Ch. 16: Acids and Bases

Dr. Namphol Sinkaset Chem 201: General Chemistry II

I. Chapter Outline I. Introduction II. Acid/Base Definitions III. Acid Strength & Molecular Structure IV. Acid Strength, Ka, and p Functions V. Weak Acid Equilibria VI. Basic Solutions VII. Ions and Salts – Acid or Base? VIII. Polyprotic Acids

I. Acids and Bases • Acid-base reactions are a huge

category within chemistry. • Acids have a sour taste, dissolve many

materials, and neutralize bases. • Bases have a bitter taste, a slippery

feel, and the ability to neutralize acids. • We have found many uses for acids and

bases.

I. Common Acids

I. Some Acid Structures

I. Carboxylic Acids

I. Common Bases

I. Household Bases

II. Acid/Base Definitions • There are actually many different

definitions for acids and bases. • Which definition we use depends on

what we are trying to describe. • We cover 3 definitions: Arrhenius Brønsted-Lowry Lewis

II. Arrhenius Definitions

• An Arrhenius acid is a substance that produces H+ (H3O+, hydronium) ions in aqueous solution.

• An Arrhenius base is a substance that produces OH- ions in aqueous solution.

• Under this definition, acids and bases neutralize each other via H+

(aq) + OH-(aq)

H2O(l).

II. Arrhenius Acid/Base

II. Strong vs. Weak

II. Brønsted-Lowry Definitions

• These definitions are more widely applicable. • A Brønsted-Lowry acid is a proton donor. • A Brønsted-Lowry base is a proton acceptor. • Why “proton” donor?* • Notice these definitions do not depend on

aqueous solutions.

II. Brønsted-Lowry Acids/Bases

• Under this definition, acids and bases always occur in pairs. HCl(aq) + H2O(l) H3O(aq) + Cl-(aq) NH3(aq) + H2O(l) NH4

+(aq) + OH-

(aq) • Substances that can act as either an

acid or base are called amphoteric.

II. Conjugate Acid/Base Pairs • Under the Brønsted-Lowry definitions, in an

acid/base reaction, the acid becomes a base, and the base becomes an acid.

• Products are said to be “conjugates” of the reactants.

II. Conjugate Acid/Base Pairs

• Conjugate acid-base pairs are two substances related to each other by a transfer of a proton.

II. Practice Problem 16.1

• Identify the conjugate acid-base pairs in the reactions below.

HNO3(aq) + H2O(l) H3O+

(aq) + NO3-(aq)

C5H5N(aq) + H2O(l) C5H5NH+(aq) + OH-

(aq)

II. Lewis Definitions

• The Lewis definitions of acids and bases are very broad and widely applicable.

• The focus is on electron pairs. • A Lewis acid is an electron pair acceptor. • A Lewis base is an electron pair donor.

II. Lewis Acids

• The Lewis definition of an acid allows substances w/out H to be classified as acids.

• A Lewis acid has an empty orbital (or can rearrange e-’s to create an empty orbital) that can accept an e- pair.

• e.g. BF3 + :NH3 F3B:NH3

II. Example Lewis Acids

II. Cationic Lewis Acids

• Small, highly charged metal ions have lost e-’s and therefore can act as Lewis acids.

III. Acid Strength & Structure

• Why are some acids strong and some acids weak?

• Depends on structure and composition of the acid.

• We examine the factors that contribute to acid strength in binary acids and oxyacids.

III. Binary Acids

• The strength of a binary acid depends on bond polarity and bond strength. The H must have

the partial positive charge. Why?* Weaker bond

leads to greater acidity. Why?*

III. Binary Group 16/17 Acids

• The combined influence of polarity and bond strength can be seen in the Group 16 and Group 17 binary acids.

III. Oxyacids

• An oxyacid (a.k.a. oxoacid) has the general form H-O-Y- in which Y is some atom which may or may not have additional atoms bonded to it.

• Oxyacid strength depends on the electronegativity of Y and the number of O atoms attached to Y.

III. Electronegativity of Y

• The more electronegative Y is, the more polar and weaker the O-H bond becomes.

• If the O kicks off the H as H+, it can claim both electrons in the bond.

III. # of O Atoms on Y • More O atoms draw electron density away

from Y, which draws electron density from the O-H bond, leading to greater acidity.

IV. Acid Strength • Recall the difference between a strong

electrolyte and a weak electrolyte. • Strong and weak acids are defined

accordingly. A strong acid completely ionizes in solution. A weak acid partially ionizes in solution.

• Thus, acid strength depends on the equilibrium: HA(aq) + H2O(l) H3O+

(aq) + A-(aq)

IV. Strong Acids

• Strong acids ionize completely; thus, a 1.25 M HCl solution will have [H3O+] = 1.25 M and [Cl-] = 1.25 M.

