Specific conductance of \(0.01 \mathrm{~N}\) solution of an electrolyte is \(0.00419\) mho \(\mathrm{cm}^{-1}\). The equivalent conductance of this solution will be: (a) \(4.19 \mathrm{mho} \mathrm{cm}^{2}\) (b) 419 mho \(\mathrm{cm}^{2}\) (c) \(0.0419 \mathrm{mho} \mathrm{cm}^{2}\) (d) \(0.209 \mathrm{mho} \mathrm{cm}^{2}\)

Specific conductance, denoted as \( \kappa \), is a measure of an electrolyte solution's ability to conduct electrical current. This inherent property depends on the presence of ions in a solution which act as carriers of electric charge. Imagine a scenario where an electrical field is applied across a solution; the ions will move towards the electrodes, thereby allowing electric current to flow. The higher the concentration and mobility of these ions, the greater the specific conductance.

While understanding this concept, students often confuse it with electrical conductivity. However, specific conductance is actually the conductivity of a solution per unit length. It is expressed in units of mho per centimeter (mho \(\mathrm{cm}^{-1}\)) or Siemens per meter (S/m). To put it into perspective, a solution with high specific conductance is like a crowd in a marathon running smoothly, with each runner (ion) quickly making their way to the finish line (electrode).

It's integral for students to grasp that specific conductance is influenced by a number a factors, including temperature, the nature of the electrolyte, and the concentration of ions. As these parameters change, the specific conductance would accordingly increase or decrease.

An electrolyte solution comprises a solvent, usually water, that contains dissolved substances (electrolytes) which are capable of conducting electricity. It's analogous to dissolving salt in water; the solid salt (the electrolyte) disintegrates and releases positively charged sodium ions and negatively charged chloride ions into the water. These free-moving charged particles make the solution conductive.

When we talk about electrolyte solutions, it's important to consider the dissociation aspect of electrolytes. Strong electrolytes like sodium chloride (\(\mathrm{NaCl}\)) dissociate completely in water to produce a larger number of ions that greatly contribute to electrical conductivity. On the other hand, weak electrolytes only partially dissociate, resulting in lower conductivity.

Understanding the behavior of electrolyte solutions is vital for various scientific and industrial applications, including battery technology, physiological processes in living organisms, and electrochemistry experiments. Students should remember that the nature of the electrolyte, the solvent, and the solution's concentration directly affect the overall conductivity of the solution.

The term 'normality' represents another way to express the concentration of a solution. It is specific to reactions in solution, where it reflects the concentration of reactive equivalents of a solute in a liter of solution. One can view normality as the molar concentration (\(\mathrm{M}\)) multiplied by the equivalent factor, which depends on the reaction under consideration.

Let's simplify this: normality is all about counting the number of particles that actually participate in a chemical reaction. For example, in acid-base reactions, normality considers the number of \(\mathrm{H^+}\) or \(\mathrm{OH^-}\) ions that can be donated or accepted. For a sulfuric acid (\(\mathrm{H_2SO_4}\)) solution, since one molecule can provide two \(\mathrm{H^+}\) ions, the normality would be twice the molarity.

Students should note the practicality of normality in titrations and other quantitative chemical analyses. It is especially useful in stoichiometry for balancing chemical equations where the reaction involves transfer of electrons, protons, or ions. Remember, the unit for normality is equivalents per liter (eq/L), which shows how this measurement is fundamentally connected to the reactants' stoichiometry within a given chemical process.