Salts Archives

Uses of different salts in daily life

December 3, 2020December 3, 2020 by Raju

Uses of different salts in daily life

Many different types of salts can be found in nature.
The sea contains many salts such as sodium chloride, potassium chloride, magnesium chloride, magnesium sulphate and potassium bromide.
The earth’s crust is made up of minerals containing various types of salts such as calcium fluoride (fluorite), magnesium sulphate (Epsom salt), lead(II) sulphide (galena) and calcium carbonate (limestone).

Uses of salts

Salts play an important role in our daily life.
The following gives examples of salts and their uses in daily life.

Agriculture

Chemical fertilisers
(a) Most of the chemical fertilisers used by farmers are salts.
(b) These salts include ammonium chloride, ammonium nitrate, ammonium phosphate, potassium chloride and NPK fertilisers.
(c) NPK fertilisers contain the elements nitrogen, phosphorus and potassium obtained from salts such as ammonium phosphate and potassium chloride.
Pesticides
(a) Certain salts are used as pesticides to kill or destroy insects, pests, weeds and fungi.
(b) Examples include copper(II) sulphate, iron(II) sulphate, mercury chloride, sodium arsenate and sodium chlorate(V).

Medical field

Hydrated calcium sulphate, CaSO₄.2H₂O, is found in plaster of Paris. It is used to make plaster casts for supporting broken bones.
Iron(II) sulphate heptahydrate, FeSO₄.7H₂O is an ingredient in ‘iron pills’. It is used as an iron supplement for patients suffering from anaemia.
Potassium chloride is used as a substitute for patients who need a low intake of sodium salt.
Magnesium sulphate heptahydrate (Epsom salt) and sodium sulphate decahydrate (Glauber salt) are used as laxatives.
Sodium hydrogen carbonate is an ingredient in anti-acids. This salt can neutralise the excess acid secreted by the stomach.
Barium sulphate is used to make barium meals for patients who need to take an X-ray of their stomach. The salt helps to make internal soft organs like intestines appear on X-ray films.
Potassium manganate(VII) can kill bacteria and hence is suitable for use as an disinfectant.

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Food industry

Sodium chloride is used as a food additive in the preparation of food, to give foods a salty taste. It is also used as a food preservative.
Monosodium glutamate (MSG) is another food additive used in food preparation to enhance the taste of foods.
Sodium hydrogen carbonate is found in baking powders. Together with acid salts such as potassium hydrogen tartrate, calcium hydrogen phosphate and sodium hydrogen phosphate, they react in the presence of water to produce carbon dioxide. As the carbon dioxide escapes, it helps to make the cake or bread rise.
Sodium nitrite and sodium benzoate are examples of salts used as food preservatives. Sodium nitrite helps to preserve processed meat such as ham and sausages. Sodium benzoate is found in tomato sauce and chilli sauce.

Industry

Sodium hypochlorite is a bleaching agent and a disinfectant. Cleaning agents contain this compound.
Tin(II) fluoride is added to toothpaste and water to prevent tooth decay.
Sodium carbonate decahydrate is used in making soda-lime glass.
Silver bromide is used in the making of photographic paper and film.

The following table lists uses of some salts:

Salts	Uses
Sodium chloride (NaCl)	An essential requirement of our food In the preservation of food In curing fish and meat In making a freezing mixture which is used by icecream vendors In the manufacture of soaps
Sodium carbonate (Na₂CO₃)	As washing soda for cleaning clothes Used in the manufacture of glass, paper, textiles, caustic soda, etc. In the refining of petroleum In fire extinguishers
Sodium bicarbonate (NaHCO₃)	Used as baking soda In fire extinguishers As an antacid in medicine
Potassium nitrate (KNO₃)	To make gunpowder, fireworks and glass As a fertilizer in agriculture
Copper sulphate (CuSO₄)	Commonly called ‘blue vitriol’, used as a fungicide to kill certain germs In electroplating In dyeing
Potash alum (K₂SO₄.Al₂(SO₄)₃.24H₂O)	Used to purify water; makes suspended particles in water settle down As an antiseptic In dyeing
Calcium carbonate (CaCO₃)	Flooring in the form of marble To make lime (CaO),cement For extraction of iron
Silver nitrate (AgNO₃)	In photography for developing films
Ammonium nitrate (NH₄NO₃)	Fertilizers and explosives