IV. Complete Ionization

• All of the strong acid molecules split into H+ and A-.

IV. Weak Acids

• Weak acids establish an equilibrium between their ionized and nonionized forms.

IV. Partial Ionization

• Only some of the weak acid molecules will split into H+ and A-.

IV. Weak Acids

• The strength of a weak acid depends on the attraction between H+ and A-, i.e. how badly HA wants to stay together.

• In general, the stronger the acid, the weaker the conjugate base. Why?*

HA(aq) + H2O(l) H3O+

(aq) + A-(aq)

IV. Weak Acid Strength

• Is there a way to characterize the strength of a weak acid?

• Since weak acids set up an equilibrium, we can calculate equilibrium constants to see how much they dissociate.

• Of course, higher values of K would mean more dissociation, more H3O+, and thus a stronger acid.

IV. Acid Dissociation Constants

• Since these are acid equilibria, we give the constants a new name: acid dissociation constants, Ka.

• For the generic weak acid equilibrium HA(aq) + H2O(l) H3O+

(aq) + A-(aq):

IV. Sample Problem 16.2

• Write the acid dissociation equation for (CH3)2NH2

+ and its Ka expression.

IV. Sample Problem 16.3

• Which acid is stronger: HF with Ka of 3.5 x 10-4 or HCN with Ka of 4.9 x 10-10? Which conjugate base is stronger?

IV. Water Reacts w/ Itself!

• We said before that water is amphoteric; it can also react with itself in an acid/base reaction.

• This is called an autoionization reaction.

IV. Water Autoionization • As you see, this is an equilibrium, so we can

write an equilibrium expression which we call the ion product constant for water, Kw.

• Kw is sometimes called the water dissociation constant.

• At 25 °C, Kw = 1.00 x 10-14.

IV. Pure Water

• In pure water, the autoionization is the only source of H+ and OH-.

• When [H+] = [OH-], the solution is neutral.

IV. Acidic/Basic Solutions

• In an acidic solution, additional H3O+ ions exist, increasing [H3O+].

• In a basic solution, additional OH- ions exist, increasing [OH-].

• However, in all aqueous solutions, the product of hydronium and hydroxide concentrations always equals Kw.

IV. Sample Problem 16.4

• Calculate the [H3O+] concentration of a solution that has [OH-] = 1.5 x 10-2 M at 25 °C. Is the solution acidic or basic?

IV. The pH Scale • pH is simply another way to specify the

acidity of a solution. pH = -log [H3O+] pH < 7 is acidic; pH = 7 is neutral; pH > 7 is basic.

• The log function has the following significant figure rule:

IV. The p Function

• The p function can be used on anything. • A p function is the mathematical

operation of taking the negative log. pOH = -log [OH-] pKa = -log Ka pNa+ = -log [Na+]

IV. pH/pOH Relationship

IV. Sample Problems 16.5

a) Calculate the H3O+ concentration for a solution that has a pH of 8.37.

b) Calculate the H3O+ concentration for a solution with a pOH of 3.42.

c) Calculate the pKa for phenol which has a Ka = 1.3 x 10-13.

V. Weak Acid Equilibria • Equilibrium problems involving weak acids

are similar to calculations we’ve done before. • Problems look for either Ka, pH, or equilibrium

concentrations. • In these problems, we ignore the contribution

from the autoionization of water. • However, they are more complicated

because information can be given in a variety of different ways!

V. Sample Problem 16.6

• In a 0.100 M solution of lactic acid, the pH is 2.44 at 25 °C. Calculate the Ka and pKa of lactic acid at this temperature.

V. Sample Problem 16.7

• Nicotinic acid is a weak acid w/ pKa = 4.85. Calculate [H3O+] and the pH of a 0.050 M solution of nicotinic acid.

V. Percent Ionization

• Instead of characterizing a weak acid by its Ka, we can calculate how much it ionizes.

• The concentration of ionized acid is simply equal to the [H+] at equilibrium.

V. Sample Problem 16.8

• Find the % ionization of a 0.200 M acetic acid solution if its pKa = 4.74.

V. Sample Problem 16.9

• In a 0.0100 M solution of butyric acid at 20 °C, the acid is 4.0% ionized. Calculate the Ka and pKa of butyric acid at these conditions.

V. Mixtures of Acids

• If there are multiple acids in solution, then there are multiple sources of H3O+.

• If one of the acids is strong, it will be the major contributor, so much so that we can ignore contributions from others. Why?*

HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)

HCOOH(aq) + H2O(l) H3O+(aq) + HCOO-

(aq)

V. Mixtures of Weak Acids

• If we have a mixture of weak acids, we need to examine the Ka’s. If the Ka’s differ by more than a factor of

several hundred, then we can just calculate based on the strongest acid.

