SALTS (Ionic compounds)

Salt is a compound formed when either all of part of the ionisable hydrogen of the acid is replaced by a metallic ion or ammonium ion. Or A salt is an ionic compound consisting of a positive metallic or ammonium ion and a negative ion derived from an acid.

Salts get their names from the acids they are derived from. Examples are in the table below.

Types of salts

Normal salt

This is a salt produced when all the ionisable hydrogen of the acid is replaced by a metallic or ammonium ion. These salts do not contain ionisable hydrogen. Examples include; sodium chloride, NaCl; ammonium nitrate, NH4NO3; Magnesium sulphate, MgSO4; lead (II) bromide, MgBr2 and sodium phosphate, Na3PO4.

All normal salts have PH of 7 except salts formed from

Strong bases and weak acids e.g. sodium carbonate (Na2CO3) and potassium ehanoate (CH3COOK). These salts in solution have PH value more than

Strong acids and weak bases e.g. ammonium chloride (NH4Cl). The salts have aPH value less than 7 in

Acid salts

An acid salt is a salt formed when only part of the ionisable hydrogen of the acid is replaced by ammonium or metallic ion. These salts contain ionisable hydrogen and examples include: sodium hydrogensulphate, NaHSO4; calcium hydrogencarbonate, Ca(HCO3)2; and potassium hydrogen carbonate, KHCO3.

These acid salts behave like salts because they contain metallic ion and a negative ion derived from an acid; they behave like acids because the negative ions are capable of further ionization to yield hydrogen ion(H+) (i.e + ).

Basic salts

Basic salt are formed when insufficient acid is present to neutralize the available base.

E.g. basic zinc chloride (Zn(OH)Cl) and basic magnesium chloride (Mg(OH)Cl).

All monobasic acids form normal salts while dibasic and tribasic acids form both normal and acid salts.

Naming of salts

Salts are named by adding the name of the radical or ion of the acid after the name of the metal or ammonium. Examples are:

SOLUBILITY OF SALTS

Solubility is the amount of solute in grams required to saturate 100g of solvent (water) at a particular temperature.

A salt is described as soluble if it can dissolve in a given solvent and insoluble if it cannot dissolve in the solvent. Salts have varying degree of solubility in water as described below: All ammonium, sodium, and potassium salts are soluble in water. All nitrate salts are soluble in water.
All chloride salts are soluble in water except silver chloride, lead (II) chloride (sparingly soluble) and mercury (I) chloride.
All sulphate salts are soluble in water except lead (II) sulphate and barium sulphate. Calcium sulphate is sparingly soluble in water.
All carbonate salts are insoluble in water except sodium, potassium and ammonium carbonates.

Determining the solubility of a salt e.g. sodium chloride

Procedure

– Take about 50cm3 of distilled water in a beaker.
– Add sodium chloride crystals to the water a little at a time while stirring continuously until when no more salt dissolves. The solution formed is saturated.
– Weigh a clean evaporating dish and pour into it a little of the clear salt solution.
– Weigh the evaporating dish with the salt solution and evaporate the solution to dryness carefully through a water bath.
– Allow the evaporating dish to cool and reweigh the dish with the dry salt.

Results

Examples

30g of sodium chloride crystals were dissolved in 75 g of water at 80 . Calculate the solubility of sodium chloride at this

Solution

There fore, the solubility of NaCl at 80 temperature is 40g/100 g of water.

12.0g of potassium chlorate was carefully evaporated to dryness, 2.4 g of potassium chlorate crystals were left on the evaporating dish. Calculate the solubility of potassium chlorate in grams per 100g of water at room temperature.

Factors that affect the rate of solubility of salts

Amount of solvent

Solubility of most salts increase with increase in the amount of solvent used.

Nature of solvent/solute

Solubility of a salt may increase or decrease depending on the nature of solvent or solute.

Temperature

Solubility of most salts increase with increase in temperature. For example, potassium chlorate and potassium nitrate. Solubility of a few salts like calcium chloride and calcium sulphate decrease with increase in temperature. The solubility of sodium hydroxide and gases as well also decrease with increase in temperature.