General Properties of Salts

November 30, 2020November 30, 2020 by Raju

General Properties of Salts

Some of the characteristic properties of salts are:

Melting and boiling points: Salts are mostly solids which melt as well as boil at high temperatures.
Solubility in water: Salts are generally soluble in water. For example, sodium chloride, potassium sulphate, aluminium nitrate, ammonium carbonate, etc., are soluble salts while silver chloride, lead chloride, copper carbonate, etc., are insoluble in water.
Water of crystallization: Generally, salts are found as crystals with water molecules present in them. This water is called water of crystallization and such salts are called hydrated salts.
For example, copper sulphate crystal has five molecules of water for each copper sulphate molecule. This is written as CuSO₄.5H₂O. This water of crystallization gives the crystal its shape. It also gives colour to some crystals. On heating, hydrated salts lose their water of crystallization and, as a result, the crystals lose their shape and colour and change to a powdery substance.
The hydrated salts that have lost their water of crystallization are called anhydrous salts.
When hydrated copper sulphate is heated, it gives out water molecules to form white powdery anhydrous copper sulphate. On addition of water, this substance can convert back to a hydrated copper sulphate solution again.

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General Properties of Salts :

1. Reaction with an acid : When a salt reacts with an acid, another salt and acid are formed. For example, when sodium chloride is heated with sulphuric acid, sodium hydrogensulphate (at low temperature) and then sodium sulphate (at high temperature) are produced and hydrogen chloride gas is evolved.

2. Reaction with a base : A salt reacts with a base to produce another salt and base.

(NH₄)₂SO₄ + 2NaOH → Na₂SO₄ + 2NH₄OH

3. Reaction with a metal : Sometimes, a salt solution may react with a metal. For example, when an iron nail is dipped into an aqueous solution of copper sulphate, copper gets deposited on the surface of the nail and the ferrous sulphate formed remains in the solution.

CuSO₄ + Fe → FeSO₄ + Cu

This reaction shows that iron is more reactive than copper.
Thus, more reactive metal can displace a less reactive metal from a solution of its salt.

4. Behaviour of salts towards water :
When a salt is dissolved in water, the solution may be neutral, acidic or alkaline. This depends upon the nature of the salt used.

(i) A normal salt derived from a strong acid and a strong base gives a neutral solution. For example, the aqueous solutions of NaCl and K₂SO₄ are neutral to litmus.
General Properties of Salts 2
(ii) A normal salt derived from a weak acid and a strong base gives an alkaline solution. For example, the aqueous solutions of both sodium carbonate (Na₂CO₃) and sodium acetate (CH₃COONa) are alkaline.

Na₂CO₃ + 2H₂O → 2NaOH + CO₂ + H₂O

CH₃COONa + H₂O → CH₃COOH + NaOH

(iii) A salt derived from a strong acid and a weak base gives an acidic solution. For example, both aluminium chloride (AlCl₃) and ammonium chloride (NH₄Cl) make acidic aqueous solutions.

AlCl₃ + 3H₂O → Al(OH)₃ + 3HCl

NH₄Cl + H₂O → NH₄OH + HCl

(iv) Solutions of acidic salts are acidic to litmus, i.e., these solutions turn blue litmus paper red. For example, a solution of sodium hydrogensulphate (NaHSO₄) turns blue litmus paper red.
Sodium hydrogencarbonate (NaHCO₃) solution, however, is slightly alkaline.

