Salts in Chemistry, their Classifications, Solubility, Preparations and Colors

Meaning of Salts in Chemistry 

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A salt is an ionic compound formed when the hydrogen ion(s) of an acid are wholly (completely) or partly replaced by a metal ion, or an ammonium ion (NH₄⁺). Generally, acids reacts with bases to form salt and water only. This reaction is termed neutralization reaction.

Acid + Base → Salt + Water (Neutralization reaction)

Below are examples of neutralization reaction:

• HCl_(aq) + NaOH_(aq) → NaCl_(aq) + H₂O_(l)

• H₂SO_4(aq) + CuO_(s) → CuSO_4(aq) + H₂O_(l)

• HNO_3(aq) + KOH_(aq) → KNO_3(aq) + H₂O_(l)

• HCl_(aq) + NH_3(aq) → NH₄Cl_(aq)

 

A salt consists of a cation or positive ion (from the metal comprising the base, or NH₄⁺) and an anion (from the acid). In the above reactions, the ions making up the salts are: Na⁺Cl⁻, Cu²⁺SO₄²⁻, K⁺NO₃^-, and NH₄⁺SO₄²⁻ respectively.

Please read the introduction to acids here.

 

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Neutralization Reaction

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Neutralization is defined as the reaction between an acid and a base to form a salt and water only, with the pH of the mixture moving toward 7.

The ionic equation for a strong acid + strong alkali (neutralization reaction) is shown below through the hydrogen ion (H⁺) from the acid, and hydroxide ion (OH^-) from the base.

• H⁺_(aq) + OH⁻_(aq) → H₂O_(l)

 

The table below shows other products that can be formed alongside salts during various chemical reaction:

Reactants	Products
Acid + Alkali (soluble base)	Salt + Water
Acid + Insoluble base (metal oxide/hydroxide)	Salt + Water
Acid + Metal	Salt + Hydrogen gas
Acid + Carbonate	Salt + Water + Carbon dioxide

 

Please read on the physical and chemical properties of acids here.

 

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Classification of Salts (Solubility)

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Salts can be classified based on their solubility. This is summarized via the table below (which shows the soluble salts, and their insoluble exceptions).

Soluble salts	Insoluble exceptions
All Group I (e.g, Na, K) and ammonium salts	None
All nitrates	None
All chlorides	AgCl, PbCl₂
All sulfates	PbSO₄, BaSO₄, CaSO₄ (is slightly insoluble)
Sodium (Na), potassium (K) and ammonium (NH4+) carbonates	All other carbonates are insoluble
Na, K, ammonium hydroxides (soluble); Ca(OH)₂ is slightly soluble	Other hydroxides insoluble

 

Note: The solubility pattern of salts can help us determine which preparation method is used.

You can attempt questions and answers in chemistry here.

 

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Classification of Salts (based on Water of Crystallization)

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Water of crystallization in a salt is defined a fixed number of water molecules locked inside the salt's crystal structure (of the chemical compound). Compounds or salts that contain this water molecules are called hydrated salts. When these water molecules are removed through heating, the leftover dry substance is called an anhydrous compound.

Salts can generally be classified into two based on the presence or absence of water of crystallization in them. These are:

1. Anhydrous salt

These are salts (ionic compounds) that do not contain water of crystallization. This imply that they contain no water molecules trapped inside their crystal structure. They are usually powdery or granular in appearance. Examples of anhydrous salts include anhydrous CuSO₄ (white), NaCl (common table salt), anhydrous CaCl₂ (a strong drying agent), anhydrous CoCl₂ (a blue solid), and anhydrous Na₂CO₃ (a dry powder used as a cleaning agent).

 

2. Hydrated salt

These are salts (ionic compounds) that has water molecules trapped inside their crystal structure. These trapped water molecule is termed water of crystallization, and it binds to the metal ions in the salt using loose chemical bonds.

Hydrated salts have a fixed number of water molecules in their crystal Structure 

Examples of hydrated salts include CuSO₄·5H₂O (blue vitriol), Na₂CO₃·10H₂O (washing soda) and MgSO₄.7H₂O (epsom salt), CaSO₄.2H₂O (gypsum), etc.

Please read on bases and alkalis, their definitions, classification, physical and chemical properties here.