Solubility curve

A solubility curve is a graph that shows how the solubility of a salt varies with temperature. The graph is obtained by plotting solubility (on the vertical axis) against temperature (on the horizontal axis).

Solubility curve of some common salts

The solubility of potassium chloride, potassium nitrate and potassium chlorate increase with increase in temperature. The solubility of potassium nitrate increases most rapidly, followed by potassium chlorate then potassium chloride.

The solubility of sodium chloride increases very slightly with increase in temperature. The solubility of calcium sulphate decreases with increase in temperature.

Uses of solubility curves

1. It can be used to find the solubility of a salt at a given
2. It gives the temperature at which a given amount of salt saturates 100g of
3. It can be used to explain the trend of solubility of salts.
4. A solubility curve can be used to calculate the mass of salt obtained by cooling a solution from a higher temperature to a lower

Mass of salt= (solubility at a higher temperature – solubility at a lower temperature)

For example, if a salt P with solubility of 180g/100g of water at 90˚C was cooled to a temperature of 30˚C where its solubility is 25g/100g of water. Calculate the mass of salt formed formed after cooling the solution.

Solution

Mass of salt= (solubility at a higher temperature – solubility at a lower temperature)

= (180-45)g = 155g

Application of solubility

1. Solubility is used to separate soluble salts from a mixture by fractional crystallization.
2. It is used in the extraction of salts from large water bodies like lakes and seas.

PREPARATION OF SALTS

The method of salt preparation depends on whether the salt is soluble in water or not. Soluble salts are prepared by crystallization and neutralization. Insoluble salts are prepared by precipitation or double decomposition. Other salts are prepared by direct synthesis.

Preparation of soluble salts

Soluble salts are prepared using dilute acids and metals, metals oxides, metal hydroxides and metal carbonates.

General procedure

1. Place some dilute acid in a beaker
2. Warm the acid and add the metal, metal oxide, metal hydroxide and metal carbonate bit by bit until in excess to ensure that the acid is completely used up.
3. Filter the excess metal, metal oxide, metal hydroxide or metal carbonate and collect the filtrate.
4. Saturate the filtrate by evaporating and allow the solution to cool as it cools to form the salt crystals.
5. Filter the crystals and wash them with water.
6. Dry the crystals in an oven, or under sun shine or between filter papers.

Preparation of salts from metals and dilute acids

Salts prepared by this method are soluble salts of iron, magnesium, aluminium and zinc. (I.e. metals higher than lead and lower than calcium in the reactivity series.)

N.B. Nitrates cannot be prepared using this method because dilute nitric acid being an oxidizing agent, does not react with metal to liberate hydrogen gas.

Example

Laboratory preparation of zinc sulphate crystals from zinc metal/powder

• Put dilute sulphuric acid in a beaker and heat it gently until when it‘s hot.
• Add zinc powder to the hot acid bit by bit while stirring until when the zinc powder is in excess.
• Filter off the excess zinc powder to obtain zinc sulphate solution as the filtrate.
• Saturate the filtrate by evaporating.
• Allow it to cool and form crystals of the salt.
• Filter the crystals and wash them with distilled water.
• Dry the crystals either in qaan oven of under the sun or between filter papers. Equation

Equation

Other salts formed in similar ways are:

Zinc chloride

Iron (II) sulphate

Preparation of salts from metal oxides and dilute acids Example

Preparation of copper (II) sulphate from copper (II) oxide in the laboratory

• Put dilute sulphuric acid in a beaker and heat it gently until when it‘s
• Add copper (II) oxide to the hot acid bit by bit while stirring until when the copper(II) oxide is in excess.