Preparation of Salts

November 30, 2020November 30, 2020 by Raju

Preparation of Salts

Preparation of salts in the laboratory

The method used to prepare a salt depends on the solubility of the salt in water.
A soluble salt can be prepared from a reaction between an acid and a metal, a base or a carbonate.
- Reaction 1: Acid + alkali → salt + water
- Reaction 2: Acid + metal → salt + hydrogen
- Reaction 3: Acid + base → salt + water
- Reaction 4: Acid + metal carbonate → salt + water + carbon dioxide
Reaction 1 is used to prepare ammonium salts, sodium salts and potassium salts.
Reactions 2, 3 and 4 are used to prepare soluble salts except ammonium salts, sodium salts or potassium salts.
Recrystallisation is carried out to obtain pure crystals of the salt from the solution.

1. By the reaction between metal and acid : Certain metals (for example, Zn and Mg) react with HCl or H₂SO₄ to form salt and hydrogen.

Zn + 2HCl → ZnCl₂ + H₂

Zn + H₂SO₄ → ZnSO₄ + H₂

2. By the reaction between an acid and a base : All acid-base reactions (neutralization reactions) produce salts.

NaOH + HCl → NaCl + H₂O

CuO + 2HCl → CuCl₂ + H₂O

3. By direct union of a metal and a nonmetal : Sodium and chlorine combine directly to form sodium chloride.

2Na + Cl₂ → 2NaCl

Similarly, when sulphur is heated with iron filings, ferrous sulphide (FeS) is formed.

4. By the union between an acidic oxide and a basic :
Preparation of Salts 1
5. By the reaction between a metal and a base : When zinc is heated with an aqueous solution of NaOH, sodium zincate (salt) is formed with the evolution of hydrogen gas.

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How to select suitable methods for preparation of salts?

The flow chart below can help you select suitable methods for preparation of salts.

Qualitative Analysis of Salts

November 30, 2020November 30, 2020 by Veerendra

Qualitative Analysis of Salts

What is qualitative analysis?

Qualitative analysis of a salt Analysis is a chemical technique used to identify the ions present in a salt by analysing its physical and chemical properties and hence determine the identity of the salt.
It determines only the presence or absence of a particular ion in a given salt. This method does not determine how much of a particular ion is present.
For example, a student did a chemical analysis on a sample of salt X. His results showed the presence of sodium ions and bromide ions. His conclusion was the salt must be sodium bromide.
The procedure for testing salts in the laboratory consists of the following general steps.
1. Make initial observations of the physical properties of the salt.
2. Study the action of heat on the salt.
3. Make aqueous solution of the salt to test for anions and cations present.
4. Carry out confirmatory tests for cations and anions.

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Examining the colour and solubility of salts in water

Preliminary examination on the physical properties of a salt such as colour and solubility in water can help to indicate that certain cations and anions might be present.
Observations of physical properties only allow you to make inferences on the possibility of the presence of certain ions in the salts. This is because it is possible to have more than one salt sharing the same physical properties. For example:
(a) Iron(II) ions, nickel(II) ions and chromium(III) ions dissolve in water to produce green solutions.
(b) Sodium chloride and potassium carbonate are white solids. Both solids dissolve in water to produce colourless solutions.
Hence, observations of physical properties of salts cannot be used to confirm the identities of the ions. Its main purpose is to help us narrow down the choices of cations and anions present in a salt. Chemical tests still have to be carried out to confirm these ions.

Table lists the colours of some common salts.

Salt	Colour
Salt	Solid	Aqueous solution
Potassium salts Sodium salts Ammonium salts Aluminium salts Calcium salts Lead(ll) salts Zinc salts (with colourless anions)	White	Colourless
Carbonate salts Chloride salts Nitrate salts Sulphate salts (with colourless cations)	White	Colourless
Iron(II) salts: Iron(II) chloride Iron(II) nitrate iron(II) sulphate	Green	Green
Iron(III) salts: Iron(III) chloride Iron(III) nitrate Iron(III) sulphate	Brown	Brown
Copper(II) salts: Copper(II) chloride Copper(II) nitrate Copper(II) sulphate	Blue	Blue
Copper(II) carbonate	Green	(Insoluble)

Table shows the solubility of some common compounds.

