Acids, Bases and Salts IGCSE

1. CHAPTER 5
ACIDS BASES
AND SALTS
Y E A R 1 0

2. YOU WILL FIND OUT ABOUT:
• Common acids – where and how they occur
• The pH scale and indicators **
• The colour changes of useful indicators **
• The ions present in acid and alkali solutions **
• The differences between acids, bases and alkalis
• The acid-base properties of non-metal oxides and metal oxides **
• Neutral and amphoteric oxides **
• Uses of common alkalis, bases and ‘antacids’, including indigestion
treatments, the treatment of acid soils and waste water treatment
• The characteristic reactions of acid
• Acids and alkalis in the analysis of salts

3. YOU WILL FIND OUT ABOUT:
• The nature and solubility of salts
• The preparations of soluble salts by various methods including titration
• The preparation of insoluble salts by precipitation **
• The nature of strong and weak acids and alkalis **
• Acids as proton donors and bases as proton acceptors **

4. 5.1 WHAT IS AN ACID?
The major acids
• Word ‘acid’ was originally applied to
substances with sour taste.
• Acids present in animal and plant
materials are known as organic acids.
What are indicators?
• Certain coloured substances
(many extracted from plants)
have been found to change
colour if added to acid solutions.
• The colour change is reversed if
the acid is ‘neutralised’ or turn to
alkali.
• Coloured extract can be made
from red cabbage or
blackberries.
• The most used indicator
historically is litmus – extracted
from lichens.

5. TYPE NAME STRONG OR
WEAK
WHERE FOUND OR USED
Organic
acids
Ethanoic
acid
Weak In vinegar
Methanoic
acid
Weak In ant and nettle stings; used in kettle
descaler
Lactic acid Weak In sour milk
Citric acid Weak In lemons, oranges and other citrus fruits
Mineral
acids
Carbonic
acid
Weak In fizzy soft drinks
Hydrochlori
c acid
Strong Used in cleaning metal surfaces; found as
the dilute acid in stomach
Nitric acid Strong Used in making fertilisers and explosives
Sulfuric acid Strong In car batteries; used in making fertilisers,
paints and detergents
Phosphoric Strong In anti-rust paint; used in making

6. • Litmus paper comes from paper that has been soaked in litmus solution.
It comes in blue and red forms. The litmus will just gives a single colour
change.
Thymolphthalein Colourless Colourless Blue
• In the table are commonly used indicators including litmus. They are
used because sometimes, the colour changes are easier to ‘see’ than
that of litmus.
Indicator Colour in
acid
Neutral
colour
Colour in
alkali
Litmus Red Purple Blue
Phenolphthalei
n
Colourless Colourless Pink
Methyl orange Red Orange Yellow

7. Universal indicator
• Another commonly used indicator, Universal Indicator (or full-range
indicator), is a mixture of indicator dyes.
• The idea of it is to imitate the colours of rainbow when measuring
acidity – this is useful because it gives a range of colours depending on
the strength of acid and alkali added.

8. The pH scale
• The pH scale to measure the strength of acid solution was worked out
by Danish biochemist Soren Sorensen. He worked in the labs of
Carlsberg breweries and was interested in checking the acidity of beer.
Rules for pH scale
• Acids have pH less than 7
• Alkalis have pH more than 7
• Neutral substances have pH of 7
• The more acidic a solution the lower the pH, and the more alkaline a
solution the higher the pH.
• Other than Universal Indicator paper or solution, pH meter is the most
accurate method to measure the pH range of substances – uses an
electrode to measure pH electrically.

9. ACID AND ALKALI SOLUTIONS
The importance of hydrogen ions
• All acids contain hydrogen.
• Acids can conduct electricity and they conduct electricity better than
distilled water. This show that acid solution contains ions.
• Water itself contains ions (H+ and OH-). All acids dissolve in water to
produce hydrogen ions and all alkalis dissolve in water to produce
hydroxide ions.
• The term pH is taken from the German ‘potenz H(ydrogen)’ which means
means the power of the hydrogen ion concentration of a solution.
• An indicator is affected by the presence of H+ and OH- ions

10. • The hydrogen ions in acid solutions make litmus go red.
• The hydroxide ions in alkali solutions make litmus go blue.
The importance of water
• Gas hydrogen chloride is made up of covalently bonded molecules. If
the gas dissolves in water, a strongly acidic solution is produced.
• Acid is substance that dissolves in water to produce hydrogen ions.
• Alkali is substance that dissolves in water to produce hydroxide ions.
• Large volume of water will produce dilute acid and alkali solutions and
less volume of water will produce concentrated acid and alkali solutions.

