Acid and Base Strength

The strength of a Brønsted-Lowry acid or base depends on the extent to which it ionizes in water. Although there are numerous solvents besides water, chemists discuss acid and base strength in relation to water because they use it so widely as a solvent. Chemists use the autoionization of pure water to determine the values for the concentrations of acidic and basic solutions. Autoionization is the reaction of two molecules of water with each other to give a hydronium ion, H3O⊕, and a hydroxide ion, -OH.

For this reaction, the amount of autoionization is extremely slight—at 25oC, it is 10–7 M (moles/liter). The concentrations of H3O⊕ and -OH are equal; that is, both measure 10–7 M. Chemists call this a neutral solution. If you add a compound that is more acidic than water, you increase the concentration of H3O⊕ ions and make the solution acidic. If you add a compound that is more basic than water, you increase the concentration of -OH ions and make the solution basic.
The product of the H3O⊕ and -OH concentrations in water is equal to 10–14 and is a constant, Kw. Chemists define Kw with the following equation.

Because the concentrations of H3O⊕ and -OH are equal in a neutral solution, you can easily calculate the concentration of both:

Because the product of the two concentrations is a constant, Kw, when one concentration increases, the other must decrease. For example, if you add -OH ions to water the concentration of the H3O⊕ decreases by whatever amount is necessary for the product of the two concentrations to still equal 10–14.
Because the hydronium ion concentrations can span a very wide range of values, from greater than 1 M down to less than 10–14 M, chemists measure the concentration of H3O⊕ on a logarithmic scale called pH. The pH values give the hydronium ion concentration of a solution. Therefore, measuring the pH of a solution is a means of quantifying the acidity of that solution. Chemists define this

measurement as the negative logarithm (base 10) of the H3O⊕ concentration, represented by the following equation:

pH = –log10[H3O⊕]

For simplicity, this book will normally refer to the H3O⊕ ion as the H⊕ ion from now on. If an equation shows the H⊕ ion present in aqueous solution, remember that it is actually the H3O⊕ ion.
This equation shows the general reaction of an acid in water:

Note that this reaction is an equilibrium. Most acid-base reactions are equilibrium reactions because the reactants only partly ionize. Strong acids and bases ionize completely in water. Weak acids and bases ionize only partly in water. An acidic, aqueous solution is any solution with a concentration of hydrogen ions greater than 10–7 M. Similarly, a basic solution is any solution with a concentration of hydroxide ions greater than 10–7 M.
To determine the relative strength of an acid or a base, you need to find out how much the acid or the base ionizes, or dissociates, in water at equilibrium. The equilibrium constant, Ke, gives this information and is defined as follows:

However, because water is the solvent and its concentration is essentially constant, a more meaningful value for acid ionization comes from multiplying the equilibrium constant by the water concentration:

Chemists call Ka the acid dissociation constant. The value of Ka specifies the strength of the acid. The stronger the acid, the larger the amount of dissociation and the larger the concentration of H3O⊕ ions. Thus, the stronger the acid, the larger the value of Ka. Strong acids completely dissociate in water and have large dissociation constants. Most organic compounds are weak acids and have dissociation constants in the range from 10–2 to 10–60.

Because acids have such a large range of values for their dissociation constants, chemists often convert those values to a logarithmic scale, similar to pH. The following equation defines this scale:

pKa = –log10 Ka