 

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Test for Hydrated and Anhydrous Salts

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This can be done via:

1. Heating a hydrated salt to know if it loses water (typically by changing its color).

2. Adding water to an anhydrous salt.

Heating a hydrated salt drives off water of crystallization (often with a color change). Adding water back to the heated hydrated salt can reverse this process. This experiment can be used as a chemical test for water.

You can read on the periodic table and classification of elements here.

 

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Recovery of Salts from Solution

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Soluble salts can be recovered from solution through heating. Insoluble salts (or salts that precipitates) can be recovered by filtering through a filter paper.

The recovery of salts from solution is further explained below:

1. Heat-stable soluble salts

These salts are stable to heat (that is; they do not decompose). When heated, the water evaporate to saturation point. We can test this process by dipping a glass rod into a hot evaporating solution during crystallization (and cooled outside the dish) to see if a tiny crystal or crust forms. 

The salt is cooled slowly in order to crystallize it. It is filtered, washed with a little cold distilled water, and dried.

Precaution: Never evaporate heat-stable soluble salt solution to complete dryness by strong heating if well-formed crystals are required.

Examples of heat soluble salts include NaCl, KNO₃, CuSO₄.

You can read on the differences between physical and chemical change in chemistry here.

 

2. Heat-sensitive salts

The solution containing the salt is heated gently so as to allow slow evaporation at room temperature. Note that the salt must not be heated strongly, in order to avoid its decomposition (or loss of water of crystallization).

Precaution: Never evaporate heat-sensitive salt under strong heat to prevent its decomposition (or prevent it from producing a fused mass instead of crystals).

Examples of heat sensitive salts include ammonium carbonate - (HH₄)₂CO₃, ammonium chloride - NH₄Cl, hydrogen bicarbonate - NaHCO₃, silver nitrate - AgNO₃, etc.

 

3. Insoluble salts (precipitates)

These salts do not dissolve in water. Instead, they remain as precipitate, and are this filtered directly, washed on a filter paper with distilled water, and dried. It is noteworthy to state that the salt should not be excessively dried if it is unstable to heat.

Examples of insoluble salts include AgCl, BaSO₄, PbCl₂, CaCO₃, PbSO₄ and Fe(OH)₃.

Please read on the preparation of bases and alkali, test and uses here.

 

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Preparation of Salts

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The method of preparing a salt is dependent on the solubility of the salt. This implies that soluble and insoluble salts can be prepared via their respective methods.

 

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Preparation of Soluble Salts

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Soluble salts dissolve in water. They can be prepared via the following methods:

1.Reaction of acid with a reactive metal. Hydrogen gas is liberated during this reaction. This reaction is shown via the equation below:

Acid + reactive metal → salt + H₂

• Zn_(s) + H₂SO₄_(aq) → ZnSO₄_(aq) + H₂_(g)

 

2. Reaction of acid with excess insoluble base or metallic oxide. This is achieved via the following steps:

• Add the solid reactant (insoluble base or metallic oxide, e.g Copper II oxide) in excess to a warm dilute acid. This is to ensure all the acid is completely used up.

• Stir the mixture until the reaction stops (fizzing or effervescence ceases).

• Filter the mixture to remove the unreacted excess solid as residue and collect the salt solution as the filtrate.

• Heat the filtrate to evaporate most of the water and create a saturated solution, then leave it to cool so crystals form.

• Filter out the crystals and dry them gently between filter papers.

Acid + excess insoluble base/oxide, filter off excess, then crystallize and dry.

CuO + H₂SO₄ → CuSO₄ + H₂O

Please read on mixtures and separation of mixtures here.

 

3. Reaction of acid with excess insoluble carbonate. This is achieved via the following steps:

• Add the solid reactant (insoluble carbonate, e.g copper carbonate or calcium carbonate) in excess to a warm dilute acid. This is to ensure that  all the acid is completely used up.

• Stir the mixture until the reaction stops (fizzing or effervescence ceases).

• Filter the mixture to remove the unreacted excess solid as residue and collect the salt solution as the filtrate.

• Heat the filtrate to evaporate most of the water and create a saturated solution, then leave it to cool so crystals form.

• Filter out the crystals and dry them gently between filter papers.

Acid + excess insoluble carbonate; filter, then crystallize and dry.

CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂)

Note: The above methods (1 - 3) of producing soluble salt can be termed crystallization. It generally follows pathway below below:

• React → Filter (if excess solid used) → Evaporate to saturation → Cool to crystallize → Filter crystals → Wash with a little cold distilled water → Dry.

Please read on liquid state of matter and characteristics of liquids here.

