Brief Notes on Uses of Normality

Normality is a concentration unit that is represented as gramme equivalent weight of solute in one litre of solution in a chemical solution. The concentration is expressed using an equivalency factor. Normality is measured in ‘N’, ‘equivalent/L’, or ‘milliequivalent/L’ units.

Normality is a measurement that is determined by the chemical process being studied. Normality, rather than molarity or any other unit of concentration of a chemical solution, is preferred in some cases. For example, in acid-base processes, normality is used to determine the concentrations of the ions ‘Hydronium’ (H3O+) and ‘Hydroxide’ (OH–). In this situation, the integer portion is 1/feq.

When is normality used?

Normality is most commonly used in three situations:

• In acid-base chemistry, it is used for determining concentrations. Normality, for example, indicates the concentrations of compound ions, hydronium ions (H3O+) or hydroxide ions (OH–) in a solution.
• In precipitation reactions, normality is used to calculate the number of compound ions that are expected to precipitate in a given reaction. 1/feq is an integer number in both acid-base as well as precipitation chemical reactions.

• It’s utilised in redox reactions to figure out how many electrons a reducing or oxidising substance can take or contribute. 1/feq might represent a fraction in redox processes.
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Examples of the use of normality

Example 1: For the following reaction, determine the normality of 0.1 M H2SO4 (sulfuric acid):

H2SO4 + 2NaOH → Na2SO4 + 2H2O

2 moles of H+ positive ions (2 equivalents) from sulfuric acid react with sodium hydroxide (NaOH) to generate sodium sulphate (Na2SO4) and water, according to the equation. Using the formula:

N = molarity x gram equivalents 

N = 0.1 x 2 N = 0.2 N

The number of moles of sodium hydroxide as well as water in the equation should not be mistaken. You don’t need the extra information because you already know the acid’s molarity. The only thing you need to know is how many moles of hydrogen ions are involved in the process.

Because sulfuric acid is such a powerful acid, you know it dissociates fully into its ions. 

Example 2: Calculate and determine the normality of the following reaction when the concentration is 1.0 M H3PO4.

H3AsO4 + 2NaOH → Na2HAsO4 + 2H2O

When we examine the reaction, we can see that just two of the H+ positive ions in H3AsO4 combine with NaOH to produce the product.

As a result, the two ions are equal in two ways. We’ll use the below-given formula to determine normality.

N = Molarity × number of equivalents

The use of normality has its limits

In acid-base chemistry, many scientists employ normality to avoid mole ratios in computations or to obtain more precise results. Normality is often employed in precipitation and redox processes, although it has significant limits. The following are the restrictions:

• In instances other than those listed above, it is not an appropriate unit of concentration. It’s a hazy measurement, and molarity or molality is a preferable option for units.
• A specified equivalency factor is required for normality.
• It isn’t a predetermined value for a certain chemical solution. Depending on the chemical process, the value might alter dramatically. To clarify, one solution can have many normalities for various reactions.
• Whilst normality is a useful measure of concentration; it cannot be utilised in all cases since its value is determined by an equivalency factor that varies depending on the kind of chemical process being studied. A solution of magnesium chloride (MgCl2), for example, maybe 1 N for the Mg2+ ion but 2N for the Cl– ion.
• While N is a useful quantity to know, it isn’t used as much in practical lab work as molality. Acid-base titrations, precipitation processes, and redox reactions all benefit from it.