Ionic equilibria, Chemistry tutorial


Define: The equilibrium established between the unionized molecules and the ions in the solution of weak electrolytes is termed as the ionic equilibrium.

Example: CH3COOH ↔ CH3COO- + H+

The equilibria comprising acids and bases are vitally significant for a broad variety of reactions. The utilization of buffer solutions for pH control is of importance in living systems, agriculture and industrial methods. Likewise, the solubility equilibrium established in the solutions of sparingly soluble salts is as well very significant. For illustration, our bones and teeth are generally calcium phosphate, Ca3(PO4)2, a slightly soluble salt. We would mainly concentrate our attention primarily on the acid-base and the solubility equilibria and a few interrelated features concerning these equilibria like pH, buffer solutions and common ion effect.

Electrolytes and Non-Electrolytes: Acids, Bases and salts

The chemical reactions in solutions play a very significant role in chemistry. For illustration, sodium chloride doesn't conduct electricity in the solid state. Though, whenever it is dissolved in water, it becomes good conductor of electricity. Michael Faraday categorized the substances into two kinds: Electrolytes and Non-electrolytes


The substances that conduct electricity in their molten states or in the form of their aqueous solutions are termed as Electrolytes. For illustration: acids, bases and salts are the examples of electrolytes.

The electrolyte is any salt or ionizable molecule that, whenever dissolved in solution, will give that solution the capability to conduct the electricity. This is because if a salt dissolves, its dissociated ions can move freely in the solution, allowing a charge to flow.

Electrolyte solutions are generally made up whenever a salt is positioned into a solvent like water. For illustration - if table salt, NaCl, is put in water, the salt (that is, a solid) dissolves into its component ions, according to the dissociation reaction:

NaCl (s) → Na+ (aq) + Cl- (aq)

It is as well possible for substances to react with water to result ions in solution. For illustration, carbon-dioxide gas, CO2, will dissolve in water to produce a solution which includes hydrogen ions, carbonate and hydrogen carbonate ions:

2 CO2 (g) + 2 H2O (l) → 3 H+ (aq) + CO32- (aq) + HCO3- (aq)

The resultant solution will conduct electricity as it contains ions. It is significant to keep in mind, though, that CO2 is not an electrolyte, as CO2 itself doesn't dissociate into ions. Only compounds which dissociate into their component ions in the solution qualify as electrolytes.

a) Strong Electrolytes:

Strong electrolytes are the substances which only exist as ions in solution. Ionic compounds are generally strong electrolytes. Strong acids, strong bases and salts are the strong electrolytes. Whenever solid NaCl is put in water, it fully dissociates to form Na+ and Cl- ions.

b) Weak Electrolytes:

A weak electrolyte merely partially dissociates in solution and generates relatively few ions. Polar covalent compounds are in general weak electrolytes. Weak acids and weak bases are the weak electrolytes.


The substances which don't conduct electricity in their molten states or in the form of their aqueous solutions are termed as Non-Electrolytes. For illustration - urea, sugar, glycerin and so on are non-electrolytes.

Non-electrolytes are the compounds which do not ionize at all in solution. As an outcome, solutions having non-electrolytes will not conduct electricity. Generally, non-electrolytes are mainly held altogether by covalent instead of ionic bonds. A general illustration of a non-electrolyte is glucose, or C6H12O6. Glucose (that is, sugar) readily dissolves in water, but as it doesn't dissociate into ions in solution, it is considered as a non-electrolyte; solutions having glucose don't, thus, conduct the electricity.

General Concepts of Acids and Bases:

We are quite familiar with the words acid, base, acidity and so on. However how do we define an acid or a base? There is no general statement or definition of acids and bases. There are three different concepts of acids and bases (introduced by Arrhenius, Bronsted and Lowry concept and Lewis correspondingly) which are familiar. Each of these emphasizes a dissimilar feature of acid-base chemistry.

1) Arrhenius Concept:

The most generally employed concept of acids and bases was introduced by Svante Arrhenius in the year 1884. According to this theory, an acid is a substance which is capable of producing hydrogen ion (H+) via dissociating in aqueous solution. The reaction can be written as:

HA (aq) → H+ (aq) + A- (aq)

Here HA symbolizes the acid and A- refers to the acid molecule devoid of the hydrogen ion. Hydrochloric acid, HCl is an illustration of an Arrhenius acid whose ionization can be symbolized as:

HCl (aq) → H+ (aq) Cl- (aq)

The proton or hydrogen ion binds itself to the water molecule and form H3O+ ion that is termed as the hydronium ion.

