SALTS

Definition of a salt
A salt is a chemical substance formed when the hydrogen ions in an acid are replaced by a metal or ammonium
ions. The replacement of hydrogen ions can be complete or partial.

For example if we replace the H+ in HCl with a potassium atom, then the salt potassium chloride is formed, KCl.
Salts are an important branch of chemistry due to the varied and important uses of this class of compounds.
These uses include fertilisers, batteries, cleaning products, healthcare products and fungicides.
TYPES OF SALTS
1. NORMAL SALTS / NEUTRAL SALTS
This is a salt formed when all the hydrogen ions in an acid are replaced by a metal or ammonium ions.
Examples of how normal or neutral salts are formed.
i. NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)
ii. Ca (s) + H2SO 4 (aq) → CaSO4 (aq) + H2 (g)
iii. 2NaOH (aq) + H2SO4 (aq) → Na2SO4 (aq) + 2H2O (l)
2. ACIDIC SALTS
This is a salt formed when part of the hydrogen ions in an acid are replaced by a metal or ammonium ions.
Partial replacement of hydrogen ions in the acid takes place in order to form an acidic salt.
Acidic salts are formed usually by a strong acid and a weak base when they react together. For example, NH4Cl
is made from a strong acid hydrochloric acid and a weak base ammonium hydroxide.
Examples of how acidic salts are formed.
i. KOH (aq) + H2SO4 (aq) → KHSO4 (aq) + H2O (l)
ii. NH4OH (aq) + H2SO4 (aq) → NH4 HSO4 (aq) + H2O (l)
iii. NaOH (aq) + H2SO4 (aq) → NaHSO4 (aq) + H2O (l)
iv. CaO (s) + 2H2CO3 (aq) → Ca (HCO3)2 (aq) + H2O (l)
Names of Common Acidic Salts

 Potassium hydrogen sulphate, KHSO4

 Sodium hydrogen sulphate, NaHSO4
 Calcium hydrogen carbonate, Ca(HCO3)2
 Sodium hydrogen carbonate, NaHCO3 (Baking/Bicarbonate salt)
3. BASIC SALT
This is a salt made from the reaction between a weak acid and a strong base. These are salts that contain
hydroxide ions (OH-) in their compounds. Example of basic salt is Copper (ii) carbonate, CuCO3.Cu (OH)2
Naming Salts
The name of a salt has two parts;
i. The first part comes from the metal part of a base, alkali, metal, or metal carbonate used in the
reaction.
ii. The second part comes from the negative ion (anion) present in acid used. For example;
Hydrochloric acid produces chloride salts. Nitric acid produces nitrate salts etc.

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The name of the salt can be determined by looking at the reactants. For instance, hydrochloric acid always
produces salts that end in chloride and contain the chloride ion, Cl –. Sodium hydroxide reacts with hydrochloric
acid to produce sodium chloride. Zinc oxide reacts with sulfuric acid to produce zinc sulphate.
SOLUBILITY OF SALTS
What is solubility? Solubility is the ability of a salt to dissolve in water.
When preparing salts, the first thing to know is the solubility of the salt you want to prepare. That is to say, Is the
salt being formed soluble or insoluble in water? The table below gives the information about solubility of different
salts. You need to memorize and know this table as it will not be provided in the exam or a test.
SOLUBILITY RULES FOR COMMON SALTS

SALTS SOLUBLE INSOLUBLE

1 Salts containing Sodium, Potassium, All None
Lithium, Ammonium
2 NITRATES All None

3 CHLORIDES Most are soluble Silver chloride

Lead (ii) chloride
4 SULPHATES Most are soluble Calcium sulphate,
Barium sulphate,
Lead (ii) sulphate
5 CARBONATES Sodium carbonate, Potassium carbonate Most Are Insoluble
Lithium carbonate, Ammonium carbonate

7 HYDROXIDES Sodium hydroxide, potassium hydroxide, Most Are Insoluble

Lithium hydroxide, Ammonium hydroxide

8 OXIDES Sodium oxide, Potassium oxide, Most Are Insoluble

Lithium oxide, Ammonium oxide

The following statements are connected to the solubility table shown above.
1) NITRATES
All nitrates are soluble.
2) CHLORIDES
All chlorides are soluble except silver chloride and lead (II) chloride which are insoluble sulphates
3) SULPHATES
All Sulphates are soluble except barium sulphate and lead (II) sulphate which are insoluble. Calcium sulphate is
slightly soluble.
4) CARBONATES
All carbonates are insoluble except potassium carbonate, sodium carbonate and ammonium carbonate which
are soluble.

