper
 eg. CuSO4 + 2NaOH
AgNO3 + NaCl
Cu(OH)2 + Na2SO4
AgCl + NaNO3
 Note:the reagents used must be in the form of aqueous solutions so
that the insoluble salt formed being the only solid present can be
completely separated from the other chemicals present.
 To prepare insoluble salt XY, use:
X nitrate + sodium Y
insoluble salt XY + sodium nitrate
Example: Preparation of Lead (II) sulfate
•
•
Reactants: sodium sulfate , lead(II) nitrate
Reaction:
•
Na2SO4 (aq) + Pb(NO3) 2 (aq)
NaNO3(aq) + PbSO4(s)
• Procedure:
1. Mix aqueous solutions of sodium sulfate and
lead(II) nitrate together.
2. Filter the mixture to remove the precipitate of
lead(II) sulfate formed.
3. Wash the residue thoroughly with distilled water.
4. Dry the salt between two pieces of filter paper.
Direct combination of two elements
 This method is confined to binary salts only ie.
salts made of two elements only.
eg. 2Fe(s) + 3Cl2(g)
2FeCl3
An excess of chlorine is added.
eg. Fe(s) + S(s)
FeS(s)
An excess of iron is added and the excess
reagent csn be removed using a magnet.
N
S