The concept of the chemical equivalent offers a perspective on measurement in chemistry that focuses on reactive capacity rather than physical count. While the mole measures a fixed quantity of particles, the equivalent serves as a unit of chemical reactivity, particularly useful in stoichiometry and concentration determinations. Understanding equivalents allows chemists to move beyond simple mass-to-mass or mole-to-mole comparisons and directly assess a substance’s function within a reaction.
Defining the Chemical Equivalent
The chemical equivalent (Eq) is defined as the amount of a substance that will react with, supply, or displace one unit of another reference substance in a given reaction. In modern chemistry, this definition is tied to the mole of reactive units involved in the process.
Specifically, an equivalent is the amount of substance that will react with or supply one mole of hydrogen ions (\(H^+\)) in an acid-base reaction, or one mole of electrons in a redox reaction. The distinction from the mole is significant because one mole of a substance does not always equal one equivalent. For example, one mole of sulfuric acid (\(\text{H}_2\text{SO}_4\)) can supply two moles of hydrogen ions, representing two equivalents in a complete neutralization reaction. The number of equivalents is context-dependent, directly related to its role in a specific chemical change, unlike the mole, which is a fixed count of particles.
Calculating Equivalent Weight
The mass of one equivalent of a substance is known as its Equivalent Weight (EW). EW is calculated by dividing the substance’s Molar Mass (MM) by the \(n\)-factor (or equivalence factor), which represents the number of reactive units the molecule provides or consumes in a specific reaction.
\(n\)-Factor for Acids and Bases
For acids, the \(n\)-factor is the number of replaceable protons (\(H^+\)) the acid can donate (its basicity). For instance, hydrochloric acid (\(\text{HCl}\)) has an \(n\)-factor of one, so its equivalent weight equals its molar mass. Phosphoric acid (\(\text{H}_3\text{PO}_4\)) typically has an \(n\)-factor of three, meaning its equivalent weight is one-third of its molar mass. The \(n\)-factor for a base is the number of hydroxide ions (\(\text{OH}^-\)) it can supply, such as one for sodium hydroxide (\(\text{NaOH}\)) and two for calcium hydroxide (\(\text{Ca}(\text{OH})_2\)).
\(n\)-Factor for Ionic and Redox Reactions
For ionic compounds, the \(n\)-factor is determined by the total positive or negative charge supplied by the ions in a solution. For a salt like aluminum sulfate, \(\text{Al}_2(\text{SO}_4)_3\), the total positive charge is \(2 \times (+3) = 6\), making the \(n\)-factor six. In redox reactions, the \(n\)-factor represents the total number of electrons gained or lost per mole of the substance, reflecting the change in oxidation state. A substance’s equivalent weight is not a fixed constant like its molar mass, as it changes depending on the specific chemical process it undergoes.
Equivalents in Acid-Base and Redox Reactions
The greatest practical utility of the equivalent concept is found in the Law of Chemical Equivalence. This law states that one equivalent of any reactant reacts exactly with one equivalent of any other reactant. This principle dramatically simplifies stoichiometric calculations, particularly in acid-base neutralization and redox reactions, because the complex molar ratios derived from a balanced chemical equation become irrelevant.
In an acid-base titration, the number of equivalents of the acid must equal the number of equivalents of the base at the equivalence point. The calculation is straightforward because the \(n\)-factor already accounts for the reactive difference between the species, such as a diprotic acid reacting with a monohydroxy base. This simplification also applies to redox reactions, where the number of equivalents of the oxidizing agent equals the number of equivalents of the reducing agent.
Equivalent Concentration (Normality)
Normality (\(N\)) is the unit of concentration based on the equivalent concept, defined as the number of equivalents of solute dissolved per liter of solution (\(\text{Eq}/\text{L}\)). This unit directly incorporates the reactive capacity of the solute, making it a highly practical measurement for titration and volumetric analysis. Normality is directly related to Molarity (\(M\)), the concentration measured in moles per liter, through the \(n\)-factor.
The mathematical relationship is expressed as \(N = M \times n\), where \(n\) is the equivalence factor. For example, a \(1.0\text{ M}\) solution of \(\text{HCl}\) (with an \(n\)-factor of one) is \(1.0\text{ N}\). Conversely, a \(1.0\text{ M}\) solution of \(\text{H}_2\text{SO}_4\) (with an \(n\)-factor of two in full neutralization) is \(2.0\text{ N}\). Normality remains in use today in fields like environmental science and clinical chemistry, particularly for measuring ion concentrations in biological fluids, often reported in milliequivalents per liter (\(\text{mEq}/\text{L}\)).