11. METAL OXIDES AND NON-METAL OXIDES
Acidic and basic oxides
The characteristics of oxides:
• Non-metals generally form acidic oxides that dissolve in water to form
acidic solutions.
• Metals form oxides that are solids. If they dissolve in water, these oxides
give alkaline solutions.These metal oxides neutralise acids and are basic
basic oxides.

12. Element How it reacts Product Effect of adding water and
testing with litmus
Non-metals
Sulfur Burns with bright blue flame Colourless gas (sulphur
dioxide)
Dissolves, turns litmus red
Phosphorus Burns with yellow flame White solid (phosphorus
(V) oxide)
Dissolves, turns litmus red
Carbon Glows red Colourless gas (carbon
dioxide)
Dissolves slightly, slowly turns
litmus red
Metals
Sodium Burns with yellow flame White solid (sodium
oxide)
Dissolves, turns litmus blue
Magnesium Burns with bright white flame White solid (magnesium
oxide)
Dissolves slightly, turns litmus
blue
Calcium Burns with red flame White solid (calcium
oxide)
Dissolves, turns litmus blue
Iron Burns with yellow sparks Blue-black solid (iron
oxide)
Insoluble
Copper Does not burn, turns black Black solid (copper Insoluble

13. Neutral and amphoteric oxides
• Water can be thought as hydrogen oxide with pH 7 and is called as
neutral oxide – do not react with either acids or alkalis. This is an
exception to the ‘rule’ that non-metals oxides are acidic oxides.
• Some metals oxides like zinc and aluminium have amphoteric
compounds that’ll produce amphoteric metal oxides (also known as
amphoteric hydroxides) – a metal oxide that can react with both acid
and alkali to give salt and water.

14. Amphoteric hydroxide example
• Sodium hydroxide is used to identify the salts produced by these
amphoteric metal oxides – zinc hydroxides and aluminium hydroxides.

15. 5.4 ACID REACTIONS IN EVERYDAY
LIFE
Indigestion, headaches and neutralisation
• Dilute hydrochloric acid in our stomach is to help digest our food.
However, excess acid causes indigestion which eventually give rise to
ulcers. To ease this, we can take antacids.
• Antacids are group of compounds with no toxic effects on the body.
• Some of the antacids contain insoluble materials to counteract the acid
or to neutralise them. For example ‘milk of magnesia’ contains insoluble
sodium hydroxide.
• Some of them, like Alka-Seltzer contain soluble material including
sodium hydrogencarbonate. On adding water, it’ll produce carbon
dioxide gas and dissolve other less soluble material.

16. Descaling kettles
• Lime-scale collects inside kettles and water heaters in hard-water areas.
Hard-water areas tend to be geographically located in limestone areas.
It contain more dissolved calcium ions than normal water.
• The limescale can be removed by treatment with an acid that is strong
enough to react with calcium carbonate but not strong enough to
damage the metal.
• Vinegar are usually used to descale kettles. Commercial ‘descalers’ use
other acid solutions such as methanoic acid.

17. Soil pH and plant growth
• Plant growth is affected by the acidity and alkalinity of the soil. Acidic
soils (soils with high peat content, or with minerals, or with rotting
vegetation and lack of oxygen) can reach as low as pH 4 while alkaline
soils (soils in limestone or chalky areas) can reach as high as pH 8.3.
• Different plants prefer different pH conditions.
• If the soil is too acidic, it is usually treated by ‘liming’ – either use
calcium oxide, calcium hydroxide or powdered chalk or limestone. They
all can neutralise the acidity of soil.
• If the soil is too alkaline, it helps to dig in some peat or decaying organic
organic matter.
• Some flowering plants carry their ‘built-in’ pH indicator. The flowers of
hydrangea bush are blue in acidic soil and are pink in alkaline soil.

