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The Theory of Acids and Bases

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The Theory of Acids and Bases THE THEORY OF ACIDS AND BASES By F. M. HALL, )1.SC. Wollongong University College, N.S. W. , Australia The theory of acids and bases, like many and the term hydrolysis relates to the inter­ other chemical theories, has undergone action of the ions of the salt with water to numerous changes in recent times. Always give (a) a weak acid or a weak base, or (b) a the changes have been such as to make the weak acid and a weak base. theory more general. The three main Applying the Law of l\llass Action to such theories in use today are: a system, the hych·olysis constant, at equili­ (l ) the Water or Arrhenius Theory; brium, may be written as (2) the Proton or Br0nsted- Lowr.v l [Base] [Acid] [OH-) [HA) Theory; (" = (Unhydrolysed salt] = [A- ] (3) the Electronic or Lewis Theory. and if x is the extent to which hydrolysis WATEl~ OR AlmHENIUS 'l'HEORY occurs and c is the molar concentration of salt The Water or Arrhenius Theory was in solution widely accepted up to the early years of this xcx xc x 2c century. It defines an acid as a hydrogen K" = (1 - x)c = (1 - x) compound ionizing in water to give hydrogen The equilibria involved for the salt of a ions, and a base as a hy<iToxyl compound weak acid and strong base in water are which gives hydroxide ions in water. The neutralization reaction between an acid and · _ rH+] [OR-) (a) H 20 ~ H+ + OR- ; li. = [HzO] a base produces a salt and water only: HA(acid) + BOH(base) ""' BA(salt) + [H +] [A - ] = H 20(water) (b) HA ~ H+ + A-; Xa [HA] and the salts so formed may be classified into (c) A-+ H 0 ~ HA + OR- ; four main groups, viz.: 2 , [I'IA) [OR - ] (i) those derived from strong acids and Xb = [A- ) [H20J strong bases ; Since [H20] is effectively constant in dilute (ii) those derived from weak acids and solutions we may replace]( and Kh' by two strong bases ; new constants, Kw and J(h respectively, (iii) those derived from strong acids and defined by weak bases; Xw = [H +) [OR-] (iv) those derived from weak acids and _ [.HA] [OR-] } 7 weak bases. and '-h - [A- ) When dissolved in water these various salts Kw is called the ionic product of water and do not necessarily give neutral solutions. 14 2 2 Interaction between the salt and water has a value of ,...._, 10- mole litre- at 25°C. (hydrolysis) accounts for the acidity, alka­ It follows that linity or neutrality of the solution. For K,. _ [H+] [OR-] [HA] _ [OH-] [HA) _ K example, a salt BA derived from a weak acid Ea - [H+) [A-] - [A ) - 11 HA and a strong base BOH gives an alkaline If two assumptions are now made: solution in water because the acid derived by hydrolysis gives few hydrogen ions and the (i) that there is only a pure aqueous base derived by hydrolysis gives many solution of the salt of the acid HA (no hydroxide ions. This may be shown as added hydrogen or hydroxide ions) ; B+ + A -+ H 20 ""' IIA + B + + OR- and 91 92 EDUCATION IN CHEl\flSTRY (ii) that the concentration of hydroxide A solution then is not necessarily neutral at ion obtained from the ionization of pH 7 but rather when the hydrogen ion con­ water is negligible compared with that centration is equal to the hydroxide ion resulting from the hydrolysis of the concentration. However, at 25°C a neutral salt, solution has a pH of 7 and all subsequent then [OH- ] = [HA] during hydrolysis of the salt, calculations refer to this temperature. [OH- F K w PROTO~ OR BR0NSTED- LO\VRY TTIEORY ancl [A-] = !{11 = K;; The Arrhenius Theory makes use of and [OH- ] = J [A -]. Kw hydroxide ions, which may not exist in non­ ] (,. aqueous solvents, and does not cover weak J(w bases. In 1923 Bnmsted and Lowry put but [H +] = [OH - J forward a more general theory of acids and bases which incorporates all protonic solvents, hence [H +] = w • JK [A-IC] and not just water. They defined an acid as a substance ·which yields a proton and a base or pH = ?-pKw + t piCa + -~ log c as a substance which can combine with a where c = [A-], the stoichiometric molar proton. concentration of the salt.