IGCSE CHEMISTRY
SECTION 4 LESSON 1
Content
The iGCSE
Chemistry
course
Section 1 Principles of Chemistry
Section 2 Chemistry of the Elements
Section 3 Organic Chemistry
Section 4 Physical Chemistry
Section 5 Chemistry in Society
Content
Section 4
Physical
Chemistry
a)
b)
c)
d)
Acids, alkalis and salts
Energetics
Rates of reaction
Equilibria
Lesson 1
a) Acids,
alkalis
and salts
4.1 describe the use of the indicators litmus, phenolphthalein and
methyl orange to distinguish between acidic and alkaline solutions
4.2 understand how the pH scale, from 0–14, can be used to
classify solutions as strongly acidic, weakly acidic, neutral, weakly
alkaline or strongly alkaline
4.3 describe the use of universal indicator to measure the
approximate pH value of a solution
4.4 define acids as sources of hydrogen ions, H+, and alkalis as
sources of hydroxide ions, OH¯
4.5 predict the products of reactions between dilute hydrochloric,
nitric and sulfuric acids; and metals, metal oxides and metal
carbonates (excluding the reactions between nitric acid and metals)
4.6 understand the general rules for predicting the solubility of
salts in water:
i all common sodium, potassium and ammonium salts are soluble
ii all nitrates are soluble
iii common chlorides are soluble, except silver chloride
iv common sulfates are soluble, except those of barium and calcium
v common carbonates are insoluble, except those of sodium,
potassium and ammonium
4.7 describe experiments to prepare soluble salts from acids
4.8 describe experiments to prepare insoluble salts using
precipitation reactions
4.9 describe experiments to carry out acid-alkali titrations.
pH scale and indicators
The pH scale is a measure of how acidic or alkaline a solution is.
pH scale and indicators
The pH scale is a measure of how acidic or alkaline a solution is.
neutral
ACIDIC
Very
Slightly
1 2 3 4 5 6
ALKALINE
Slightly
Very
7 8 9 10 11 12 13 14
pH scale and indicators
The pH scale is a measure of how acidic or alkaline a solution is.
neutral
ACIDIC
Very
Slightly
1 2 3 4 5 6
ALKALINE
Slightly
Very
7 8 9 10 11 12 13 14
A substance forms an aqueous solution when it dissolves
in water. Water itself is neutral.
When substances dissolve in water, they dissociate into
individual ions.
pH scale and indicators
Water, H2O
H+(aq)
OH-(aq)
hydrogen ion hydroxide ion
acid
alkali
pH scale and indicators
Indicators are special dyes that change colour
according to whether they are in acidic, alkaline
or neutral solutions.
pH scale and indicators
Indicators are special dyes that change colour
according to whether they are in acidic, alkaline
or neutral solutions.
Three common indicators are:
Litmus
Phenolphthalein
Methyl Orange
pH scale and indicators
Litmus
5.0
ALKALINE
neutral
ACIDIC
Very
Slightly
1 2 3 4 5 6
Slightly
Very
7 8 9 10 11 12 13 14
8.0
pH scale and indicators
Methyl Orange
4.4
6.2
neutral
ACIDIC
Very
Slightly
1 2 3 4 5 6
ALKALINE
Slightly
Very
7 8 9 10 11 12 13 14
pH scale and indicators
Phenolphthalein
ALKALINE
neutral
ACIDIC
Very
Slightly
1 2 3 4 5 6
Slightly
Very
7 8 9 10 11 12 13 14
8.3
10
pH scale and indicators
Universal indicator is a mixture of dyes, and
shows a complete colour range across the pH
scale.
Common acids and alkalis
ACIDS
Name
Hydrochloric acid
Sulphuric acid
Nitric acid
Formula
HCl
H2SO4
HNO3
Common acids and alkalis
ACIDS
Name
Hydrochloric acid
Sulphuric acid
Nitric acid
Formula
HCl
H2SO4
HNO3
ALKALIS
Name
Sodium hydroxide
Potassium hydroxide
Calcium hydroxide
Formula
NaOH
KOH
Ca(OH)2
Salt formation
When acids and alkalis react together,
salts are formed.
The general equation is:
Acid + Base  Salt + Water
Salt formation
When acids and alkalis react together,
salts are formed.
The general equation is:
Acid + Base  Salt + Water
This is known as a neutralisation reaction
because the products are neutral.
Salt formation
Salts consist of two parts – a metal part,
and the non-metal ion from the acid.