18. Effluent (liquid waste and sewage) and waste water treatment
• Liquid waste from factories is often acidic and if they get into the river,
the habitats living there will be killed.
• Slaked lime is often added to neutralise it.
• The same method is used to treat streams, rivers and lake affected by
acid rain.
• To reduce emission of sulphur dioxide, many modern factories now
spray acidic waste gases with jets of slaked lime in a flue-gas
desulfuriser to neutralise them before they leave the chimneys.

19. 5.5 ALKALIS AND BASES
What types of substance are alkalis and bases
• Alkalis are substance that dissolve in water to give solutions with pH
more than 7 and turn litmus blue. The solutions contain excess of
hydroxide ions.
• However, among the antacids we use is insoluble magnesium hydroxide
which shows that all metal oxides and hydroxides can also neutralise
acids whether they dissolve in water or not.
• These substances are known as bases.
• Bases react with acid to produce salt and water.
• Most bases are insoluble in water. Thus, the few bases that dissolve in
water are known as alkalis.

20. Common alkalis are:
• Sodium hydroxide solution
• Potassium hydroxide solution
• Calcium hydroxide solution
• Ammonia solution (also known as ammonium hydroxide)
The first 2 are stronger alkalis than others.

21. Properties and uses of alkalis and bases
Bases:
Neutralise acids to give salt and water only
Are the oxides and hydroxides of metals
Are mainly insoluble in water
Alkalis are bases that dissolve in water:
Feel soapy to skin
Turn litmus blue
Give solutions with pH greater than 7
Give solutions that contain hydroxide ions

22. 5.6 CHARACTERISTIC REACTIONS OF
ACIDS
The reactions of acids
• A reactive metal (for example, magnesium or zinc)
• A base (or alkali) – neutralisation reaction
• A metal carbonate (or metal hydrogencarbonate)
All these reactions will produce salt – a compound made from an acid
when a metal displace the hydrogen in the acid. The acid from which the
salt is made is called parent acid.

23. The reaction of acids with metals
metal + acid salt + hydrogen
This reaction can only be done using the quite reactive metals, not the
very reactive ones like sodium or calcium – the reaction is too violent.
The salt made depends on the acid:
• Hydrochloric acid – chloride HCl
• Nitric acid – nitrate HNO3
• Sulfuric acid – sulphate H2SO4
• Ethanoic acid – ethanoate CH3COOH
For example: magnesium + nitric acid magnesium nitrate +
hydrogen
calcium + hydrochloric acid ------- calcium chloride +

24. The reaction of acids with bases and alkalis
acid + base/alkali salt + water
The salt produced will still depend on the acid used. To make particular
salt, choose a suitable acid and base to give out the salt.
Base = metal oxide
Alkali = metal hydroxide
For example:
sodium hydroxide + hydrochloric acid sodium chloride + water
More examples of salts produced by bases and acids in page 133 (Table
5.8)

25. The reaction of acids with carbonates
acid + metal carbonate salt + water + carbon dioxide
The normal method of preparing carbon dioxide in the laboratory is based on this
reaction. Dilute hydrochloric acid is reacted with marble chips (calcium carbonate).
For example:
hydrochloric acid + calcium carbonate calcium chloride + water + carbon
dioxide

26. 5.7 ACIDS AND ALKALIS IN CHEMICAL
ANALYSIS
The test for carbonates using acid
All carbonates will react with acids to give off carbon dioxide. We can use
this as a test to find out if an unknown substance is a carbonate or not:
A piece of rock that we think is limestone can be checked by dripping a
few drops of vinegar on it. If it fizzes, then it could be limestone.
Add dilute hydrochloric acid to a powdered substance and any gas
given off would be passed into lime water. If it turns cloudy, the gas is
carbon dioxide and the substance is carbonate

27. Test for metal ions in salts using alkalis
All salts are ionic compounds made up of metal ions and non-metal ions.
We have seen that most metal hydroxides are insoluble. Thus, we can
identify the metal ions of the unknown salt by adding an alkali to the
solution.
Coloured hydroxide precipitates
Copper(II) salts give a light blue precipitate of copper(II) hydroxide.
Iron(II) salts give a light green precipitate of iron(II) hydroxide.
Iron(III) salts give a red-brown precipitate of iron(III) hydroxide.
Chromium(III) salts give a grey-green precipitate of chromium(III)
hydroxide.
For example:
iron(II) sulphate + sodium hydroxide iron(II) hydroxide + sodium

