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A. Acids and bases Historically there have been several influential definitions of acids and bases. The main definition in current use and the one we'll be using is Brønsted-Lowry definition, which defines acids as compounds capable of releasing a proton (H +) while bases are compounds that can accept a proton. Acid-base reactions are therefore proton-transfer reaction, where an acid releases a proton to form its conjugated base and a base accepts the proton to form its conjugated acid. The more general Lewis definition defines acids as compounds capable of receiving an electron pair while bases are electron pair donors. An older definition by Svante Arrhe-nius defined acids as compounds releasing a proton and bases as releasing a hydroxide anion. The process of releasing a proton (or any other ion) is called dissociation. As any other chemical reaction, dissociation will reach equilibrium, which can be characterised by an equilibrium constant, K. For a generic acid HA dissociating in water Protons in aqueous solutions become hydrated to H 3 O + called the hy-dronium or hydroxonium ion—the conjugated acid of water. HA + H 2 O −→ A − + H 3 O + (1) the equilibrium constant is K = [H 3 O + ][A − ] [HA][H 2 O] (2) Since the concentration of water in a dilute aqueous solution is pretty much constant we can rewrite the expression as K A = [H 3 O + ][A − ] [HA] (3) where K A is the dissociation constant of the acid (note that it is not the same as the original equilibrium constant). The value of K A is a good measure of the strength of an acid, i.e its ability to dissociate in solution: acids with K A > 10 −2 are usually called strong acids and the rest are weak acids. Please note that strength of an acid is not the same as its corrosive power or dangerousness! K A is often expressed as its negative decadic logarithm pK A = − log 10 K A. For example, HF is a relatively weak acid (K A = 6.8×10 −4) but can cause serious burns or dissolve glass. As the majority of acid-base reactions in biochemistry occur in aqueous solutions it is important to understand the role of water in these processes. Water is an amphoteric compound which means that it can both accept and donate a proton, i.e. it can act as an acid or as a base. In reaction (1) it accepts a proton from HA and thus behaves as a base. Similarly, water could behave as an acid and donate a proton to ammonia NH 3 + H 2 O − − − − NH + 4 + OH − (4) Moreover water can even donate protons to itself in a process called self-dissociation H 2 O + H 2 O − − − − H 3 O + + OH − (5)
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