Ch. 14: Acids and BasesCh. 14: Acids and Bases

Dr. Namphol Sinkaset

Chem 152: Introduction to General Chemistry

I. Chapter OutlineI. Chapter Outline

I. Introduction

II. Acid/Base Definitions

III. Titrations

IV. Strong vs. Weak

V. The pH Scale

VI. Buffers

I. IntroductionI. Introduction

I. AcidsI. Acids

• Acids have characteristic properties. Taste sour. Dissolve many metals. Turn blue litmus paper red.

• Common acids include: hydrochloric, sulfuric, nitric, acetic, carbonic, and hydrofluoric.

I. Two AcidsI. Two Acids

I. BasesI. Bases

• Bases have characteristic properties. Taste bitter. Feel slippery. Turn red litmus paper blue.

• Common bases include: sodium hydroxide, potassium hydroxide, sodium bicarbonate, and ammonia.

II. Acid/Base DefinitionsII. Acid/Base Definitions

• There are several different definitions for acids and bases.

• What definition you use depends on what kinds of compounds you are studying and what’s convenient.

• We will cover the two most commonly used definitions.

II. The Arrhenius DefinitionsII. The Arrhenius Definitions

• An acid is a substance that produces H+ ions in aqueous solution.

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

• Note that these definitions are restricted to water-based solutions.

II. An Arrhenius AcidII. An Arrhenius Acid

• HCl is an example of an Arrhenius acid.

• Note that H+ always attaches to a water molecule to form H3O+, the hydronium ion.

• H+(aq) = H3O+

(aq)

II. An Arrhenius BaseII. An Arrhenius Base

• Sodium hydroxide is an example of an Arrhenius base.

II. Brønsted-Lowry DefinitionsII. Brønsted-Lowry Definitions

• An acid is a proton (H+ ion) donor.

• A base is a proton (H+ ion) acceptor.

• Notice that the focus in these definitions is on transfer of H+.

• Notice that there is no dependence on aqueous solutions, so this definition is more widely applicable.

II. Acid/Base PairsII. Acid/Base Pairs

• To use the Brønsted-Lowry definitions, you have to analyze an entire reaction and see what’s giving up H+ and what’s accepting the H+.

• Under this definition, acids and bases always occur together!

II. A Brønsted-Lowry AcidII. A Brønsted-Lowry Acid

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

• We see that the HCl gives up an H+; HCl is the acid.

• We see that H2O accepts an H+; H2O is the base.

II. A Brønsted-Lowry BaseII. A Brønsted-Lowry Base

NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

• We see that NH3 accepts an H+; NH3 is the base.

• We see that H2O gives up an H+; H2O is the acid.

II. Water is AmphotericII. Water is Amphoteric

• Notice in the last two slides that H2O was acting as a base in one and as an acid in the other.

• H2O is amphoteric, a substance that can act as either an acid or a base.

• Another example would be bisulfate, HSO4

-.

II. Conjugate Acid-Base PairsII. Conjugate Acid-Base Pairs

• Under Brønsted-Lowry: The acid loses H+

to become a conjugate base.

The base gains H+ to become a conjugate acid.

II. Conjugate Acid-Base PairsII. Conjugate Acid-Base Pairs

• The formulas of conjugate pairs differ by only one H+!

II. Practice ProblemII. Practice Problem

• 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)

III. Acid/Base TitrationIII. Acid/Base Titration

• When an acid reacts with a base, the product is always water and a salt.

• We can use the stoichiometry of the reaction to figure out the concentration of one if we know the concentration of the other.

• Titration is a technique in which a solution of known [ ] is reacted with another solution of unknown [ ].

III. A Typical TitrationIII. A Typical Titration

• Solution of known [ ] is added through a buret.

• An indicator tells you when to stop.

• At the equivalence point, moles acid = moles base.

III. Calculating the Unknown [ ]III. Calculating the Unknown [ ]

1) Use the volume added from the buret and the concentration to find moles of known.

2) Use the balanced equation to convert moles of known to moles of unknown.

3) Divide moles of unknown by the sample volume.

III. Sample ProblemIII. Sample Problem

• A 25.0-mL sample of sulfuric acid is titrated with a 0.225 M solution of sodium hydroxide. If it takes 21.27 mL to reach the endpoint, what is the molarity of the sulfuric acid solution?

IV. Acid/Base StrengthIV. Acid/Base Strength

• Different acids and bases have different strengths.

• There are actually more weak acids and bases than strong acids and bases.

• Acid/base strength is related to whether they are strong or weak electrolytes.

IV. Strong Acids and BasesIV. Strong Acids and Bases

• Strong acids and bases are strong electrolytes; they dissociate completely.

IV. Strong Acids and BasesIV. Strong Acids and Bases

IV. Weak Acids and BasesIV. Weak Acids and Bases

• Weak acids and bases are weak electrolytes; they do not dissociate completely.

IV. Indicating WeaknessIV. Indicating Weakness

• Equations showing weak acids or bases use a double arrow to indicate incomplete dissociation.

V. Water Reacts w/ Itself!V. Water Reacts w/ Itself!

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

V. Water Ion Product ConstantV. Water Ion Product Constant

• In pure water at 25 °C, there’s always a little H3O+ and OH- in equal amount.

• Specifically, [H3O+] = [OH-] = 1.0 x 10-7 M.

• When these concentrations are multiplied, you get the ion product constant for water, Kw.

• Kw = [H3O+][OH-]

• At 25 °C, Kw = 1.0 x 10-14.

V. Acidic/Basic SolutionsV. 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.

V. Sample ProblemV. Sample Problem

• 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?

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

acidity or basicity of a solution. pH < 7 is acidic; pH = 7 is neutral; pH

> 7 is basic.

V. pH is a log ScaleV. pH is a log Scale

• pH = -log [H3O+]

• Since it’s a log scale, a one unit change is actually a 10x change.

• log is a different type of math, so it has its own sig fig rule…

V. Sig Figs for logV. Sig Figs for log

V. Sample ProblemsV. Sample Problems

• Perform the following calculations.a) Calculate the pH of a solution in which the

hydronium concentration is 4.2 x 10-3 M.

b) Calculate the pH of a solution in which the hydroxide concentration is 7.89 x 10-8 M.

c) If the pH of a solution is 4.67, calculate the concentration of hydronium.

VI. Resisting Changes in pHVI. Resisting Changes in pH

• The only things that affect the pH are free H3O+ and OH-.

• If we can create a solution that “captures” any added H3O+ or OH-, then we can resist changes in pH.

• A solution that can do this is called a buffer.

VI. Example of a BufferVI. Example of a Buffer