28. White hydroxide precipitates
Certain hydroxide precipitates are white – calcium, zinc and aluminium
hyroxides.
zinc sulphate + sodium hydroxide zinc hydroxide + sodium sulfate
ZnSO4(aq) + 2NaOH(aq) Zn(OH)2(s) + Na2SO4(aq)
Even though the precipitates are all white, the test is still useful:
When an excess of sodium hydroxide added, only calcium hydroxide
does not re-dissolve, both zinc and aluminium hydroxide will re-dissolve.
dissolve.
To check between aluminium and zinc ion, the test needs to be repeated
repeated with ammonia solution. The same white precipitates of their
hydroxides produced.
However, with excess ammonia solution it is only the zinc hydroxide that

29. Test for ammonium salts using alkali
Ammonium salts are important as fertilisers like ammonium sulphate and
ammonium nitrate. They are salts containing ammonium ions that produce
ammonia gas when reacted with alkali solutions. They can be detected because
they turn damp red litmus paper blue.
Example:
ammonium nitrate + sodium hydroxide sodium nitrate + water +
ammonia
This reaction occurs because ammonia is more volatile base than sodium
hydroxide, so it is easily displaced by sodium hydroxide from its salt. This
reaction can be used to determine ammonium ions or to prepare ammonia in
lab.

31. 5.8 SALTS
The importance of salts – an introduction
• A salt is an ionic compound formed from an acid by the replacement of
the hydrogen in the acid by a metal.
• While a number of salts can be obtained by mining, others must be
made by industry. Therefore, it is worth considering the methods
available to make salts.
• Two things are important in working out method of salt preparation:
– Is the salt soluble or insoluble in water?
– Do crystals of the salt contain water of crystallisation?

32. The solubility of salts
Soluble salts are made by neutralising acids. Insoluble salts are made by
other methods.
Salts Soluble Insoluble
Sodium salts ALL NONE
Potassium salts ALL NONE
Ammonium salts ALL NONE
Nitrates ALL NONE
Ethanoates ALL NONE
Chlorides Most are soluble Silver chloride,
Lead(II) chloride
Sulfates Most are soluble Barium sulphate,
Lead(II) sulphate,
Calcium sulphate
Carbonates Sodium, potassium
and ammonium
carbonates
Most are insoluble

33. Water of crystallisation
The crystals of some salts contain water of crystallisation. This water gives the
crystals their shape. In some cases it also gives them their colour. Such salts are
known as hydrated salts.
When these hydrated salts are heated, their water of crystallization is driven off
as steam. The crystals lose their shape and become powder.
Example: copper(II) sulphate crystals are blue, but when we heat them, they
dehydrated to form white powder:
copper(II) sulphate crystals anhydrous copper(II) sulphate + water vapour
CuSO4.5H2O(s) CuSO4(s) + 5H2O(g)

34. Crystals that have lost their water of crystallisation are called anhydrous.
If water is added back, the powder will turns to their colour again and
heat is given out. This can also be used to test the presence of water.
Hydrated salt Colour
Copper(II) sulphate Blue
Cobalt(II) chloride Pink
Iron(II) sulphate Green
Magnesium sulphate White
Sodium carbonate White
Calcium sulphate white

35. 5.9 PREPARING SOLUBLE SALTS
Method A – Acid plus solid metal, base or carbonate
• Stage 1: an excess of the solid is added to the acid and allowed to react
and make sure that all the acid is used up.
• Stage 2: the excess solid is filtered out.
• Stage 3: the filtrate is gently evaporated to concentrate the salt solution.
This can be done on a heated water bath or sand tray.
• Stage 4: when crystals can be seen forming, heating is stopped and the
solution is left to crystallise.
• Stage 5: the concentrated solution is cooled to let the crystals form.
 Refer figures in page 139

37. Method B – Acid plus alkali by titration
• Stage 1: the acid solution is poured into a burette. A known volume of
alkali solution is placed in a conical flask using a pipette. A few drops of
an indicator (universal indicator, methyl orange, etc.) are added to the
flask.
• Stage 2: the acid solution is run into the flask from the burette until the
indicator just changes colour. When the end-point is found, the volume
of acid run into the flask is noted. The experiment is repeated without
indicator and will use the volume of acid as noted.
Alternatively, charcoal can be added to remove the coloured indicator and
it can be filtered off leaving the colourless solution behind.
• Stage 3: the salt solution is evaporated and cooled to form crystals as
described in Method A.

