Prentice Hall ©2004 Chapter 14 Aqueous Equilibria: Acids and Bases

Prentice Hall ©2004 Acid–Base Concepts01 Arrhenius Acid: A substance which dissociates in water to form hydrogen ions (H + ) in solution. HA(aq) + H 2 O(l)  H 3 O + (aq) + A – (aq) Arrhenius Base: A substance that dissociates in, or reacts with water to form hydroxide ions (OH – ). MOH(aq)  M + (aq) + OH – (aq)

Prentice Hall ©2004 Acid–Base Concepts02 Brønsted–Lowry Acid: Substance that can donate H + Brønsted–Lowry Base: Substance that can accept H + Chemical species whose formulas differ only by one proton are said to be conjugate acid–base pairs.

Prentice Hall ©2004 Acid–Base Concepts03

Prentice Hall ©2004 Acid–Base Concepts04

Prentice Hall ©2004 A Lewis Acid is an electron-pair acceptor. These are generally cations and neutral molecules with vacant valence orbitals, such as Al 3+, Cu 2+, H +, BF 3. A Lewis Base is an electron-pair donor. These are generally anions and neutral molecules with available pairs of electrons, such as H 2 O, NH 3, O 2–. The bond formed is called a coordinate bond. Acid–Base Concepts05

Prentice Hall ©2004 Acid–Base Concepts06

Prentice Hall ©2004 Acid–Base Concepts07 Problems 14.1,14.2 Write balanced equations for the dissociation of each of the following Brønsted–Lowry acids. (a) H 2 SO 4 (b) HSO 4 – (c) H 3 O + Problems Identify the Lewis acid and Lewis base in each of the following reactions: (a) AlCl 3 (s) + Cl – (aq) æ AlCl 4 – (aq) (b) SO 2 (aq) + OH – (aq) æ HSO 3 – (aq) (c) Ag + (aq) + 2 NH 3 (aq) æ Ag(NH 3 ) 2 + (aq)

Prentice Hall ©2004 Dissociation of Water01 Water can act as an acid or as a base. H 2 O(l) æ H + (aq) + OH – (aq) This is called the autoionization of water. H 2 O(l) + H 2 O(l) æ H 3 O + (aq) + OH – (aq)

Prentice Hall ©2004 Dissociation of Water02 This equilibrium gives us the ion product constant for water. K w = K c = [H 3 O + ][OH – ] = 1.0 x 10 –14 If we know either [H 3 O + ] or [OH – ] then we can determine the other quantity.

Prentice Hall ©2004 Dissociation of Water03 Problem 14.6 The concentration of OH – in a sample of seawater is 5.0 × M. Calculate the concentration of H 3 O + ions, and classify the solution as acidic, neutral, or basic. Problem 14.7 At 50°C the value of K w is 5.5 × What are the concentrations of H 3 O + and OH – in a neutral solution at 50°C?

Prentice Hall ©2004 pH – A Measure of Acidity01 The pH of a solution is the negative logarithm of the hydrogen ion concentration (in mol/L). pH = –log [H 3 O + ] pOH = –log [OH – ] pH + pOH = 14 Acidic solutions:[H + ] > 1.0 x 10 –7 M, pH 7.00 Neutral solutions:[H + ] = 1.0 x 10 –7 M, pH = 7.00

Prentice Hall ©2004 pH – A Measure of Acidity02 Problem 14.8 Calculate the pH of each of the following Problem 14.9 Calculate the concentrations of H 3 O + and OH – in each of the following solutions: (a) Human blood (pH 7.40) (b) A cola beverage (pH 2.8) Problem Calculate the pH of (a) M HClO 4 (b) 6.0 M HCl (c) 4.0 M KOH (d) M Ba(OH) 2 Problem Calculate the pH of a solution prepared by dissolving 0.25 g of BaO in enough water to make L of solution

Prentice Hall ©2004 pH – A Measure of Acidity03

Prentice Hall ©2004 Strength of Acids and Bases01 Strong acids and bases: are strong electrolytes that are assumed to ionize completely in water. Weak acids and bases: are weak electrolytes that ionize only to a limited extent in water. Solutions of weak acids and bases contain ionized and non-ionized species.