* This is only Thus an acid HB dissociates to give a justified if the hydrolysis of the salt is very proton and its conjugate base. Alternately a small. base B can combine with a proton to yield the This equation then allows calculation of the conjugate acid BH+ of the base. In general pH of an aqueous soh1tion of a salt or the pH terms, this may be written as at the equi valence point of a weak acid- strong Acid 1 + Base 2 "" Acid 2 + J:3ase1 base titra.tion. Similarly for a strong acid-weak base in which the proton is pat•titioned between system the pH is given by two bases and the equilibrium constant is determined by the relat ive affinities for the pH = t pi(,,. - ·~· pKb - -~ log c proton. where J(b is the ionization constant of the A Br0nsted acid may be an electrically base and c is the molar concentration of the neutral molecule, e.g. HCI, a cation , e.g. salt. C6H5~TH3+ , or an anion, e.g. H S04- , whilst a It should be noted that ICv , the ionic Br0nsted base may be a neutral molecule, product of water, like any other equilibrium e.g . aniline C6H 5NH2 , or an anion, e.g. Cl-. constant, varies with temperature. Table I One important result of this definition of an illustrates this point. acid is that the strength of an acid depends upon its environment. The acidic strength TABLE I of a weak acid is etihanced by its solution in a strongly basic solvent and the basic strength pH of neutral solutiort of a weak base is enhanced by its solution in Temp. ' C Kw x 101'1 [H+] = [OH- ] ----·-1----- 1--------- a strongly acidic solvent . In fact, all acids 0 0·1139 7·47 tend to become indistinguishable in strength 10 0·2920 7·27 in strongly basic solvents. This is known as 20 0·6809 7·085 25 1·008 6·998 the levelling effect of the solvent. 30 1 · ~69 6·915 Solvents may be protophilic, protogenic or 40 2·919 6·77 aprotic. If a solvent exhibits both proto­ 50 5·474 6·63 60 9·614 6·51 philic and protogenic properties it is termed a.mphiprotic. Examples of each form are *pH is here defined as pH = - log[H +], i.e. in protophilic solvents-ethers, ammines terms of concentration rather t han activity of the (basic substances); hydrogen ion. This is done for the sake of sim ­ plicity in t his and in all subsequent calculations in protogenic solvents- sulphuric acid this article. (acidic substances); THE THEORY OF ACIDS A.1'<D U.A.SES 93 amphiprotic solvents-water, acetic acid, strengths of what, in water, are normally alcohols; termed strong acids are to be determined, aprotic solvents-benzene, chloroform then solvents other than water have to be ('inert' substances). used. This necessitates the use of other Actually, these definitions involve a modern acidity sca.les, the most familiar of which is extension of the Bn'insted- Lowry theory, probably the Hammett acidity scale. which might be termed the 'auto-protolysis For example, the acids perchloric, hydro­ theory,' viz. that in a solventS in which the chloric and nitric appear equally strong in equilibrium water where the strongest acid that can exist 2S .= A++ B- is the hydroxonium ion, H 30 +, but in other solvents jheir strengths differ because of occurs, an acid is any substance whose dis­ variation in the ease of formation of the solution increases the concentration of A+ and solvated proton in that particular solvent. a base is any substance which increases the The basicity and dielectric constant* appear concentration of B- . to be th<: principal factors in the manifesta­ Thus when water ionizes, the equation may tion of acidity. be written as Since the extent of acidic and basic dis­ H 20 + H 20 <=' H 3 0 + + OR- sociation is influenced by the ability of the where one molecule of water is behaving as solvent to accept or donate protons, the same solute may be dissociated to widely different an acid and another as a base (amphiprotic). degrees in different solvents. For example, Similarly for acetic acid ammonia is not highly protonatccl in water, CH3COOH + CH 3COOH <=' CH 3COOH2 + + but glacial acetic acid, a solvent more proto­ CJ:I. 3COO- genic than water, induces extensive protolysis. one acetic acid is acting as a proton donor .::\TH3 -! C H.3COOH .= ~ H, - + CH3COO­ (acid) and the other as a proton acceptor Again, sulphuric acid is over four hundred (base). times as acidic in acetic acid as in water at Again, ammonia ionizes as follows: the same concentration. Such systems are ~1-J3 + NH 3 <=' NH,,+ + .::\TH2 - often referred to as 'super-acid systems.' and acids arc those substances which, in P erchloric acid may be dissolved in acetic liquid ammonia, increase the concentration of acid to give a solution containing CH3COOH2+ ions, and, as this ion can readily give up a NH4+. It follows then that a substance which proton to react with a base, the solution is ftmctions as an acid in one solvent docs not strongly acidic. On the other hand acetic necessarily react in this way in another acid may itself donate protons to a suitable solvent.
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