38. Indicator Colour in
acid
Neutral
colour
Colour in
alkali
Universal
Indicator
Red Green Blue
Phenolphthalei
n
Colourless Colourless Pink

39. 5.10 PREPARING INSOLUBLE SALTS
Choosing a method of salt preparation
Start
Does the metal react
with acids
Does it react
safely?
Is the base or carbonate
soluble?
Method A: can
prepare salt by
reacting acid with
excess solid, followed
by filtration, e.g.
CuSO4
Method A: can
prepare salt by using
excess metal and
acid, followed by
filtration, e.g. MgSO4
Method B: can use
titration method,
e.g. NaCl
Salt crystals
prepared by
evaporation
and
crystallisatio
n
Yes
No
Yes
No
No
Yes

40. Making salts by precipitation (usually used for chlorides, iodides and
sulfates salts)
• Some salts are insoluble in water. Such salts cannot be made by
crystallisation methods.
• They are generally made by ionic precipitation – the sudden formation
of a solid, either when two solutions are mixed or when a gas is bubbled
into a solution.
• For example, barium sulphate is made when a solution of soluble
sulphate (sodium sulphate) is added to a solution of soluble barium salt
(barium chloride). The insoluble barium sulphate is formed immediately.
• The precipitate will ‘fall’ to the bottom of the beaker and then filtered
off. Then, we wash with distilled water and dried in warm oven.
BaCl2(aq) + Na2SO4(aq) BaSO4(s) + 2NaCl(aq)
Ba2+(aq) + SO4
2-(aq) BaSO4(s) (ionic equation)
• The ions that remain in the solution are known as spectator ions.

42. 5.11 STRONG AND WEAK ACIDS AND
ALKALIS
Strong and weak acids
• The strength of an acid depends on its concentration of hydrogen ions
(H+).
• The higher the concentration of hydrogen ions, the higher the acidity
and the lower the pH.
• Each pH unit means a ten-fold difference in H+ ion concentration. An
acid of Ph 1 has ten times the H+ ion of an acid of pH 2.
• Strong acids are completely ionised (complete separation into ions) in
solution in water. For example, hydrochloric acid and sulphuric acid.
• Weak acids are partially dissociated into ions in solution in water. For
example, ethanoic acid, carbonic acid and methanoic acid.

43. Strong and weak alkalis
• Alkalis are also the same as acids except that they are classified based
on the concentration of hydroxide ions (OH-).
• The higher the concentration of hydroxide ions, the higher the alkalinity
and the higher the pH.
• Strong alkalis like sodium hydroxide and potassium hydroxide are
completely ionised in water.
• Weak alkalis like ammonia solution only partially dissociated in water,
thus it only have a low concentration of ammonium ions and hydroxide
ions.

44. • High concentration of ions make the strong acids and alkalis conduct
electricity well. They can be used as strong electrolytes in solution.
• While, the weak acids and alkalis are weak electrolytes.
• For weak acids and alkalis, the dissociation process is reversible: it can
go in either direction.
• For example, in ethanoic acid, molecules are dissociating into ions, but
the ethanoate ions and hydrogen ions are also re-combining to be
ethanoic acid compound back.
CH3COOH(aq) H+(aq) + CH3COO-(aq)
H2O

45. What happens to the ions in neutralisation?
acid + alkali salt + water
For example,
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
• All these compounds are completely ionised, except for water produced.
• The hydrogen ions from the acid and hydroxide ions from the alkali have
combined to form water.
H+(aq) + OH-(aq) H2O(l)
• This is the ionic equation in the neutralisation. The spectator ions
(sodium and chloride) remains in the solution and become a solution of
sodium chloride.
• Acid is a molecule or ion that is able to donate a proton (H+ ion) to a
base.

46. The ‘basicity’ of acids
• Monobasic acid – has one replaceable hydrogen atom per molecule and
produce only one salt.
• Dibasic acid – has two replaceable hydrogen atoms per molecule and
produce two different salts.
• Tribasic acid – has three replaceable hydrogen atoms per molecule and
produce three different salts.
Refer table 5.13 on page 147 for the examples of ‘basicity’ of common
acids.