Prentice Hall ©2004 Strength of Acids and Bases02 HClO 4 HI HBr HCl H 2 SO 4 HNO 3 H 3 O + HSO 4 – HF HNO 2 HCOOH NH 4 + HCN H 2 O NH 3 ClO 4 – I – Br – Cl – HSO 4 – NO 3 – H 2 O SO 4 2– F – NO 2 – HCOO – NH 3 CN – OH – NH 2 – ACID CONJ. BASE Increasing Acid Strength

Prentice Hall ©2004 Strength of Acids and Bases03 Stronger acid + stronger base  weaker acid + weaker base Problems 14.4 Predict the direction of the following: HF(aq) + NO 3 – (aq) æ F – (aq) + HNO 3 (aq) NH 4 + (aq) + CO 3 –2 (aq) æ HCO 3 – (aq) + NH 3 (aq)

Prentice Hall ©2004 Acid Ionization Constants01 Acid Ionization Constant: the equilibrium constant for the ionization of an acid. HA(aq) + H 2 O(l) æ H 3 O + (aq) + A – (aq)

Prentice Hall ©2004 Acid Ionization Constants x 10 –4 4.5 x 10 –4 3.0 x 10 –4 1.7 x 10 –4 8.0 x 10 –5 6.5 x 10 –5 1.8 x 10 –5 4.9 x 10 – x 10 –10 HF HNO 2 C 9 H 8 O 4 (aspirin) HCO 2 H (formic) C 6 H 8 O 6 (ascorbic) C 6 H 5 CO 2 H (benzoic) CH 3 CO 2 H (acetic) HCN C 6 H 5 OH (phenol) F – NO 2 – C 9 H 7 O 4 – HCO 2 – C 6 H 7 O 6 – C 6 H 5 CO 2 – CH 3 CO 2 – CN – C 6 H 5 O – ACID K a CONJ. BASE K b 1.4 x 10 – x 10 – x 10 – x 10 – x 10 – x 10 – x 10 – x 10 –5 7.7 x 10 –5

Prentice Hall ©2004 Calculating Equilibrium Concentration in Solutions of Weak Acids

Prentice Hall ©2004 HA æ H + +A (M): (M):–x+x+x Equilib (M): 0.50–xxx Acid Ionization Constants04 I nitial C hange E quilibrium T able : Determine the pH of 0.50 M HA solution at 25°C. K a = 7.1 x 10 –4. Initial Change (aq) -

Prentice Hall ©2004 Acid Ionization Constants05 pH of a Weak Acid (Cont’d): 1. Substitute new values into equilibrium expression. 2. If K a is significantly (>1000 x) smaller than [HA] the expression (0.50 – x) approximates to (0.50). 3. The equation can now be solved for x and pH. 4. If K a is not significantly smaller than [HA] the quadratic equation must be used to solve for x and pH.

Prentice Hall ©2004 Acid Ionization Constants06 The Quadratic Equation: The expression must first be rearranged to: The values are substituted into the quadratic and solved for a positive solution to x and pH.

Prentice Hall ©2004 Acid Ionization Constants08 Percent Dissociation: A measure of the strength of an acid. Stronger acids have higher percent dissociation. Percent dissociation of a weak acid decreases as its concentration increases.

Prentice Hall ©2004 Base Ionization Constants01 Base Ionization Constant: The equilibrium constant for the ionization of a base. The ionization of weak bases is treated in the same way as the ionization of weak acids. B(aq) + H 2 O(l) æ BH + (aq) + OH – (aq) Calculations follow the same procedure as used for a weak acid but [OH – ] is calculated, not [H + ].

Prentice Hall ©2004 Base Ionization Constants x 10 –4 4.4 x 10 –4 4.1 x 10 –4 1.8 x 10 –5 1.7 x 10 –9 3.8 x 10 – x 10 –14 C 2 H 5 NH 2 (ethylamine) CH 3 NH 2 (methylamine) C 8 H 10 N 4 O 2 (caffeine) NH 3 (ammonia) C 5 H 5 N (pyridine) C 6 H 5 NH 2 (aniline) NH 2 CONH 2 (urea) C 2 H 5 NH 3 + CH 3 NH 3 + C 8 H 11 N 4 O 2 + NH 4 + C 5 H 6 N + C 6 H 5 NH 3 + NH 2 CONH 3 + BASE K b CONJ. ACID K a 1.8 x 10 – x 10 – x 10 – x 10 – x 10 –6 2.6 x 10 – Note that the positive charge sits on the nitrogen.