 

4. Neutralization reaction. This is the reaction of acids with alkali (through titration with an appropriate indicator). This is carried out via the steps below:

Step 1

• Use a volumetric pipette to measure and place a fixed volume of alkali (e.g., (25cm³) of sodium hydroxide) into a conical flask.

• Add a few drops of a suitable indicator; e.g, phenolphthalein or methyl orange.

• Carefully fill a burette with the acid (e.g., hydrochloric acid) and slowly add it to the alkali while swirling the flask until the indicator changes color permanently at the endpoint (neutralization reaction completed). This entire process is termed titration.

• Take note (write down) the exact volume of acid needed for complete neutralization. 

Step 2

• Repeat the above (step 1) but this time, without an indicator. Mix the exact measured volumes of acid and alkali together in a clean flask or beaker, but do not add any indicator this time to keep the salt pure. 

Step 3

• Evaporate the solution by heating the resulting neutral solution gently in an evaporating dish on a bunsen burner (or over a water bath) until most of the water has evaporated, and the remaining solution becomes saturated.

• Cool and crystallize: Allow the hot solution to cool slowly at room temperature so that the pure salt crystallizes out.

• Filter out the salt crystals and dry them gently between filter papers.

Acid + alkali (titration method, since both are soluble): find exact neutralizing volume with an indicator, repeat without indicator, then crystallize.

HCl + NaOH → NaCl + H₂O

Please read on the concept of ions and electrolytes here.

 

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Preparation of Insoluble Salts

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Insoluble salts can be prepared through a precipitation reaction called double decomposition; and through the direct combination of certain elements.

1. Precipitation (Double Decomposition)

Below are the steps to follow in double decomposition reaction:

• Dissolve two suitable soluble salts or solutions in water. Both soluble salts reacts by the exchange of radicals, with one providing the positive metal ion and the other, the negative non-metal ion). Stir thoroughly. This is shown via the equations below:

AgNO₃(aq) + NaCl(aq) → AgCl_(s)↓ + NaNO₃_(aq)

The AgCl salt is insoluble.

• Filter the mixture by pouring it through filter paper in a funnel. The insoluble salt stays on the paper as a residue or precipitate, while the other soluble salt passes through as a filtrate.

• Rinse the solid residue with distilled water while it is still on the filter paper to remove any remaining soluble impurities.

• Dry the salt on the filter paper.

In double decomposition, mix two soluble solutions whose ions combine to form an insoluble salt, then filter, wash, and dry.

AgNO₃_(aq)+ NaCl_(aq) → AgCl_(s)↓ + NaNO₃_(aq).

The reactants AgNO₃ and NaCl are soluble salts reacting via the exange of their radicals. The product, that is; silver chloride salt (AgCl) produced in the reaction above is an insoluble salt.

You can read on the differences between physical and chemical change in chemistry here.

 

2. Direct combination of elements

It follows the principle of heating certain reactive metals and non-metals together, causing them react directly and form an insoluble binary salt. Examples of insoluble salts produced this way include iron(II) sulfide (FeS) and iron III chloride (FeCl₃).

• Procedure for preparation iron(II) sulfide (FeS): Heat iron powder and sulfur powder directly in a test tube. They react, giving off heat to the surrounding (exothermic reaction) to form solid iron(II) sulfide, which does not dissolve in water.

Fe_(s) + S_(s) + heat →FeS_(s)

• Procedure of preparing Iron(III) chloride (FeCl₃): Pass dry chlorine gas over heated iron.

2Fe_(s) + 3Cl_2(g) + heat →2FeCl_3(s)

Please read on the physical properties of metals in chemistry here.

 

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Color of Salts

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The color of a salt generally reflects the presence of a transition metal ion. This is so because transition metals are characterized with colored ions. However, note that group I, II and ammonium salts are usually white or colorless.

The table below shows the salts of various metals and their respective colors.

Salt	Color
Copper(II) salts	Blue
Iron(II) salts	Pale green
Iron(III) salts	Yellow/brown
Zinc, lead, calcium, sodium, potassium salts	White
Nickel(II) salts	Green
Manganate(VII) salts (KMnO₄)	Purple
Chromate(VI) salts	Yellow
Dichromate(VI) salts	Orange

 

Please read the continuation of this topic here (types of salts, properties of salts, deliquescence, efflorescence, hygroscopy, drying agents, hydrolysis and uses of salts here.