Classwork Series and Exercises {Chemistry – SS2}: Salts

Chemistry, SS 1 Week 3

Topic: Salts

Introduction
A salt is an ionic compound formed when the hydrogen of an acid is partly or completely replaced by a metal ion or ammonium ion. All salts are chemically and electrically neutral.
Example:

Diagram above shows that when the hydrogen ion in nitric acid is replaced by Na⁺, Ca²⁺, NH₄⁺ or Al³⁺ ions, salts are formed.

Other examples:
Barium nitrate, zinc sulphate and tin nitrate are salts

There are 4 types of salt, these are:

• Nitrate
• Chloride
• Sulphate
• Carbonate

Classification of Salt

Salts are classified into four different types:

• Normal salts
• Double salts
• Mixed salts
• Complex salts

NORMAL SALTS: Salts which produce one simple cation and one simple anion in aqueous solution are called normal salts.

The ions present in simple salt can be tested easily. Based on the nature of ions produced they are further classified into:

1. Neutral salts
2. Acidic salts
3. Basic salts

Neutral salts: Salt which is formed by complete neutralization of strong acid and base or weak acid and weak base is called neutral salt and it neither produces H⁺ or OH^– in solution

Examples:

• NaCl (formed by neutralization of NaOH and HCl)

HCl_(aq) + NaOH_(aq) —–> NaCl_(aq) + H2O_(l)

• K₂SO₄(formed from KOH and H₂SO₄)

H₂SO_4(aq) + 2KOH_(aq) —–> K₂SO_4(aq) + 2H₂O_(l)

Acidic salts: Salt formed by partial neutralization of poly basic acid with a base. Acidic salt produces H⁺ in solution.

Example:

• NaHSO₃(formed When poly basic acid H₂SO₃ is partially neutralized by NaOH)

NaOH_(aq) +H₂SO_3(aq) ——–> NaHSO_3(aq) +H₂O_(aq) ( Acidic salt formed by partial neutralization)

2NaOH_(aq) +H₂SO_3(aq) ———> Na₂SO_3(aq) + 2H₂O_(l) (Neutral salt formed by complete neutralization)

Basic salt: Salt which is formed by partial neutralization of poly acidic base (Ca(OH)₂,Fe(OH)₃ etc.) with an acid. Such a salt produces OH^– ion in solution

Example:

• Ca(OH)Cl (formed by partial neutralization of Ca(OH)₂ with HCl)

Ca(OH)_2(aq) + HCl_(aq) ——-> Ca(OH)Cl_(aq) + H₂O_(l) (Basic salt)

Ca(OH)_2(aq) + 2HCl_(aq) ——–> CaCl_2(aq) + H₂O_(l) (Neutral salt)

Double salt: A salt formed from two different salts and whose solution gives test for all the ions present in it. In other words, double salt contains two different positive metallic ions and common negative acid radical or a positive metallic ion and ammonium ion and a common negative acid radical

Examples:

• FeSO₄ (NH₄)₂SO₄ .6H₂O – Ammonium Iron (III) tetraoxosulphate (VI) or Mohr salt
• K₂SO₄ Al₂(SO₄)₃ .12H₂O – Potassium Aluminium tetraoxosulphate (VI) dodecahydrate (Potash alum)

Mixed salt: When an acid is simultaneously neutralized by two bases or when a base is neutralized by two acids. They produce two cations or two anions and one cation.

Example: Ca(OCl)Cl -bleaching powder

Complex salt: Salt which produces one simple ion and a complex ion in aqueous solution. A complex salt does not answer the ions present in complex ion.

Example: K₄(Fe(CN)₆ – Potassium hexacyanoferrate (II)

Solubility of Salts

Solubility is the ability of a compound to dissolve in a solvent.

Table below shows the solubility of the salts of nitrate, sulphate, chloride and carbonate.

Salt	Solubility
Salt of potassium, sodium and ammonium	All soluble in water
Salt of nitrate	All soluble in water
Salt of sulphate	Mostly soluble in water except: (Pb) Lead sulphate (Ba) Barium sulphate (Ca) Calcium sulphate
Salt of chloride	Mostly soluble in water except: (Pb) Lead chloride (Ag) silver chloride (Hg) mercury chloride
Salt of carbonate	Mostly insoluble in water except: Potassium carbonate Sodium carbonate Ammonium carbonate

 

Notes: Lead halide such as lead(II) chloride (PbCl₂), lead(II) bromide (PbBr₂) and lead (II) iodide (PbI₂) are insoluble in cold water but soluble in hot water.

Preparation of Soluble Salts

There are 2 things to be considered when preparing a salt:

• What are the chemical used?
• How to separate the salt from other substance?

Method used to prepare salt depends on the solubility of the salt. Soluble salts are prepared from the reactions between an acid with a metal/ base/ metal carbonate.

Diagram below shows the chemical reaction that can be used to prepare the soluble salts.

Preparing Salts of Potassium, Sodium and Ammonium

Potassium, sodium and ammonium salts are usually prepared through the reactions of acids with alkalis. Reacting acid with alkali will produce salt and water. The salt is prepared by titration method of acid and alkali using an indicator.