Compound	Solubility in water
Sodium, potassium and ammonium salts	All are soluble
Nitrate salts	All are soluble
Ethanoate salts	All are soluble
Chloride salts	All are soluble except AgCl, HgCl and PbCl₂
Sulphate salts	All are soluble except BaSO₄, PbSO₄ and CaSO₄
Carbonate salts	All are insoluble except Na₂CO₃, K2CO₃ and (NH₄)₂CO₃
Metal oxides	All are insoluble except Na₂O, K₂O and CaO (slightly soluble)
Metal hydroxides	All are insoluble except NaOH, KOH and Ba(OH)₂

Example: Preliminary examination of solid X gave the following observations. Identify solid X. Explain your answer.

Green solid that insoluble in water
Dissolved in dilute acid with effervescence to form a blue solution

Solution:

The colour of an aqueous solution of a substance gives a more accurate inference about the nature of the ion present. Blue solution indicates presence of Cu²⁺ ion.
When a metal reacts with an acid, effervescence due to hydrogen gas liberated is observed. Effervescence due to carbon dioxide gas is seen when an acid reacts with a carbonate salt. Most metals are grey solids. Hence, the green solid X has to be a carbonate salt. X is copper(II) carbonate.

Test for gases

Gases are often produced from reactions carried out during laboratory tests on salts. Gases can be evolved when
(a) salts are heated.
(b) salts are reacted with acids or alkalis.
For example:
(a) Heating a carbonate salt produces carbon dioxide gas.
(b) Reacting a metal with dilute acids produces hydrogen gas.
By identifying the gases evolved, it is possible to infer the types of cations or anions present in a salt.
A gas can be identified by its colour, smell, effect on litmus paper and reactions with special reagents.

Gas	Characteristic
Ammonia Method: Bring a piece of moist red litmus paper to the mouth of the test tube. Observation: Red litmus paper turns blue.	Colourless gas Pungent smell Alkaline gas Turns moist red litmus paper blue Forms dense white fumes with hydrogen chloride Produced by heating a mixture of ammonium salt and alkali NH₄Cl(s) + NaOH(aq) → NaCl(aq) + NH₃(g) + H₂O(l)
Oxygen Method: Lower a glowing wooden splint into the test tube. Observation: Glowing wooden splint is relighted.	Colourless gas No effect on litmus paper Supports combustion Relights a glowing wooden splint Produced by heating a chlorate(V) or nitrate salt 2KClO₃(s) → 2KCl(s) + 3O₂(g) 2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)
Carbon dioxide Method: Bubble the gas through limewater. Observation: Limewater turns milky (chalky).	Colourless gas Acidic gas Turns moist blue litmus paper red Turns limewater milky Produced by heating a metal carbonate or acid- carbonate reaction CaCO₃(s) → CaO(s) + C02(g) CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
Hydrogen Method: Bring a lighted wooden splint to the mouth of the test tube. Observation: Gas burns with a ‘pop’ sound.	Colourless gas No effect on litmus paper Forms an explosive mixture with air Burns with a ‘pop’ sound Produced by an acid-metal reaction Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Chlorine Method: Bring a piece of moist blue litmus paper to the mouth of the test tube. Observation: Blue litmus paper turns red. It is then bleached.	Greenish-yellow gas Pungent smell Acidic gas Turns moist blue litmus paper red Bleaches litmus paper Produced by heating a mixture of manganese(IV) oxide and concentrated hydrochloric acid MnO₂(s) + 4HCl(aq) → MnCl₂(aq) + Cl₂(g) + 2H₂0(l)
Sulphur dioxide Method: Bubble the gas through acidified K2Cr207 solution or acidified KMn04 solution Observation: The orange K₂Cr₂O₇solution turns green or the purple KMnO₄ solution turns colourless.	Colourless gas Pungent smell Acidic gas Turns moist blue litmus paper red Reduces orange acidified dichromate(VI) ion to green chromium(III) ion Decolourises purple acidified manganate(VII) ion Produced by an acid-sulphite reaction Na₂SO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + SO₂(g)
Hydrogen chloride Method: Use a glass rod to bring a drop of concentrated ammonia solution to the mouth of the test tube. Observation: Dense white fumes are formed.	Colourless gas Forms fumes in moist air Acidic gas Turns moist blue litmus paper red Forms dense white fumes with ammonia gas Produced by heating a mixture of common salt and concentrated sulphuric acid H₂SO₄(aq) + NaCl(s) → NaHSO₄(aq) + HCl(g)
Nitrogen dioxide Method: Bring a piece of moist blue litmus paper to the mouth of the test tube. Observe the colour of gas evolved. Observation: Blue litmus paper turns red. Brown gas is evolved.	Brown gas Pungent smell Acidic gas Turns moist blue litmus paper red Reacts with water to form colourless solution Produced by heating a nitrate salt 2Cu(NO₃)O₂(s) → 2CuO(s) + 4NO₂(g) + O₂(g)