Prentice Hall ©2004 Diprotic & Polyprotic Acids01 Diprotic and polyprotic acids yield more than one hydrogen ion per molecule. One proton is lost at a time. Conjugate base of first step is acid of second step. Ionization constants decrease as protons are removed.

Prentice Hall ©2004 Diprotic & Polyprotic Acids02 Very Large 1.3 x 10 –2 6.5 x 10 –2 6.1 x 10 –5 1.3 x 10 –2 6.3 x 10 –8 4.2 x 10 –7 4.8 x 10 – x 10 –8 1 x 10 – x 10 –3 6.2 x 10 –8 4.8 x 10 –13 H 2 SO 4 HSO 4 – C 2 H 2 O 4 C 2 HO 4 – H 2 SO 3 HSO 3 – H 2 CO 3 HCO 3 – H 2 S HS – H 3 PO 4 H 2 PO 4 – HPO 4 2– ACID K a CONJ. BASE K b HSO 4 – SO 4 2– C 2 HO 4 – C 2 O 4 2– HSO 3 – SO 3 2– HCO 3 – CO 3 2– HS – S 2– H 2 PO 4 – HPO 4 2– PO 4 3– Very Small 7.7 x 10 – x 10 – x 10 – x 10 – x 10 –7 2.4 x 10 –8 2.1 x 10 –4 1.1 x 10 –7 1 x 10 –5 1.3 x 10 – x 10 –7 2.1 x 10 –2

Prentice Hall ©2004 Molecular Structure and Acid Strength01 The strength of an acid depends on its tendency to ionize. For general acids of the type H–X: 1. The stronger the bond, the weaker the acid. 2. The more polar the bond, the stronger the acid. For the hydrohalic acids, bond strength plays the key role giving: HF < HCl < HBr < HI

Prentice Hall ©2004 Molecular Structure and Acid Strength02 The electrostatic potential maps show all the hydrohalic acids are polar. The variation in polarity is less significant than the bond strength which decreases from 567 kJ/mol for HF to 299 kJ/mol for HI.

Prentice Hall ©2004 Molecular Structure and Acid Strength03 For binary acids in the same group, H–A bond strength decreases with increasing size of A, so acidity increases. For binary acids in the same row, H–A polarity increases with increasing electronegativity of A, so acidity increases.

Prentice Hall ©2004 Molecular Structure and Acid Strength04 For oxoacids bond polarity is more important. If we consider the main element (Y): Y–O–H If Y is an electronegative element, or in a high oxidation state, the Y–O bond will be more covalent and the O–H bond more polar and the acid stronger.

Prentice Hall ©2004 Molecular Structure and Acid Strength05 For oxoacids with different central atoms that are from the same group of the periodic table and that have the same oxidation number, acid strength increases with increasing electronegativity.

Prentice Hall ©2004 Molecular Structure and Acid Strength06 For oxoacids having the same central atom but different numbers of attached groups, acid strength increases with increasing central atom oxidation number. As shown on the next slide, the number of oxygen atoms increases the positive charge on the chlorine which weakens the O–H bond and increases its polarity.

Prentice Hall ©2004 Molecular Structure and Acid Strength07 Oxoacids of Chlorine:

Prentice Hall ©2004 Molecular Structure and Acid Strength08 Predict the relative strengths of the following groups of oxoacids: a) HClO, HBrO, and HIO. b) HNO 3 and HNO 2. c) H 3 PO 3 and H 3 PO 4.

Prentice Hall ©2004 Acid–Base Properties of Salts01 Salts that produce neutral solutions are those formed from strong acids and strong bases. Salts that produce basic solutions are those formed from weak acids and strong bases. Salts that produce acidic solutions are those formed from strong acids and weak bases.