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 METHODS USED TO PREPARE SALTS
There are 3 methods used to prepare salts, namely titration, neutralisation and precipitation.
 TITRATION METHOD – used to prepare soluble salts which contains a metal from group 1 or
ammonium salts
E.g. sodium chloride, potassium nitrate, ammonium sulphate
 NEUTRALISATION METHOD - used to prepare soluble salt which has a metal not from group 1
E.g. copper (ii) sulphate, zin chloride, calcium nitrate
 PRECIPITATION METHOD is used to prepare insoluble salts.
Remember there are only 5 insoluble salts, which are BaSO4, CaSO4 and PbSO4 under sulphates and
Silver chloride (AgCl) and Lead (ii) chloride (PbCl2) under chlorides.

1) NEUTRALISATION METHOD
This method is used to prepare any soluble salt that does not contain a group 1 metal or ammonium ion.
This means that if the metal part of the salt is from group 2 or 3 or a transition metal then you use this
method.
Reagents or chemicals used
 An acid
 Insoluble substance – there are 3 options to choose from under this.
i. A metal oxide
ii. A carbonate
iii. A metal in powder form.
(Note: For copper, a copper metal can’t be used because copper is below hydrogen in the
reactivity series)

 Take note that NO INDICATOR IS USED during neutralisation method.

Steps To Take In Doing Neutralisation Method
i) Put dilute acid in a beaker, warm the acid but don’t boil. The acid is warmed in order to speed up
the reaction/ increase the rate of the reaction.
ii) Add an excess the insoluble substance to the acid in the beaker and stir until it dissolves
completely. Excess of the insoluble substance is added to ensure that all the acid is completely
neutralised.
iii) Filter off excess metal oxide/carbonate/ metal powder to get a salt solution.
iv) Evaporate the filtrate/ salt solution to obtain a saturated solution
v) Cool the salt solution to allow crystals of to form.
A summary of all these steps is illustrated by the diagrams below.

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2) TITRATION METHOD
It is used to prepare salts which have a group 1 metal or ammonium ion. Examples of such salts are
NaCl, KNO3, LiCl, NH4Cl, Na2SO4 etc
Reagents or chemical used for titration
 An acid
 An alkali
Steps To Do To Prepare A Salt
i. Using a pipette, measure 25 cm3 of an alkali and put it into a conical flask.
ii. Add two or three drops of indicator to sodium hydroxide using a dropper
iii. Fill the burette to the zero reading with an acid.
iv. Place the conical flask on a white tile below the burette
v. Add acid from the burette to alkali in the conical flask, while swirling, until the mixture just
changes colour

vi. From the titration result, we can know the exact volume of acid needed to react with 25.0cm3
of the alkali. Volume of acid used is calculate by doing this: V = V2 – V1
Note: The indicator added contaminates the salt (it’s an impurity). So a second titration is done
without adding an indicator.
Steps in doing a second titration is carried out without the indicator
The exact volume obtained from the preliminary or first titration is used.
i. Using a pipette, measure of an alkali and put it into a conical flask. This time no indicator is
added.
ii. Add an acid from the burette to an alkali.
iii. Evaporate the mixture to obtain a saturated solution.
iv. Cool the saturated solution to obtain crystals of the salt

3) PRECIPITATION METHOD
Precipitation is the formation of an insoluble product and may occur on mixing two solutions.
Precipitation is an example of double decomposition. In double decomposition, two solutions are mixed
to form an insoluble salt and a soluble solution.
Soluble salt + soluble salt → insoluble salt + soluble salt
In double decomposition reactions, cations and anions are exchanged.

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 Precipitation is also an example of ionic association which is the attraction of oppositely charged ions to
one another to form a solid called precipitate abbreviated as ppt.
Preparation of insoluble salts by precipitation (Double decomposition)
Example: Preparation of silver chloride
Reagents
 Silver nitrate, AgNO3
 Sodium chloride, NaCl
Reaction equation
AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq)

Method of Preparation
i. Mix silver nitrate solution with sodium chloride solution in a beaker or test tube
ii. A white precipitate of silver chloride forms.
iii. Allow the precipitate to settle.
The figure below shows what is obtained when two salts are mixed and the resulting mixture is
allowed to settle to obtain a precipitate

Sodium nitrate solution

Silver chloride precipitate

iv. Filter off the precipitate and wash it with distilled water to remove any of sodium nitrate left.
v. Dry the precipitate on the filter paper. Pure dry sample of silver chloride forms.