Action of Heat on Salts

November 30, 2020November 30, 2020 by Veerendra

Action of Heat on Salts

Heating a salt may cause it to decompose. The decomposition may result in
(a) a colour change
(b) evolution of a gas
(c) liberation of water vapour
Gases such as carbon dioxide, sulphur dioxide, nitrogen dioxide, ammonia or oxygen can be evolved. By identifying the gas or gases liberated, it is possible to pinpoint the anion present in the salt.
Examination of the residue can provide information to identify the cation in the salt.

Action of Heat on Carbonate Salts

Most metal carbonates decompose on heating to produce metal oxides and carbon dioxide gas.
Metal carbonate → metal oxide + carbon dioxide
When the carbon dioxide gas is bubbled through limewater, it will turn the limewater milky.

Table shows the action of heat on carbonate salts.

Carbonate salt	Action of heat
Potassium carbonate Sodium carbonate	Do not decompose
Calcium carbonate Magnesium carbonate Aluminium carbonate Zinc carbonate Iron(III) carbonate Lead(II) carbonate Copper(II) carbonate	Decompose to produce metal oxide and carbon dioxide Metal carbonate → metal oxide + carbon dioxide For example, CaCO₃(s) → CaO(s) + CO₂(g)
Silver carbonate	Decomposes to produce metal, oxygen and carbon dioxide 2Ag₂CO₃(s) → 4Ag(s) + O₂(g) + 2CO₂(g)
Ammonium carbonate	Decomposes to produce ammonia, water and carbon dioxide (NH₄)₂CO₃(s) → 2NH₃(g) + H₂O(l) + CO₂(g)

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Action of Heat on Carbonate Salts Experiment

Aim: To investigate the action of heat on carbonate salts.
Materials: Limewater, sodium carbonate, magnesium carbonate, calcium carbonate, zinc carbonate, lead(II) carbonate, copper(II) carbonate and potassium carbonate.
Apparatus: Test tubes, tongs, spatula, Bunsen burner and stopper with delivery tube.
Procedure:

About two spatulaful of copper(II) carbonate are placed in a test tube.
The colour of the carbonate salt is noted.
The test tube is stoppered with a delivery tube dipping into limewater as shown in Figure.
The carbonate salt is then heated strongly.
Any changes that occur to the limewater and also the colour of the residue when it is hot and when it is cold are recorded.
Steps 1 to 5 are repeated using each of the carbonate salts listed in Table to replace the copper(II) carbonate.

Observations:

Carbonate salt	Colour of salt before heating	Colour of residue		Effect on limewater
Carbonate salt	Colour of salt before heating	Hot	Cold	Effect on limewater
Copper(II) carbonate	Green	Black	Black	Limewater turns milky.
Sodium carbonate	White	–	–	No change.
Potassium carbonate	White	–	–	No change.
Calcium carbonate	White	White	White	Limewater turns milky.
Magnesium carbonate	White	White	White	Limewater turns milky.
Zinc carbonate	White	Yellow	White	Limewater turns milky.
Lead(II) carbonate	White	Brown	Yellow	Limewater turns milky.

Discussion:

Alkali metal carbonates such as sodium carbonate and potassium carbonate are stable to heat.
Most metal carbonates decompose on heating to produce metal oxides and liberate carbon dioxide gas.
The carbon dioxide gas forms a white precipitate with limewater, making the limewater milky.