H+ + H2O- → H3O

The hydronium ion is as well termed as oxonium ion or the hydroxonium ion. In the light of this fact, the first equation can be represented as:

HA (aq) + H2O (aq) → H3O+ (aq) + A (aq)

Base on the other hand is stated as a substance able of giving a hydroxyl ion (HO-) on dissociation in the aqueous solutions.

MOH (aq) → M+ (aq) + OH- (aq)

Here, M+ refers to the base molecule devoid of the hydroxyl ion. Sodium hydroxide is the illustration of the Arrhenius base, dissociating as,

NaOH (aq) → Na+ (aq) + OH- (aq)

Arrhenius theory is quite helpful and illustrates the acid-base behavior to a good extent. Though it has some demerits such as:

a) This is limited to just aqueous solutions and needs dissociation of the substance.

b) It doesn't describe the acidic behavior of certain substances which don't have hydrogen. For illustration: AlCl3. Likewise it doesn't describe the fundamental character of substances such as NH3 and Na2CO3 which don't encompass a hydroxide groups.

2) Lowry and Bronsted concept:

Bronsted-Lowry theory, as well termed as proton theory of acids and bases, a theory, introduced independently in the year 1923 by the Danish chemist Johannes Nicolaus Bronsted and the English chemist Thomas Martin Lowry, stating that any compound which can transfer a proton to any other compound is an acid and the compound which accepts the proton is a base. A proton is a nuclear particle having a unit positive electrical charge; it is symbolized by the symbol H+ as it comprises the nucleus of a hydrogen atom.

According to the Bronsted-Lowry system a substance can function as the acid only in the presence of the base; likewise, a substance can function as a base merely in the presence of an acid. Moreover, whenever an acidic substance loses a proton, it forms a base, termed as the conjugate base of an acid, and if a basic substance gains a proton, it makes an acid termed as the conjugate acid of a base. Therefore, the reaction between the acidic substance, like hydrochloric acid and a basic substance, like ammonia, might be symbolized by the equation:

HCl + NH3 ↔ NH4+ + Cl-

In the above equation, the ammonium ion (NH4+) is the acid conjugate to the base ammonia and the chloride ion (Cl-) is the base conjugate to the hydrochloric acid.

The Bronsted-Lowry theory expands the number of compounds considered to be acids and bases to comprise not merely the neutral molecules (example - sulphuric, nitric and acetic acids, and the alkali metal hydroxides) however as well some atoms and molecules having positive and negative electrical charges (that is, cations and anions). The ammonium ion, the hydronium ion, and a few hydrated metal cations are considered as acids. The phosphate, acetate, carbonate, sulphide and halogen ions are considered as bases.

3) Lewis Concept:

As from the above we are familiar that the Bronsted- Lowry concept doesn't based on the nature of the solvent (that is, a short coming of the Arrhenius concept erected). Though, similar to Arrhenius concept it doesn't describe the acidity of the substances which don't encompass a hydrogen atom (example - AlCl3) and the basicity of the substances without OH group (example - Na2CO3). G.N.Lewis stated in the year 1923 a yet another theory of acids and bases which comprises such substances as well. According to him, an acid might be stated as, any atom, molecule or ion which can accept an electron pair from any other atom, molecule or ion. A Lewis base on the other hand can be stated as, any atom; molecule or ion which can donate a pair of electrons let us take an illustration:

AlCl3 + NH3 → Cl3Al ← NH3

In the illustration above, AlCl3 is an electron deficient species. This accepts an electron pair from a molecule of NH3 that consists of a lone pair of electrons on N atom. Therefore, AlCl3 is a Lewis acid and NH3 is the Lewis base.

Attributes of acids and bases:

General features of Acids:

  • pH < 7
  • Sour taste (however you must never use this feature to recognize an acid in the lab)
  • React by a metal to form hydrogen gas
  • Increases the H+ concentration in water
  • Contributes H+ ions
  • Turns blue litmus indicator red

General features of Bases:

  • pH > 7
  • Bitter taste
  • Slippery feel
  • Raises the OH- concentration in water
  • Accepts the OH- ions
  • Turns red litmus indicator blue

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