The lewis concept

The Lewis Concept (G.N. Lewis 1923)

Lewis acid - species which can accept electrons (Electron acceptor)

Lewis base - species which can donate electrons (Electron donor).

A lewis acid - base reaction results in the formation of a coordinate covalent bond between the two species. Generally, lewis acid is identified by the presence of vacant orbitals which can accommodate electrons whereas a lewis base will have lone pair of electrons available for shairing.

Thus, lewis definition extends beyond reactions which involve protons. Its most important application is in describing complex ion formation e.g.

 The strength of the base is in a given solvent depends on the nature of the acid used as a reference.

Note:Brfnsted-Lowry base (accepts a proton by donating e’s) is also a Lewis base BUT a Brfnsted -Lowry acid is not necessarily a Lewis acid.For reactions which occur in aqueous solution, Brfnsted -Lowry definitions are more applicableIn systems ;which do not involve protons, Lewis definitions is more valuable

Classes of solvents

w.r.t Brfnsted-Lowry acid-base properties, solvents are either aprotic or amphiprotic. Aprotic solvents are neither acidic nor basic , e.g. CCl4, C6H6

Amphiprotic solvents - can act as both proton acceptors and proton donors. They range from predominantly acidic to predominantly basic. Water and ethanol are neither strongly acidic or strongly basic.

All Amphiprotic solvents undergo Autoprotolysis or self ionization reaction:(spontaneous reaction of molecules to give pair of ions)

Where SH2+ is solvated proton;

 Self ionization (Autoprotolysis) of water

Water exhibits both acidic and basic properties:

The reverse reaction proceeds to a greater extent than the forward reaction and therefore the equilibrium lies to left.

K = [H3O+][OH-] = Kw

Kw – ion product of water

At 25oC, Kw = 10-14. In Pure water, [H3O+]  = [OH-] = 10-7

Note: self ionization of water contributes to the hydrogen (hydronium ion) and hydroxide ion concentration in solution but the effect depends on the concentration and the nature (strong/ weak acid/ base) of the solute. In General, if [solute] > 10-6 M, the contribution of either OH- or H3O+ is negligible

pH Scale;Potential of hydrogen ion” is the – ve potential of the hydronium ion

pH = -log [H3O+]

Examples

Calculate pH of 0.005 M HNO3 acid,

pH = -log [H3O+] = -log 0.005 = 2.3

 Calculate [H3O+] for HCl solution at pH 3.5:

3.5 = -log [H3O+]; [H3O+] = antilog -3.5 = 3.16 X 10-4 M

Similarly, pOH = -log [OH-]

Thus

Kw = [H3O+][OH-] = 10-14 Þ [H3O+] = [OH-] = 10-7

-Kw = -([H3O+][OH-])

-log Kw = -log [H3O+] + -log [OH-]

Since -log [H3O+] = pH and -log [OH-] =  pOH

pKw = pH + pOH =14

hence for a neutral solution, pH = pOH = 7

Note:

Self ionization of water contributes to the amount of H3O+ and OH-; The overall effect of self ionization of water depends on whether  the acid or base is a strong or weak electrolyte

Consider 0.01 M HNO3, both the HNO3 and H2O contribute H3O+

 

But nitric acid (strong acid) dissociates completely to produce 0.01 M H3O+. By Le Chatelier’s principle, the amount of H3O+ (from autoprotolysis of water) will be diminished at equilibrium and will be expected to be less than 10-7 M. Thus the contribution of H3O+ from water will be negligible.

Numerically,

For 0.01 M HNO3 , [H3O+] = 0.01 M; At equilibrium,

 Kw = [H3O+][OH-] = 10-14

Thus [OH-] <<< 10-7 ([OH-]  in pure water.

Since the only source of OH- is self ionization water, [H3O+] contributed by water = 10-12 M. Thus the contribution of [H3O+] contributed by self ionization of water is negligible.

Similarly, for strong base e.g. 0.01 M NaOH, the contribution of self ionization of water to [OH-] is negligible as it is less than 10-7 M

In 0.01 M NaOH, [OH-] = 0.01 M

Thus [OH-] from self ionization of water = 10-12 M

Generally, if the dissolved acid or base contributes a [H3O+] or [OH-] ³ 10-6 M, the contribution of water to the total concentration of either H3O+ or OH- is negligible.

Examples

1.The pH of rainwater is found to be 4.35 while that of ammonia solution is found to be 11.28. calculate the concentration of the hydronium and hydroxide ions in (a) rainwater and (b) ammonia solution

2. What volume of water is required to raise the pH of 1.00 L of HCl solution from 2.5 to 3.10

3. The solubility of Ca(OH)2 is 0.165 g/ 100 ml of solution at 20 °C. What is the pH of the saturated solution of calcium hydroxide at 20 °C.

4. Calculate the pH of a solution that is 3.00% KOH by mass and has a density of 1.0242 g/ml

Weak acids and Weak bases

 Partially ionized in solution

For weak acid, HA, which dissociates according to

The equilibrium constant, known as acid ionization constant, Ka, is given by;

For weak base, BH

For acetic acid, CH3COOH

pKa = -log Ka = -log 1.8 X 10-5 = 4.74

The larger the Ka , the farther the equilibrium lies to the right and greater the concentration of the anion produced.

Examples

1. A 0.250 M aqueous solution of butyric acid showed a pH of 2.72. Determine the acid dissociation constant of the organic acid.

 solution

Init. Conc  0.250 M               -  -       -

D       -x M  x M                       x M

Equi.Conc  (0.250-x) M                            x M     x M

pH = -log [H3O+]

[H3O+] = antilog 2.72 = 1.9 X 10-3 = x

Qtn Cocain (C17H21O4N) is an alkaloid, with a characteristic bitter taste, has a solubility of 0.17g/ 100 ml. If the pH of the saturated solution is 10.08, calculate the Kb value for cocaine.

References

 ·  Christian Laurence and Jean-François Gal "Lewis Basicity and Affinity Scales : Data and Measurement" Wiley, 2009.

 ·  Lewis, G.N., Valence and the Structure of Atoms and Molecules (1923) p. 142.

 ·  Miessler, L. M., Tar, D. A., (1991) p166 - Table of discoveries attributes the date of publication/release for the Lewis theory as 1923. 

March, J. “Advanced Organic Chemistry” 4th Ed. J. Wiley and Sons, 1992: New York

  ·  Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. 

·  Jensen, W.B. (1980). The Lewis acid-base concepts : an overview. New York:.

·  Yamamoto, Hisashi (1999). Lewis acid reagents : a practical approach. New York: Oxford University Press.