Conclusion:
Heating a metal carbonate will decompose it into a metal oxide and liberate carbon dioxide. Group 1 metal carbonates are not decomposed by heat.

Action of Heat on Nitrate Salts

Nitrate salts also undergo decomposition on heating.
Most metal nitrates decompose to produce a metal oxide, nitrogen dioxide and oxygen.
Sodium nitrate and potassium nitrate decompose to produce nitrite salts and oxygen.
Nitrogen dioxide is a brown gas. It is an acidic gas that turns moist blue litmus paper red. Hence, dissolving it in water produces a colourless acidic solution.
2NO₂(g) + H₂O(l) → HNO₂(aq) + HNO₃(aq)
The colourless oxygen gas rekindles a glowing wooden splint.

Table: Action of heat on nitrate salts

Nitrate salt	Action of heat
Potassium nitrate Sodium nitrate	Decompose to produce a nitrite salt and oxygen 2KNO₃(s) → 2KNO₂(s) + O₂(g) 2NaNO₃(s) → 2NaNO₂(s) + O₂(g)
Calcium nitrate Magnesium nitrate Aluminium nitrate Zinc nitrate Iron(II) nitrate Iron(III) nitrate Lead(II) nitrate Copper(II) nitrate	Decompose to produce metal oxide, nitrogen dioxide and oxygen Metal nitrate → metal oxide + nitrogen dioxide + oxygen For example, 2Cu(NO₃)₂(s) → 2CuO(s) + 4NO₂(g) + O₂(g)
Silver nitrate	Decomposes to produce metal, nitrogen dioxide and oxygen 2AgNO₃(s) → 2Ag(s) + 2NO₂ (g) + O₂(g)
Ammonium nitrate	Decomposes to produce nitrous oxide and water NH₄NO₃(s) → N₂O(g) + 2H₂O(l)

Action of Heat on Nitrate Salts Experiment

Aim: To investigate the action of heat on nitrate salts.
Materials: Sodium nitrate, magnesium nitrate, calcium nitrate, zinc nitrate, lead(II) nitrate, copper(II) nitrate, potassium nitrate, iron(III) nitrate, iron(II) nitrate, blue litmus paper and wooden splint.
Apparatus: Test tubes, tongs, spatula and Bunsen burner.
Procedure:

About two spatulaful of copper(II) nitrate are placed in a test tube.
The colour of the nitrate salt is noted.
The nitrate salt is then heated strongly as shown in Figure.
The gases liberated are tested by
(a) lowering a glowing wooden splint into the test tube.
(b) bringing a piece of moist blue litmus paper to the mouth of the test tube.
The colour of the residue when it is hot and when it is cold are recorded.
Steps 1 to 5 are repeated using each of the nitrate salts listed in Table to replace the copper(II) nitrate.

Observations:

Nitrate salt	Colour of salt before heating	Colour of residue		Tests for gases evolved
Nitrate salt	Colour of salt before heating	Hot	Cold	Colour of gas	Glowing splint	Blue litmus paper
Copper(II) nitrate	Blue	Black	Black	Brown gas and colourless gas	Rekindles	Turns red
Sodium nitrate	White	White	White	Colourless	Rekindles	No change
Potassium nitrate	White	White	White	Colourless	Rekindles	No change
Calcium nitrate	White	White	White	Brown gas and colourless gas	Rekindles	Turns red
Magnesium nitrate	White	White	White	Brown gas and colourless gas	Rekindles	Turns red
Zinc nitrate	White	Yellow	White	Brown gas and colourless gas	Rekindles	Turns red
Iron(II) nitrate	Green	Black	Black	Brown gas and colourless gas	Rekindles	Turns red
Iron(III) nitrate	Brown	Brown	Brown	Brown gas and colourless gas	Rekindles	Turns red
Lead(II) nitrate	White	Brown	Yellow	Brown gas and colourless gas	Rekindles	Turns red

Discussion:

When nitrate salts are heated, they decompose to liberate nitrogen dioxide and oxygen.
Only sodium nitrate and potassium nitrate decompose to liberate oxygen.
Nitrogen dioxide is a brown gas that turns moist blue litmus paper red.
Oxygen is a colourless gas that relights a glowing wooden splint.

