Reactants and products coexist in equilibrium, therefore reactant conversion to product is never greater than 100%. Equilibrium reactions may entail the breakdown of a covalent (non-polar) reactant or the ionisation of ionic compounds in polar solvents into their ions. This part will teach us about ionic equilibrium in ionic solutions. Substances in Ionic Equilibrium can be divided into two groups based on their capacity to conduct electricity, as shown below.
- Non-Electrolytes: These are molecules that have no electric charge, do not dissolve into constituent ions, and thus do not conduct electricity in an aqueous solution or molten form. As an example, consider the sugar solution.
- Electrolytes: These are compounds that dissolve into constituent ions in aqueous solutions and thus conduct electricity in aqueous solutions or molten states. For instance, salt solution, acid solution, base solution, and so on. Strong and weak electrolytes are the two types of electrolytes in ionic equilibrium.
Strong electrolytes are chemicals that totally ionize upon dissociation in their ionic solution, whereas weak electrolytes only partially ionize. In its aqueous solution, NaCl, for example, undergoes complete ionisation to produce sodium ions (Na+) and chloride (Cl–) ions, whereas acetic acid experiences partial ionisation to produce some acetate ions (CH3COO–) and hydrogen (H+) ions. In the event of a strong electrolyte, the dissociation reaction is said to be complete, proceeding only forward, whereas in the case of a weak electrolyte, the reaction is said to be reversible. In the case of a weak electrolyte, an ionic equilibrium is achieved between the ions and the unionised molecules, which is referred to as ionic equilibrium.
Ostwald’s Dilution Law: Degree of Dissociation
The application of the law of mass action to weak electrolytes in solution is known as Ostwald’s dilution law. AB is a binary electrolyte that dissociates into A+ and B– ions.
AB ⇌ A+ + B–
- For very weak electrolytes, since α <<< 1, (1 – α) = 1
- Concentration of any ion = Cα = √CK. Dilution causes a rise in the degree of ionisation. As a result, the degree of dissociation of a weak electrolyte is proportional to the square root of the dilution.
Limitations of Ostwald’s Dilution law: The law is only valid for weak electrolytes and fails totally for strong electrolytes.
Ionic Equilibrium Formulas
At equilibrium, it is vital to know what fraction of the initial amount of reactants is converted into products. The degree of dissociation/ionization refers to the fraction of the starting molecules that are transformed at equilibrium.
- Degree of dissociation or ionization = α = (At the start, the number of reactant molecules dissociated) / (At the start, the number of reactant molecules)
- In Ionic equilibrium, the degree of dissociation can be stated as a percentage.
% Degree of dissociation or ionization = α = (The number of reactant molecules that were dissociated or ionised at the start of the reaction.)/(At the start, the number of reactant molecules) × 100
Factors Influencing the Degree of Dissociation
An electrolyte’s degree of dissociation is determined by the following factors:
- Nature of Solute: Mineral acids, alkalies, and the majority of salts ionise almost entirely in aqueous solutions. These are known as strong electrolytes. Compounds such as organic acids and bases, as well as some inorganic acids such as HCN and inorganic bases such as NH4OH, on the other hand, ionise to a lesser amount. These are known as weak electrolytes.
- Nature of Solvent: Water, which has a high dielectric constant (i.e., insulating power), causes more ionisation than alcohol, which has a low dielectric constant. For example, an aqueous solution of hydrochloric acid rapidly conducts electricity, whereas its solution in toluene (an organic solvent) scarcely permits any electricity to pass through because few or no ions are generated in the latter situation.
- Dilution: The greater the amount of solvent utilised, the greater the amount of ionisation induced by it. As a result, the degree of ionisation in dilute solutions is greater than that in concentrated solutions.
- Temperature: Ionization increases as temperature rises.
- Nature of Other Substances Present in Solution: The presence of additional electrolytes with the same ion influences the degree of ionisation of one electrolyte. Ionization of ammonium hydroxide, for example, is inhibited by the presence of ammonium chloride in the solution. The common ion effect refers to the suppression of one electrolyte’s degree of ionisation when another electrolyte with a common ion is added.
Question 1: What is the Arrhenius Theory of Electrolytic Dissociation?
Arrhenius proposed an electrolyte behaviour theory in 1887 to explain a number of phenomena connected to electrolytes, such as higher values of colligative qualities of electrolysis compared to non-electrolytes, conductance, electrolysis, and so on.
Question 2: What are the Three Postulates of the Arrhenius Theory?
This theory’s three key postulates are:
- When an electrolyte dissolves in water, it separates into charged particles known as ions.
- Ions recombine to create unionised electrolyte. As a result, a dynamic or ionic equilibrium is achieved between the generated ions and the ionised electrolyte.
- When some electrolytes are dissolved in water, they virtually fully ionise. These are known as strong electrolytes.
Question 3: What are the Examples of Equilibrium?
It is an illustration of equilibrium when you are calm and stable. Because both hot and cold air enter the room at the same time, the overall temperature of the room remains constant. This is an illustration of equilibrium. Emotional or mental equilibrium A state of equilibrium or equality between opposing powers.
Question 4: What is a strong electrolyte?
A strong electrolyte is a chemical that completely ionises upon dissociation in its ionic solution, whereas weak electrolytes only partially dissociate and retain some of its ions in the solution as undissociated ions.
Share your thoughts in the comments
Please Login to comment...