• If the Ka’s are too close together, we must solve two weak acid equilibria! Start with the strongest weak acid and use

that result in the second weak acid equilibrium.

V. Sample Problem 16.10

• Calculate the pH of a solution that is 0.100 M in acetic acid, CH3COOH (Ka = 1.8 x 10-5), and 0.200 M in formic acid, HCOOH (Ka = 1.8 x 10-4).

VI. Basic Solutions

• Just like acidic solutions, there are strong bases and weak bases.

• Everything we learned about weak acids applies to weak bases.

• The two systems are analogous to each other.

VI. Strong Bases

• Strong bases ionize completely. • Strong bases are typically ionic compounds

containing the hydroxide anion.

VI. Weak Bases

• Weak bases typically do not produce OH- by partially ionizing.

• Weak bases produce OH- by pulling a proton off water.

B(aq) + H2O(l) BH+(aq) + OH-

(aq)

VI. Strong/Weak Bases

VI. Weak Bases

• The strength of a weak base depends on its base ionization constant, Kb.

• For the generic weak base B(aq) + H2O(l) BH+

(aq) + OH-(aq):

VI. Common Weak Bases

VI. Weak Base Structures

• Weak bases tend to have lone pair e-’s that can accept a proton.

VI. Weak Base Problems

• The method of solving weak base problems is no different than the method of solving weak acid problems!

• Instead of Ka, you use Kb. • To get to pH, remember that you can go

through pOH.

VI. Sample Problem 16.11

• The pain reliever morphine is a weak base. In a 0.010 M morphine solution, the pH is 10.10. Calculate the Kb and pKb of morphine.

VI. Sample Problem 16.12

• The pKb for pyridine is 8.77. What’s the pH of a 0.010 M aqueous solution of pyridine?

VII. Ions as Weak Acids/Bases • Some ions can act as either weak acids

or weak bases. e.g. NH4

+ and CH3COO- • These ions must be introduced into a

solution as a salt. e.g. NH4Cl and CH3COONa

• These ionic salts ionize, and then the weak acid/base sets up its equilibrium.

VII. Anions as Weak Bases • Any anion can be thought of as the conjugate

base of an acid. • Anions that are conjugate bases of weak

acids are themselves weak bases. • Anions that are conjugate bases of strong

acids are pH neutral.

VII. Finding pH of Ionic Solutions

• If we know that a salt will dissociate and form a conjugate base in solution, we can calculate the pH of the solution.

• Problem: We generally will have a list of Ka’s, but what about the Kb of the conjugate?

VII. Conjugate Ka/Kb Pairs

VII. pKa/pKb Relationship

VII. Sample Problem 16.13

• What’s the pH of a 0.10 M NaNO2 solution? Note that Ka for nitrous acid is 7.1 x 10-4.

VII. Cations as Weak Acids

• When cations go into aqueous solutions,we need to examine whether or not they will set up an equilibrium. Cations of strong bases do nothing and are

thus pH neutral. Cations that are conjugate acids of weak

bases will establish an acid equilibrium, e.g. NH4

+ is the conjugate acid of NH3. Small, highly-charged metal cations form

weakly acidic solutions.

VII. The Case of Al3+

• Al3+(aq) will form Al(H2O)6

3+ which will establish an acid equilibrium.

VII. pH of Salt Solutions

• To determine whether a salt solution will be acidic, basic, or neutral, we need to consider the nature of the cation and anion.

• There are 4 possibilities: Neither cation nor anion acts as acid or base. Cation acts as acid, anion is neutral. Anion acts as base, cation is neutral. Cation acts as acid, and anion acts as base.

VII. Analyzing Salt Solutions • Use the following steps to determine

whether an aqueous solution of a salt will be acidic, basic, or neutral.

1) Break up salt into its cation and anion. 2) Ask yourself whether the cation can

donate a proton or whether it is small and highly charged. If so, it is a weak acid.

3) Ask yourself whether the anion can accept a proton. If so, it is a weak base.

4) Consider the combined effect of having the cation and anion in solution.

VII. Sample Problem 16.14

• For each compound, predict whether its 0.1 M solution in water will be acidic, basic, or neutral.

a) NaNO2 b) KCl c) NH4Br d) Fe(NO3)3

e) NH4CN

VIII. Acids w/ More than one H+

• Some acids have more than one acidic proton; these are called polyprotic acids.

• Generally, the Ka of the 2nd proton is much smaller than the 1st, so we generally just solve for the 1st ionization. Exceptions: H2SO4 and when Ka’s are within

a few hundred of each other. For exceptions, it’s a double equil. problem!

VIII. Some Polyprotic Acids

VIII. Sample Problem 16.15

• What is the pH and [SO42-] of a 0.0075

M sulfuric acid solution if Ka2 = 0.012 for sulfuric acid?