Acid
Ions in solution
Salts formed
Hydrochloric
acid
H+ Cl-
chlorides
Sulphuric acid
2H+ SO42-
sulphates
Nitric acid
H+ NO3-
nitrates
Salt formation
Examples of salts
Salt
Formula
Metal ion
Non-metal
ion
Copper
sulphate
CuSO4
Cu2+
SO42-
Sodium
chloride
Potassium
nitrate
NaCl
Na+
Cl-
KNO3
K+
NO3-
Calcium
sulphate
CaSO4
Ca2+
SO42-
Reactions of salts
Acids + Metals
Acid + Metal  Salt + Hydrogen
Reactions of salts
Acids + Metals
Acid + Metal  Salt + Hydrogen
Magnesium + Hydrochloric  Magnesium + Hydrogen
Acid
chloride
Mg + 2HCl  MgCl2 + H2
Reactions of salts
Acids + Metals
Have you got that?
Are you really
sure? Let’s try a
few examples.
Reactions of salts
Acids + Metals
Magnesium + Sulphuric Acid 
Iron + Hydrochloric Acid

Lead

+ Sulphuric Acid
Reactions of salts
Acids + Metals
Magnesium + Sulphuric Acid  Magnesium sulphate +
Hydrogen
Iron + Hydrochloric Acid
Lead
+
 Iron chloride + Hydrogen
Sulphuric Acid  Lead sulphate + Hydrogen
Reactions of salts
Acids + Metal oxides
Acid + Metal oxide  Salt + Water
Reactions of salts
Acids + Metal oxides
Acid + Metal oxide  Salt + Water
Copper + Sulphuric  Copper + Water
oxide
Acid
sulphate
CuO + H2SO4  CuSO4+ H2O
Reactions of salts
Acids + Metal oxides
Have you got that?
Are you really,
really sure? Let’s
try a few more
examples.
Reactions of salts
Acids + Metal oxide
Magnesium + Hydrochloric Acid 
oxide
Iron + Sulphuric Acid
oxide
Lead +
oxide

Hydrochloric Acid

Reactions of salts
Acids + Metal oxide
Magnesium + Hydrochloric Acid  Magnesium + Water
oxide
chloride
Iron + Sulphuric Acid
oxide
Lead +
oxide
 Iron sulphate + Water
Hydrochloric Acid
 Lead chloride + Water
Reactions of salts
Acids + Metal carbonate
Acid + Metal  Salt + Carbon + Water
carbonate
dioxide
Reactions of salts
Acids + Metal carbonate
Acid + Metal  Salt + Carbon + Water
carbonate
dioxide
Copper + Hydrochloric  Copper + Carbon + Water
Carbonate acid
chloride dioxide
CuCO3 + 2HCl  CuCl2 + CO2 + H2O
Reactions of salts
Acids + Metal carbonate
Guess what?
That’s right, no
more examples!
Solubility of salts
If a substance is soluble, then this means that it
will dissolve in a solvent.
Solubility of salts
If a substance is soluble, then this means that it
will dissolve in a solvent.
The most common solvent you will come across is
WATER.
Solubility of salts
If a substance is soluble, then this means that it
will dissolve in a solvent.
The most common solvent you will come across is
WATER.
Solute (the solid) + Solvent (water)
 Solution (aqueous)
Solubility of salts
There’s no easy way
around this – you’ve
just got to learn the
relative solubility of
salts!
Solubility of salts
All ammonium,
potassium and
sodium salts are
soluble in water.
Solubility of salts
All ammonium,
potassium and
sodium salts are
soluble in water.
All nitrates are
soluble in water
Solubility of salts
All ammonium,
potassium and
sodium salts are
soluble in water.
Most chlorides
are soluble in
water (except
lead and silver).
PbCl2 is soluble
in hot water.
All nitrates are
soluble in water
Solubility of salts
All ammonium,
potassium and
sodium salts are
soluble in water.
Most sulphates
are soluble in
water (except
barium, calcium
and lead)
Most chlorides
are soluble in
water (except
lead and silver).
PbCl2 is soluble
in hot water.
All nitrates are
soluble in water
Solubility of salts
All ammonium,
potassium and
sodium salts are
soluble in water.
Most carbonates
are insoluble in
water (except
sodium, potassium
and ammonium)
Most sulphates
are soluble in
water (except
barium, calcium
and lead)
Most chlorides
are soluble in
water (except
lead and silver).
PbCl2 is soluble
in hot water.