BEHAVIOUR OF SOME SALTS WITH WATER AND AIR

Many salts that are soluble tend to absorb water vapour from the air. They are said to be hygroscopic. You will
have noticed that the table salt at home in a salt shaker becomes damp and tends to stick together in humid
weather. This shows that the salt is a hygroscopic material. Other hygroscopic salts include potassium carbonate
(K2CO3), anhydrous calcium chloride (CaCl2) and anhydrous sodium sulphate (Na2SO4).

HYGROSCOPIC SUBSTANCE. A hygroscopic substance absorbs water from the air but does not the change its
state. Anhydrous cobalt chloride is a hygroscopic salt. Water changes anhydrous cobalt chloride from blue to pink.
CoCl2 (s) + H2O (l) → CoCl2 .6H2O (s) + heat

(Blue) (Pink) \

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This reaction is often used as a test for the presence of water. The process can be reversed by heating the pink
hydrated salt.

Concentrated sulphuric acid is also hygroscopic. It can be used to dehydrate or to remove water from blue
crystals of hydrated copper (ii) sulphate forming the pale blue anhydrous salt.

DELIQUESCENCE SUBSTANCE

These are substances that absorb water from the atmosphere to form a solution. Calcium chloride is a
deliquescent salt. It is used as a drying agent in desiccators. A desiccator is a piece of equipment used to dry
substances.

HYDRATED SALT
Solids salts usually form crystals when the cation and anion combine. Sometimes some water gets trapped
within the structure of the crystal. This trapped water is called the water of crystallization or water of hydration.
A hydrated salt is a salt that contains water of crystallization in its crystals. They contain a fixed amount of water
in their crystal lattice. The water of crystallization is part of the structure. If this water is removed, by heating for
example, the colour and shapes of the crystals often changes. Examples of hydrated salts are in the table below
Name of crystallized salt Chemical Formula
1 Copper (ii) sulphate -5 water CuSO4.5H2O
2 Sodium carbonate -10 water Na2CO3.10H2O
3 Cobalt (ii) chloride - 6 water CoCl2.6H2O
4 Iron (ii) sulphate - 7 water FeSO4.7H2O

The water of crystallization is shown in the formula by writing the number of water molecules in each salt crystal,
after the formula of the salt. For example, Tin chloride – 2- water or hydrated tin (ii) chloride is written as
SnCl2.2H2O.
ANHYDROUS SALTS
A salt which has lost its water of crystallization is called an anhydrous salt.
When water is added to an anhydrous salt, the salt becomes hydrated. For example, when blue copper (II)
sulphate crystals are heated, stem is produced and a pale- blue or white powder.
CuSO4.5H2O (s) CuSO4(s) + 5H2O (g)

(Hydrated copper sulphate) (Anhydrous copper sulphate) (Steam)

When water is added to anhydrous copper (ii) sulphate heat is produced and a blue solution is formed:

CuSO4 (s) + 5H2O (l) → CuSO4.5H2O (aq) + heat

This process is called hydration

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EFFLORESCENT SALT
Efflorescence is the loss of water of crystallization to the atmosphere by a salt. Hence salts that lose water of
crystallization to the atmosphere when exposed to the air as efflorescent salts.
For example, crystals of sodium carbonate – 10 − water turn into powder when exposed to air. This happens
because they lose the water of crystallization to the atmosphere so easily. This process is shown by the equation
below: Na2CO3.10H2O (s) → Na2CO3.H2O (s) + 9H2O (l)

Subtopic: OXIDES
What are oxides? Oxides are compounds containing oxygen and another element (a metal or non-metal)
You have seen already that metal oxides act as bases. Here we look more closely at different types of oxides,
and their behaviour
TYPES OF OXIDES
1. BASIC OXIDES – these are oxides of metals.
Metallic oxides are known as basic oxides because they are bases in nature because they neutralise
acids to form salt and water only. Basic oxides belong to the larger group of compounds called bases.
Look at the following situation. When copper is heated in a stream of the oxygen gas, its surface turns
black and a black substance, copper (ii) oxide, is formed as shown:
Cu(s) + O2 (g) → CuO.
Dilute HCl acid turns blue litmus paper red, like all acids do. Copper (ii) oxide reacts with HCl acid, when
it is warmed. The resulting liquid or product in the flask has no effect on blue litmus, the colour of the
indicator stays the same. This means the copper (ii) oxide has neutralised the HCl acid, hence it is
acted like a base.
Copper (ii) oxide is called a basic oxide since it can neutralise an acid as shown below;
CuO (s) + 2HCl (aq) → CuCl2 (aq) + H2O (l)
*In general, metals react with oxygen to form basic oxides