Acid + Alkali → Salt + Water

Steps to Prepare the Salts of Potassium, Sodium and Ammonium through Titration

Step 1 – Titration to Find the End Point

The end point is the point in a titration at which the two reactants have completely reacted. It is often marked by a colour change.

Step 2 – Titrate Without Indicator

The product obtain in step 1 is contaminate by the indicator. The reaction is repeated by using the same amount of reactants as in step 1, without using any indicator.

Step 3 – Crystallisation

Step 4 – Filtration and Drying

Preparing Salts of Non-“Potassium, Sodium and Ammonium”

The salt non-potassium, sodium and ammonium is prepared by reacting acid with insoluble metal/metal oxide/metal carbonate:

Acid + Metal Salt + Hydrogen (Displacement reaction)

Acid + Metal oxide Salt + Water (Neutralisation Reaction)

Acid + Metal carbonate Salt + Water + Carbon Dioxide

Below is the steps in preparing the soluble non-potassium, sodium and ammonium salts

Step 1 – The Reaction

Add metal/metal oxide/metal carbonate powder until excess into a fixed volume of the heated acid

Step 2 – Filtration 1 to Remove Excess Reactant

Filter the mixture to remove excess metal/metal oxide/metal carbonate

Step 3 – Crystallisation

Evaporate the filtrate until it becomes a saturated solution

Dip in a glass rod, if crystals are formed, the solution is saturated.

Step 4 – Filtration 2 to Collect the Solid Salt

Cooled at room temperature

Filter and dry the salt crystals by pressing them between filter papers

Chemical equation(s) for the reaction that can be used to prepare the following salts:

• Sodium Chloride

NaOH + HCl → NaCl + H₂O

• Ammonium Nitrate

NH₃ + HNO₃ → NH4NO₃

• Potassium sulphate

KOH + H₂SO₄ → K₂SO₄ + H₂O

• Zinc Sulphate

ZnO +  H₂SO₄ → ZnSO₄ + H₂O

Zn  + H₂SO₄ → ZnSO₄ + H₂

ZnCO₃  + H₂SO₄ → ZnSO₄ + H₂O + CO₂

• Lead(II) nitrate

PbO +  2HNO₃ → Pb(NO₃)₂ + H₂O

Pb  + 2HNO₃ → Pb(NO₃)₂ + H₂

PbCO₃  + 2HNO₃ → Pb(NO₃)₂ + H₂O + CO₂

• Copper sulphate

CuO +  H₂SO₄ → CuSO₄ + H₂O

CuCO₃  + H₂SO₄ → CuSO₄ + H₂O + CO₂

Preparing Insoluble Salts

Insoluble salts can be made by ionic precipitation (is also called double decomposition/double displacement). This involves mixing a solution that contains its positive ions with another solution that contains its negative ions.
The equation of the reaction that can be used to prepare the following salt:

• Calcium sulphate

CaCl₂ + NaSO₄ → CaSO₄ + 2NaCl

Ca(NO₃)₂ + ZnSO₄ → CaSO₄ + Zn(NO₃)₂

• Lead chloride

Pb(NO₃)₂ + 2NaCl → PbCl₂ + 2NaNO₃

• Copper carbonate

CuSO₄ + Na₂CO₃ → CuCO₃ + Na₂SO₄

CuCl₂ + K₂CO₃ → CuCO₃ + 2KCl

Cu(NO₃)₂ + Na₂CO₃ → CuCO₃ + 2NaNO₃

Efflorescent, Deliquescent and Hygroscopic Substances

These terms are used to describe the behaviour of substances when exposed to the atmosphere. These substances either give up water to the atmosphere or absorb water from the atmosphere

Efflorescent compounds: These are hydrated salts that lose part or all their water of crystallisation to form a lower hydrate salt when exposed to atmosphere. Examples are washing soda i.e. sodium trioxocarbonate (IV) decahydrate (Na₂CO₃ .10H₂O) which changes to monohydrate Na₂CO₃.H₂O

Equation: Na₂CO₃ .10H₂O ——> Na₂CO₃.H₂O + 9H₂O

Another example is sodium tetraoxosulphate (VI) decahydrate known as Glawber’s salt. It changes to the anhydrous salt on exposure to air for some days.

Na₂SO₄.10H₂O —–> Na₂SO_4(aq) + 10H₂O

Deliquescent compounds: These are compounds which when exposed to the atmosphere will absorb moisture and dissolve in it to form solutions

Examples are:

• Sodium hydroxide pellets
• Iron (III) chloride
• Phosphorous (V) oxide
• Potassium hydroxide
• Calcium chloride

Hygroscopic Compounds: These are compounds which when exposed to the atmosphere will absorb moisture, but will not dissolve in it and will only become sticky

Examples are:

• Copper (II) oxide
• Quicklime (Calcium oxide)
• Sodium trioxonitrate (V)

The only liquid hygroscopic compound is concentrated tetraoxosulphate (VI) acid. It will absorb moisture from the atmosphere and then becomes dilute