• Filter off the excess copper (II) oxide to obtain copper sulphate solution as the filtrate.
• Saturate the filtrate by
• Allow it to cool and form crystals of the
• Filter the crystals and wash them with distilled
• Dry the crystals either in an oven or under sunshine or between filter

Equation

Other examples of salts formed from metal oxides are:

Magnesium chloride

Copper (II) nitrate

Aluminium sulphate

Preparation of salts from insoluble metal carbonates Example

Preparation of lead (II) nitrate from lead (II) carbonate

• Pour dilute nitric acid in a beaker and warm it gently.
• Add lead (II) carbonate a little at a time. Effervescence occurs as carbondioixde is evolved.
• Continue adding the carbonate until when it is in excess and no more effervescence occurs.
• Filter off the excess carbonate to get a colourless filtrate.
• Evaporates the filtrate by heating gently to obtain a saturated solution.
• Cool the saturated solution to obtain white crystals of lead (II) nitrate salts.
• Wash the crystals with cold distilled water and dry either on sun shine, in an oven or between filter papers.

Equation
Lead (II) nitrate

Other salts prepared in similar ways are:

Barium chloride

Copper (II) sulphate

Preparation of salts from metal hydroxides Example

Preparation of lead (II) nitrate starting from lead (II) hydroxide

• Pour dilute nitric acid in a beaker and warm it gently.
• Add lead (II) hydroxide a little at a time while until when it is in excess.
• Filter off the excess hydroxide to get a colorless filtrate.
• Evaporates the filtrate by heating gently to obtain a saturated solution.
• Cool the saturated solution to obtain white crystals of lead (II) nitrate salts.
• Wash the crystals with cold distilled water and dry them either on sun shine, in an oven or between filter papers.

Other salts prepared in similar ways are:

Iron (II) nitrate

Copper (II) chloride

Zinc sulphate

Laboratory preparations of salts whose carbonates, oxides and hydroxides are insoluble.

These salts include potassium, sodium and ammonium salts. The salts can be prepared by titration method (neutralization).
Neutralization is a reaction between an acid and a base to produce a salt and water only.

General procedure

• Put a known volume of hydroxide of a metal in a conical flask.
• Add 2 or 3 drops of an indicator.
• Run a suitable acid from the burette until when the color of the mixture just changes. Note and record the volume of acid used.
• Measure accurately the same volume of hydroxide as before and titrate with exactly the same volume of acid as recorded above.
  NB. An indicator is not used in the second titration, since the volume of acid required to neutralize the fixed volume of base was already got.
• Stir and heat the solution to make it saturated.
• Allow the hot saturated solution to cool as it forms salt crystals.
• The crystals are filtered off, washed with cold distilled water and dried in an oven, under sun shine or between filter papers.

Preparation of sodium chloride crystals in the laboratory Procedure

• Put a known volume of sodium hydroxide in a conical flask.
• Add 2 or 3 drops of an indicator.
• Titrate the sodium hydroxide with hydrochloric acid from the burette until when the end point is reached (when the indicator changes color). Note and record the volume of acid used.
• Measure accurately the same volume of sodium hydroxide as before and titrate with exactly the same volume of hydrochloric acid as recorded above without using an indicator.
• Stir and heat the solution to make it saturated.
• Allow the hot saturated solution to cool as it forms salt crystals.
• The crystals are filtered off, washed with cold distilled water and dried in an oven, under sun shine or between filter papers.

Other salts prepared in similar ways are:

Ammonium sulphate ((NH4)2SO4)

Potassium nitrate, KNO3

Sodium sulphate

Preparation of insoluble salts

Insoluble salts are prepared by double decomposition or precipitation method. In this method, two soluble salts are mixed to form two new salts by exchange of radicals. One of the new salts formed is a soluble salt and one is an insoluble salt that appear as precipitates. The precipitate is filtered off and washed then dried.

Precipitation is the formation of solids when solutions are mixed.

A precipitate is the solid formed when two or more solutions are mixed.

Example

Preparation of barium sulphate (by reacting barium nitrate and sodium sulphate) Procedure

• Put a solution of barium nitrate in a beaker and add a solution of sodium sulphate to it. A white precipitate of barium sulphate immediately appears,
• Filter off the precipitate and wash with distilled
• Dry the precipitate (salt formed) under sun shine, in an oven or between filter papers.

Equation

This method can be used to prepare salts such as lead sulphate, aluminium chloride, silver chloride, silver carbonate and barium sulphate.