2. ACIDIC OXIDES – Acidic oxides are oxides of non-metals. They are formed when non-metals react with
oxygen.
For instance, when carbon is heated over a Bunsen burner until red-hot, then plunged into a jar of
oxygen. It glows bright red, and the gas carbon dioxide is formed according to the equation below:
C (s) + O2 (g) → CO2 (g)
Carbon dioxide is slightly soluble in water. The solution will turn litmus red: it is acidic. The weak acid
carbonic acid has formed: CO2 (g) + H2O (l) → H2 CO3 (aq)
Sulphur dioxide and phosphorus pentoxide also dissolve in water to form acids. So they are all called
acidic oxides. In general, non-metals react with oxygen to form acidic oxides.

3. NEUTRAL OXIDES
Neutral oxides are oxides of some non-metals. These oxides have neither acidic nor basic properties so
they do not form salts when they are mixed with acids or bases.
Examples of neutral oxides are water (H2O), carbon monoxide (CO) and dinitrogen oxide (N2O). (*Note
that other oxides of nitrogen such as nitrogen dioxide are acidic)

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 4. AMPHOTERIC ACID – These are some metallic oxides that show properties of acids and bases. They
can react with either an acid or a base.
For instance, you would expect aluminium oxide to be a base. In fact it is both acidic and basic. It acts
as a base with hydrochloric acid as shown:
Al2O3 (s) + HCl (aq) → AlCl3 (aq) + H2O (l)
But it acts as an acidic oxide with sodium hydroxide, giving a compound called sodium aluminate, a salt,
as shown: Al2O3 (s) + NaOH (aq) → Na3AlO3 (aq) + H2O (l)
From these two reaction results aluminium oxide is called an amphoteric oxide. An amphoteric oxide will
react with both acids and alkalis.
Other examples of amphoteric oxides are Zinc oxide and Lead (ii) oxide. They will also react with both
acids and alkalis
QUALITATIVE ANALYSIS TESTS (or IDENTIFICATION OF IONS)
Qualitative analysis is the analysis of chemical substances using their colour, smell, melting point boiling point or
solubility. It involves the identification of the solids, liquids, gases or ions produced in a chemical reaction.
TEST FOR CATIONS (or METAL IONS) IN SOLUTIONS
To test for a cation or metal ion we use two chemicals which are sodium hydroxide and ammonia solution. The
following are the steps on how the tests are done;
1. Take a small portion (about 2mℓ or 2cm3) of the sample/ salt solution and put it in a clean test tube.
Using a dropper, add few drops of sodium hydroxide. Record the colour of the precipitate formed.
To the same mixture you had in step 1, add excess or a lot of sodium hydroxide solution. Record what
happens to the precipitate, whether it dissolves (soluble) or does not dissolve (Insoluble).
2. Take another clean test tube, add a small portion of the sample/ salt solution. Using a dropper add a
few drops of ammonia solution to the sample solution. Record the colour of the precipitate that forms.
Add excess of ammonia solution to the same test tube containing the mixture.
The table below summaries the different tests for some common cation or metal ions using sodium hydroxide
solution and ammonia solution (or aqueous ammonia)
Cation ion present in Effect of aqueous sodium hydroxide Effect of aqueous ammonia when
the salt solution when added to the salt solution added to salt solution
1. Aluminium ion, Aℓ3+  White ppt is formed.  White ppt is formed.
 The ppt dissolves or ppt soluble  The ppt is insoluble in excess
in excess giving a colourless
solution

2. Calcium ion, Ca2+  White ppt is formed.  No is ppt is formed or very

 The ppt is insoluble in excess slight ppt is seen

3. Copper (ii) ion, Cu2+  Blue ppt is formed.  Blue ppt is formed.
 The ppt is insoluble in excess  The ppt is soluble in excess
(it does not dissolve in excess) giving a dark blue solution

4. Iron (ii) ion, Fe2+  A green ppt is formed  A green ppt is formed
 The ppt is insoluble in excess  The ppt is insoluble in excess

5. Iron (iii), Fe3+  A red-brown ppt is formed  A red-brown ppt is formed

 The ppt is insoluble in excess  The ppt is insoluble in excess
6. Zinc ion, Zn2+  A white ppt is formed  A white ppt is formed
 The ppt is soluble (or dissolves)  The ppt is soluble in excess
in excess

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 7. Ammonium ion, NH4+  Ammonia is produced after  No reaction occurs
warming the mixture of the
sample/ salt solution with a
sodium hydroxide. A damp/
moist red litmus paper
Ammonia gas turns damp (or
wet) red litmus paper blue

