Chemistry: TestAS practice testlets

Many reactions do not run to completion but reach a state of dynamic equilibrium, in which the forward and reverse reactions proceed at the same rate. The concentrations of all substances then remain constant, although reactions continue at the particle level. A classic example is esterification, the reaction of a carboxylic acid with an alcohol to form an ester and water. Esters are responsible for the smell of many fruits and are used as solvents and flavourings.

acetic acid + ethanol ⇌ ethyl acetate + water K = [ester]·[water] ÷ ([acid]·[alcohol])

The position of the equilibrium is described by the equilibrium constant K, which relates the concentrations of products and reactants at equilibrium. For the esterification of acetic acid with ethanol, K is about 4 at room temperature. If one mole each of acid and alcohol are mixed in a given volume and x moles of ester are formed, then x moles of water are also formed, and 1 − x moles each of acid and alcohol remain. Setting x² ÷ (1 − x)² = 4 gives x ÷ (1 − x) = 2, and so x = 2/3: at equilibrium, two thirds of the acid have been converted into ester. The reaction is slow, and a few drops of concentrated sulphuric acid are added as a catalyst.

Le Chatelier's principle predicts how an equilibrium responds to disturbances: the system shifts so as to counteract the change. Adding more of one reactant, for example an excess of the cheaper alcohol, shifts the equilibrium towards the products and increases the yield of ester relative to the acid. Removing a product, for example by distilling off the ester or binding the water, also shifts the equilibrium to the right. A catalyst, by contrast, speeds up both the forward and the reverse reaction equally: equilibrium is reached faster, but its position does not change.

The reverse reaction, the splitting of an ester by water, is called hydrolysis. In the presence of a base such as sodium hydroxide, the acid formed is converted into its salt and removed from the equilibrium, so the hydrolysis runs almost to completion. This reaction, applied to fats, which are esters of glycerol and fatty acids, is the basis of soap-making and is known as saponification. Soaps are therefore the salts of fatty acids.

In aqueous solutions, acids release hydrogen ions (protons), which combine with water molecules to form oxonium ions, H₃O⁺. The acidity of a solution is expressed by its pH value, the negative decimal logarithm of the oxonium ion concentration in mol/L. A solution with [H₃O⁺] = 0.01 mol/L = 10⁻² mol/L has a pH of 2. Because the scale is logarithmic, a difference of one pH unit corresponds to a factor of ten in concentration: a solution of pH 3 contains ten times fewer oxonium ions than one of pH 2.

pH = −lg [H₃O⁺] [H₃O⁺] = 10⁻ᵖᴴ neutral at 25 °C: pH 7 pH + pOH = 14

Pure water at 25 °C has a pH of 7; acidic solutions have lower values, basic (alkaline) solutions higher values. Strong acids such as hydrochloric acid dissociate almost completely, so the oxonium ion concentration equals the acid concentration. Weak acids such as acetic acid dissociate only partly, so a solution of the same concentration has a higher pH. Diluting a strong acid tenfold with water raises its pH by one unit; however, no amount of dilution can make an acid solution basic, since the pH approaches 7 but does not exceed it.

A buffer solution resists changes in pH when small amounts of acid or base are added. It consists of a weak acid and its conjugate base in similar amounts, for example acetic acid and sodium acetate. Added oxonium ions are taken up by the acetate ions, forming acetic acid; added hydroxide ions react with acetic acid, forming acetate and water. The pH therefore changes only slightly. A buffer works best when the acid and its conjugate base are present in equal concentrations; the pH is then equal to the pKa of the acid, 4.76 for acetic acid.

Buffers are vital in living organisms. Human blood is kept within a narrow range around pH 7.4, mainly by the carbonic acid–hydrogen carbonate buffer system; values below 7.35 or above 7.45 are dangerous. Buffers are also used in laboratories, in the food industry and in cosmetics. When the buffer capacity is exhausted, because one of its components has been used up, the pH changes sharply. The capacity is larger the higher the concentrations of the two buffer components, so a concentrated buffer can absorb more acid or base than a dilute one of the same pH.