Conclusion:
Most metal nitrates decompose to produce a metal oxide, nitrogen dioxide and oxygen except sodium nitrate and potassium nitrate which decompose to produce nitrite salts and oxygen.

Action of Heat on Sulphate Salts

The normal sulphate salts are more stable to heat compared to the carbonates and nitrates.
Group 1 metal sulphates such as sodium sulphate and potassium sulphate do not decompose on heating. Group 2 metal sulphates such as calcium sulphate also do not decompose when heated.
The sulphates of heavy metals are decomposed into metal oxides and sulphur trioxide when heated.
Sulphur trioxide is a typical acidic oxide and dissolves in water to form sulphuric acid.
SO₃(g) + H₂O(l) → H₂SO₄(aq)
An exceptional case is iron(II) sulphate because it also forms sulphur dioxide gas.
2FeSO₄(s) → Fe₂O₃(s) + SO₃(g) + SO₂(g)
The green crystals of iron(II) sulphate turn into a brown solid of iron(III) oxide.
Ammonium sulphate sublimes when first heated. Further heating decomposes the salt into ammonia and hydrogen sulphate.
(NH₄)₂SO₄(S) – 2NH₃(g) + H₂SO₄(g)

Action of Heat on Chloride Salts

Chloride salts are stable to heat except ammonium chloride.
Initial heating of ammonium chloride causes the salt to sublime.
NH₄Cl(S) → NH₄Cl(g)
On further heating, decomposition takes place to produce ammonia and hydrogen chloride.
NH₄Cl(g) → NH₃(g) + HCl(g)
When ammonium chloride is heated in a test tube, the lighter ammonia gas will emerge first and turn a piece of moist red litmus paper blue. Hydrogen chloride, coming up next, will change the litmus paper from blue back to red.

Identification of salts by action of heat

When a salt is heated strongly, it may decompose. One or more gases may be liberated.
Each gas can be identified by
- noting its colour.
- testing it with moist blue or red litmus paper.
- testing it with limewater.
- testing it with glowing wooden splint.
- testing it with acidified potassium dichromate(VI) solution or acidified potassium manganate(VII) solution.
The colour of the residue when hot and cold must be noted to help in the identification of the salt.
The following table shows how to identify salts P, Q, R, S and T through gases liberated by the action of heat.

Test	Observation	Inference
Heat P in a test tube. Identify the gas/gases given off.	A colourless gas is liberated. It forms a white precipitate with limewater, i.e. limewater turns milky.	Carbon dioxide gas, CO₂, is produced. Carbonate ion, CO₂^2-, is present. P is a carbonate salt.
Heat Q in a test tube. Identify the gas/gases given off.	A brown gas is liberated which turns damp blue litmus paper red. A colourless gas is evolved which relights a glowing wooden splint. Solid residue is yellow when hot and white when cold.	Brown gas is nitrogen dioxide, NO₂. Colourless gas is oxygen, O₂. Nitrate ion, NO₃^–, is present. Residue is zinc oxide. Q is zinc nitrate.
Heat R in a test tube. Identify the gas/gases given off.	A colourless gas is given out. It turns orange acidified potassium dichromate(VI) green.	Sulphur dioxide gas, SO₂ is produced. Sulphate ion, SO₄^2-,is present. R is a sulphate salt.
Heat S in a test tube. Identify the gas/gases given off.	A colourless pungent gas is evolved. It turns damp red litmus paper blue. It forms dense white fumes with concentrated hydrochloric acid.	Ammonia gas, NH₃, is produced. Ammonium ion, NH₄⁺, is present. S is an ammonium salt.
Heat Tin a test tube. Identify the gas/gases given off.	A yellow-green gas is liberated. It turns damp blue litmus paper red and then white.	Chlorine gas, Cl₂, is produced. Chloride ion, Cl^–, is present. T is a chloride salt.