Chapter 17 Acid–Base (Proton Transfer) Reactions

Arrhenius Acid–Base Theory An acid ( HCl) is a substance which produces hydrogen ions in water solution. The properties of an acid is the properties of the hydrogen ions. A base (NaOH) is a substance which produces hydroxide ions in water solution. The properties of a base are the properties of the hydroxide ions. The net reaction between a strong acid and a strong base is :  H+ + OH-  H2O   The Arrhenius concept of acids and base is limited because it applies only to aqueous solutions.

Arrhenius Acid–Base Theory An example of an Arrhenius acid: Gaseous hydrogen chloride dissolved in water: HCl(g) H+(aq) + Cl–(aq) An example of an Arrhenius base: Solid sodium hydroxide dissolved in water: NaOH(s) Na+(aq) + OH–(aq)

Arrhenius Acid–Base Theory Properties of an acid must be due to properties of H+. Properties of a base must be due to properties of OH–. Thus the cause of the sour taste of acids is the H+ ion, the cause of the bitter taste of bases is the OH– ion. Other characteristic properties of acids and bases are also due to the H+ and OH– ions in water solutions.

Arrhenius Acid–Base Theory

Brønsted–Lowry Acid–Base Theory An acid is a proton donor. A base is a proton acceptor. An acid-base reaction is a proton-transfer reaction in which the proton is transferred from the acid to a base with formation of another acid and base.

Brønsted–Lowry Acid–Base Theory The base formed when the acid has donated a proton is called the conjugate base of the acid. Acid A ↔ H+ + Conjugate base of acid A The sign ↔ is used to show that the reaction is reversible. The stronger the acid, the weaker the conjugate base, and the weaker the acid, the stronger the conjugate base.

Brønsted–Lowry Acid–Base Theory An acid base reaction is a proton transfer reaction in which a proton is transferred from a stronger acid to a stronger base with formation of a weaker acid and weaker base. Stronger Acid1+ Stronger Base2↔Weaker Acid2+ Weaker Base1 HNO3 + NH3 ↔ NH4+ + NO3- HCl + CH3COO- ↔ HCH3COO + Cl-

Brønsted–Lowry Acid–Base Theory H3O+ is called hydronium ion. The conjugate base of acid HNO3 is NO3-

Brønsted–Lowry Acid–Base Theory Water which can behave as a base in one case and an acid in another is said to be amphoteric.

Brønsted–Lowry Acid–Base Theory

Conjugate Acid–Base Pairs B + HA HB+ + A– base acid acid base proton proton proton proton remover source source remover Conjugate Acid–Base Pair Two species that transform into each other by gain or loss of a proton, H+. HB+ and B and HA and A– are conjugate acid–base pairs

Conjugate Acid–Base Pairs

Conjugate Acid–Base Pairs

Lewis Acid–Base Theory Lewis Theory of Acids and Bases Acid Electron-pair acceptor. Base Electron-pair donor.

Lewis Acid–Base Theory

Relative Strengths of Acids & Bases

Predicting Acid–Base Reactions

Autoionization of water The Water Equilibrium Autoionization of water

Kw is the water constant or equilibrium constant for water The Water Equilibrium H2O(l) H+(aq) + OH–(aq) Kw = [H+] [OH–] = 1.0 × 10–14 Kw is the water constant or equilibrium constant for water If [H+] = [OH–] = x Kw = [H+] [OH–] = 1.0 × 10–14 = x2 x = = 10–7 moles/liter

The Water Equilibrium For water or water solutions: If [H+] = [OH–] = 10–7 M, the solution is neutral. If [H+] > [OH–], the solution is acidic. If [H+] < [OH–], the solution is basic.

The Water Equilibrium Example: What is the hydrogen ion concentration in a solution of 10–4 M sodium hydroxide in which the hydroxide ion concentration is 10–4 M? Is the solution acidic or basic? Solution: GIVEN: [OH–] = 10–4 M WANTED: [H+] EQUATION: Kw = [H+] [OH–] = 10–14 Since [H+] = 10–10 M< [OH–] = 10–4 M, the solution is basic

pH and pOH By definition pH and pOH are given by pH ≡ -log [H3O+] pOH ≡ -log [OH-]   [H3O+] ≡ antilog(-pH) ≡ 10-pH [OH-] ≡ antilog(-pOH) ≡ 10-pOH

pH and pOH What is the pH of a solution with [H+] = 10–5 M? Solution: pH = – log [H+] = – log 10–5 = 5 What is the [OH–] of a solution with pOH = 6? [OH–] = antilog (–pOH) = antilog (–6) = 10–6 M

pH and pOH Kw = [H+] [OH–] = 1.0 × 10–14 [H+] [OH–] = 1.0 × 10–14 – log ([H+] [OH–]) = – log (1.0 × 10–14) – log ([H+] [OH–]) = 14 – log [H+] + (– log [OH–]) = 14 pH + pOH = 14

pH and pOH

pH and pOH Example: The hydrogen ion concentration of a 10–3 M HCl solution is 10–3 M. What are the pH, pOH, and [OH–] of the solution? Solution: pH = – log [H+] = – log 10–3 = 3 pH + pOH = 14 pOH = 14 – pH = 14 – 3 = 11 [OH–] = antilog (–pOH) = antilog (–11) = 10–11 M

pH and pOH A solution is neutral if [H+] = 10–7 M A solution is acidic if [H+] > 10–7 M A solution is basic if [H+] < 10–7 M Using pH = – log [H+] and pOH = – log [OH–], A solution is neutral if pH = 7 A solution is acidic if pH < 7 A solution is basic if pH > 7

Significant Figures and Logarithms In a logarithm, the digits to the left of the decimal are not counted as significant figures. Counting significant figures in a logarithm begins at the decimal point.

pH and pOH Example: The hydrogen ion concentration of a solution is 2.7 × 10–6 M. What are the pH, pOH, and hydroxide ion concentration? Solution: pH = – log [H+] = – log (2.7 × 10–6) = – log (10–6) – log (2.7) = = 6 – log (2.7) = 5.57 (2 significant figures) pH + pOH = 14.00 pOH = 14.00 – pH = 14.00 – 5.57 = 8.43 [OH–] = antilog(–pOH) = antilog(–8.43) = 10–8.43 M =3.7 × 10–9 M

pH and pOH

pH and pOH

pH and pOH Measurement of pH

Homework 15, 17, 21, 23, 39, 41, 55, 59, 64.