Take Note Of The Following Concerning Test For Cations:

The cations react with hydroxide ions present in aqueous sodium hydroxide or ammonia to form insoluble
hydroxides. These hydroxides appear as precipitates. The following ionic equations show how these precipitates
are formed.

i. Al3+ + OH- → Al (OH) 3 (a white ppt)

ii. Ca2+ + OH- → Ca (OH) 2 (a white ppt)
iii. Cu2+ + OH- → Cu (OH) 2 (a white ppt)
iv. Fe2+ + OH- → Fe (OH) 2 (a green ppt)
v. Fe3+ + OH- → Fe (OH) 3 (a red-brown ppt)
vi. Zn2+ + OH- → Zn (OH) 2 (a white ppt)

Some of these precipitates dissolve in excess aqueous sodium hydroxide or aqueous ammonia to form soluble
complex salts. These appear as colourless solutions, or in the case of copper (II) ions in excess aqueous
ammonia, a dark blue solution.
Copper (II), iron (II) and iron (III) ions are easily identified by the characteristic colour of their precipitates.
Aluminium ion, Calcium ion and Zinc ion all give the same observations when a few drops of aqueous sodium
hydroxide is used. However, only zinc ions will give a white precipitate soluble in excess aqueous ammonia.
TESTING FOR ANIONS (NEGATIVE IONS)
Sometimes we want to analyse a salt and find out what is in it. There are simple chemical tests which allow us to
identify the anion part (or negative ion) of the salt. These are often called spot tests – because they show the
results immediately. To test for anions or negative ions, different reagents (chemical solutions) are used unlike
for cations were only two solutions, which are sodium hydroxide and ammonia solution. The table below gives a
summary of how different anions are identified.
ANION IN SOLUTION TEST METHOD POSITIVE RESULT
1. Carbonate ion, CO32- Add any dilute acid Effervescence occurs and bubbles of
(preferably hydrochloric acid) carbon dioxide gas are produced.
The ionic equation shows how the reaction
takes place.
CO3 2- (aq) + H + (aq) → CO2 (g) + H2O (l)

2. Chloride ion, Cℓ - Add a few drops of dilute A white precipitate of silver chloride forms.
nitric acid to make the The ionic equation shows how the ppt is
solution acidified. Then add formed:
a few drops of silver nitrate. Ag + (aq) + Cl - (aq) → AgCl (s)

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 3. Iodide ion, I - Add a few drops of dilute A yellow precipitate of Lead (ii) iodide is
nitric acid to make the formed.
solution acidified. Pb 2+ (aq) + 2I - (aq) → PbI2 (s)
Then add a few drops of
aqueous lead (ii) nitrate

4. Sulphate ion, SO42- Add a few drops of dilute A white precipitate of barium sulphate
nitric acid to make the forms.
solution acidified. Then add Ba2+ (aq) + SO42- (aq) → BaSO4 (s)
a few drops of barium nitrate

5. Nitrate ion, NO3 - Add a few drops aqueous Damp or moist red litmus paper turns blue.
sodium hydroxide then Ammonia gas is produced
aluminium foil. Warm the
mixture carefully. Introduce
red litmus paper into the test
tube.

TEST FOR GASES

When recording observations for gases, it is important to record the following:

 Presence of effervescences
 Colour, if any and smell of the gas, if any
 Chemical test for the gas and test result
You have a sample of gas. You think you know what it is, but you’re not sure. So you need to do a test. Below
are some tests for common gases. Each is based on particular properties of the gas, including its appearance,
and sometimes its smell.

NAME OF GAS PROPERTIES, TEST DONE AND POSITIVE RESULT

Ammonia, NH3 Properties: Ammonia is a colourless alkaline gas with a strong sharp smell.
Test: Hold damp red litmus paper in it.
Result: The gas turns damp red litmus paper turns blue
Carbon dioxide gas, Properties: Carbon dioxide is a colourless, weakly acidic gas.
CO2 It reacts with limewater (a solution of calcium hydroxide in water) to give a white
precipitate of calcium carbonate:
CO2 (g) + Ca(OH) 2 (aq) → CaCO3 (s) + H2O (l)
Test: Bubble the gas through limewater.
Result: Limewater turns cloudy or milky.

Hydrogen gas, H2 Properties: Hydrogen is a colourless gas which combines violently with oxygen
when lit.
Test: Collect the gas in a tube and hold a lighted splint to it.
Result: The gas burns with a squeaky pop

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