All nitrates are
soluble in water
Preparing
soluble salts
from acids
Preparing
insoluble salts
using
precipitation
reactions
Carrying out
acid-alkali
titrations
Preparing soluble salts from acids
Eg. the preparation of sodium chloride by neutralization
Preparing soluble salts from acids
Eg. the preparation of sodium chloride by neutralization
Dilute
sodium
hydroxide
+ indicator
solution
Dilute
hydrochloric
acid
Preparing soluble salts from acids
Eg. the preparation of sodium chloride by neutralization
Filtrate
Decolourising
charcoal
Evaporating
basin
Preparing soluble salts from acids
Eg. the preparation of sodium chloride by neutralization
Crystals of
sodium
chloride
forming)
Filtrate
Acid + Alkali 
Salt + Water
Steam
Water bath
Bunsen
burner
HCl(aq) + NaOH(aq) 
NaCl(aq) + Water(l)
Preparing insoluble salts using
precipitation reactions
Eg. the preparation of lead iodide
Preparing insoluble salts using
precipitation reactions
Eg. the preparation of lead iodide
Lead
nitrate
Potassium
iodide
Yellow
precipitate
of lead
iodide
Preparing insoluble salts using
precipitation reactions
Eg. the preparation of lead iodide
Filter to separate the
precipitate. Wash with
distilled water and dry to
get the pure product.
Lead
iodide
Preparing insoluble salts using
precipitation reactions
Eg. the preparation of lead iodide
Filter to separate the
precipitate. Wash with
distilled water and dry to
get the pure product.
Salt + Salt  Insoluble + Soluble
salt
salt
Lead
iodide
Lead + Potassium  Lead + Potassium
nitrate iodide
iodide
nitrate
Pb(NO3)2(aq) + 2KI(aq)  PbI(s) + 2KNO3(aq)
Acid – Alkali Titrations
A titration is a
very accurate way
of adding an acid
to an alkali to get
a salt.
Acid – Alkali Titrations
To carry out an acid-alkali titration we need
the right bits of kit.
Conical
flask
Acid – Alkali Titrations
To carry out an acid-alkali titration we need
the right bits of kit.
Pipette
Acid – Alkali Titrations
To carry out an acid-alkali titration we need
the right bits of kit.
Burette
Acid – Alkali Titrations
Stage 1
Dilute sodium hydroxide (NaOH)
solution is sucked up into a pipette
using a pipette filler. The pipette
contains exactly 25.0cm3 of
solution when the bottom of the
meniscus is level with the pipette
mark. The sodium hydroxide
solution is then released in to the
conical flask.
meniscus
Acid – Alkali Titrations
Stage 1
Dilute sodium hydroxide (NaOH)
solution is sucked up into a pipette
using a pipette filler. The pipette
contains exactly 25.0cm3 of
solution when the bottom of the
meniscus is level with the pipette
mark. The sodium hydroxide
solution is then released in to the
conical flask.
meniscus
Acid – Alkali Titrations
Stage 2
Two or three drops of an acidalkali indicator such as litmus
solution are added to the sodium
hydroxide solution using a teat
pipette. The alkali turns the
litmus blue.
Acid – Alkali Titrations
Stage 3
The burette is filled with dilute
hydrochloric acid to the zero
reading. (Again, look for the
meniscus)
Acid – Alkali Titrations
Stage 4
The conical flask is placed on a
white tile beneath the burette.
Acid is added from the burette
until the colour of the solution
turns from blue to red. The point
at which the colour changes is
called the end-point. At this
stage the conical flask only
contains salt and water.
Acid – Alkali Titrations
Stage 4
The conical flask is placed on a
white tile beneath the burette.
Acid is added from the burette
until the colour of the solution
turns from blue to red. The point
at which the colour changes is
called the end-point. At this
stage the conical flask only
contains salt and water.
At the end-point the volume of
acid added can be measured by
reading off the volume used in the
burette – this is the titre
Preparing
soluble salts
from acids
Preparing
insoluble salts
using
precipitation
reactions
Carrying out
acid-alkali
titrations
Preparing
soluble salts
from acids
Preparing
insoluble salts
using
precipitation
reactions
Carrying out
acid-alkali
titrations
Preparing
soluble salts
from acids
Preparing
insoluble salts
using
precipitation
reactions
Carrying out
acid-alkali
titrations
End of Section 4 Lesson 1
In this lesson we have covered:
pH scale and indicators
Common acids and alkalis
Salt formation and reactions
Preparing salts