Preparation of lead (II) sulphate (by reacting lead (II) nitrate and sulphuric acid) Procedure

• Add dilute nitric acid to lead (II) nitrate solution in a beaker and stir the
• White precipitates of lead (II) sulphate is formed.
• Filter off the precipitates and wash with distilled water to remove traces of the
• Dry the precipitates in a steam oven or leave it to dry in

If any of the compounds to be used in the preparation of the salt is insoluble in water, it must first be made to dissolve in a mineral acid. For example, in the preparation of lead (II) sulphate using lead (II) oxide, the lead (II) is first divvolved in nitric acid to form lead (II) nitrate.

Equation

The lead (II) nitrate formed reacts with sulphuric acid to form lead (II) sulphate.

Preparation of salts by direct synthesis

Salts consisting of two elements (binary salts) can be prepared by direct synthesis/ direct combination.

Example:

In the preparation of sodium chloride from sodium and chlorine, burning sodium is lowered in a gas jar of chlorine. Sodium continues to burn in chlorine forming white fumes which settle into white solids (sodium chloride).

Equation

Other salts prepared by direct synthesis include:

Magnesium chloride

Effects of heat on salts

Carbonates

Potassium and sodium carbonates are very stable and are not decomposed by heat. But if the salts are hydrated, they lose their water of crystallization. In such a process, salts lose their crystalline nature and become amorphous.

All the other metallic carbonates decompose upon heating to give the oxide of the metal and carbondioxide gas.

Example

When white zinc carbonate is heated, it produces a colorless gas that turns lime water milky leaving a yellow residue when hot which turns white on cooling.

When lead (II) carbonate is heated, a brown residue (when hot) which becomes yellow on cooling and a colorless gas that turns lime water milky are produced.

When copper (II) carbonate is heated, black solids of copper (II) oxide is formed and a colorless gas that turns lime water milky evolved.

Ammonium carbonate decomposes to give ammonia gas, carbondioxide and water vapor.

Hydrogen carbonate of metals decompose to form carbonate of metals, carbondioxide gas and water vapor.

Example

Sulphates

Sulphates of sodium and potassium do not decompose on heating. When hydrated sulphates of potassium or sodium is heated, it loses its water of crystallization and becomes amorphous.

Sulphates of heavy metals decompose to give metal oxides and white fumes of sulphur trioxide gas. When heated more strongly, the sulphur trioxide gas decomposes to give sulphur dioxide and oxygen gas.

Examples

When hydrated copper (II) sulphate crystals are heated, they lose their water of crystallization and changes from blue crystals to white powder. The water condenses as a colorless liquid on the cooler parts of the test tube.

On further heating,the white powder gives off white fumes of a gas (SO3) and a black residue (CuO) is left.

When green solid of iron (II) sulphate is heated, it loses it water of crystallization and changes from green to dirty-yellow anhydrous solids.

When heated more strongly, it gives off sulphur dioxide (a colorless gas that turns potassium dichromate solution from yellow to green), white fumes of sulphur trioxide and brown residue of iron (III) oxide is left.

Nitrates

All nitrates decompose upon heating.

Sodium and potassium nitrates melt into colorless liquids then decompose upon heating to give their corresponding nitrites that form yellow solids on cooling and oxygen gas.

All nitrates from calcium down to copper decompose to give their corresponding oxides, brown fumes of nitrogen dioxide gas and oxygen gas.

Examples

Mercury (II)nitrate and silver nitrate decompose to give their corresponding metals, brown fumes of nitrogen dioxide gas and a colorless gas that relights a glowing spling (oxygen gas).

Ammonium nitrate sublimes upon heating to give dinitrogen oxide and water vapor.

Chlorides

Metallic chlorides are not affected by heat because hey are very stable. However, if they are hydrated, the lose their water of crystallization.

For example

Ammonium chloride sublimes on slight heating and on further heating, it decomposes to give ammonia and hydrogen chloride gases.

Equation

Effects of heat on hydroxides

Hydroxides of sodium and potassium are not decomposed by heat. If they are in solid forms, they absorb moisture and melt to form solutions.

However, the hydroxides of other metals decompose to give the corresponding oxides and water vapor.

END OF UNIT