Acid-base equilibrium pH scale, weak acids bases, salt of weak acid and bases

pH scale pH is the measure of the acidity or alkalinity of a solution pH is a measurement of the concentration of hydrogen ions in a solution. So, low pH values are associated with solutions with high concentrations of hydrogen ions, while high pH values occur for solutions with low concentrations of hydrogen ions.

pH scale The pH scale is an inverse logarithmic representation of hydrogen proton (H+) concentration. The pH of solution was defined by Danish chemist S. P. L. Sørensen as pH = − log10[H + ]

pH scale Example 1 Calculate the pH of a 2.0 x 10-3 M solution of HCl.

pH scale Solution [H+] = 2.0 x 10-3 M pH = − log10[H + ] = − log10[2.0 x 10-3 M ] = 3 - log10 2.0 = 3 – 0.30 = 2.7

pH scale There is also pOH, in a sense the opposite of pH, which measures the concentration of OH− ions, or the alkalinity. A similar definition is made for the hydroxyl ion concentration pOH = − log10[OH- ]

pH scale -log Kw = - log [H+][OH-] = -log [H+] – log [OH-] pKw = pH + pOH edit

pH scale PLEASE NOTE THAT AT 25˚C [H+] x [OH-] = 10-14 pH + pOH = 14

pH scale Example 2 Calculate the pOH and the pH of a 5.0 x 10-2 M solution of NaOH

pH scale Solution Method 1 [OH-] = 5.0 x10-2 M pOH = - log(5.0 x10-2 ) = 1.3

pH scale pH + pOH = 14 Solution continued Since pH + 1.3 = 14.00

pH scale [H+] x [OH-] = 10-14 Solution Method 2 Since [H+] = 1.0 x 10-14 = 2.0 x 10-13 5.0 x 10-12 pH = - log (2.0 x 10-13) = 13 – log 2.0 = 13 – 0.30 = 12.7 [H+] x [OH-] = 10-14

pH scale When [H+] = [OH-] the solution is NEUTRAL. pH@pOH of 7 When [H+] < [OH-] the solution is ALKALINE pH@pOH greater than 7 When [H+] > [OH-] the solution is ACIDIC pH@pOH less then 7

pH scale

Weak Acid A weak acid is an acid that does not completely donate all of its hydrogens when dissolved in water. These acids have higher pH compared to strong acids, which release all of their hydrogens when dissolved in water.

Weak Acid The acidity constant for acetic acid at 25oC is 1.75 x 10-5 [H+][OAc-] = 1.75 x 10-5 [HOAc] When acetic acid ionizes, it dissociates to equal portion of H+ and OAc- by such an amount will always be equal 1.75 x 10-5 The equilibrium concentrations of reactants and products are related by the Acidity constant expression (Ka) [H+][OAc-] = Ka

Weak Acid example 3 Calculate the pH of a 1.00 x 10-3 M solution Of Acetic Acid

Weak Bases A weak base is a chemical base that only partially ionize in water.

Weak Acid and Weak Base Example 4 Calculate the pH and pOH for a 1.00 x 10-3 M solution of ammonnia. Kb ,basicity constant is 1.75 x 10-5 at 25oC

Salts of Weak Acids and Bases The salt of a weak acid for example NaOAc is strong electrolyte, like all salt and completely ionizes. In addition, the anion of the salt of a weak acid is a Brønsted base which will accept protons. It partially hydrolyzed in water to form hydroxide ion and the corresponding undissociated acid.

Salts of Weak Acids and Bases If the salt hydrolyzes that salt is consider as a weak base. The weaker the conjugate acid, the stronger the conjugate base, that is, the more strongly the salt will combine with a proton, as from the water , to shift the ionization to the right.

Salts of Weak Acids and Bases Example 5 Calculate the pH of a 0.25 M solution of ammonium chloride.

Salts of Weak Acids and Bases Solution. Write the equilibria NH4Cl NH4+ + Cl- (ionization) NH4+ + H20 ↔ NH4OH + H+ (hydrolysis) (NH4+ + H20 ↔ NH3 + H30 + )

Salts of Weak Acids and Bases [NH4OH][H+ ] = Ka = Kw = 1.0 x 10 -14 NH4 Kb 1.7 x 10-5 = 5.7 x 10-10 Let x represent the concentration of [NH4OH] and [H+ ] at equilibrium. Then at equilibrium, [NH4OH] = [H+ ] = x [NH4+] = CNH4+ - x = 0.25 - x

Salts of Weak Acids and Bases Since CNH4+ » Ka , neglect x compared to CNH4+ Then, (x)(x) = 5.7 x 10^-10 0.25 X = √ 5.7 x 10^-10 x 0.25 = 1.2 x 10-5 M The NH4OH formed is undissociated and does no contribute to the pH [H+] = 1.2 x 10-5 M pH = - log (1.2 x 10-5 M) = 5 – 0.08 = 4.92

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