Chapter 17: Chemistry of Acids & Bases

Chapter 17: Chemistry of Acids & Bases

“The end is near”

17.0 Objectives• 1. To recognize the properties of strong and

weak acids and strong and weak bases.

• 2. To become familiar with Arrhenius and Bronsted-Lowry theories of acids and bases.

• 3. Be able to calculate and understand the following quantities: [H+], [OH-], pH, pOH, Kw, Ka, and Kb.

• 4. Calculate the equilibrium concentrations of weak acids and bases.

• 5. Classify aqueous salt solutions as acidic, basic, or neutral.

• 6. Determine the concentration of unknown acid or base, or the Ka of a weak acid from a titration procedure.

Homework

• HW #1 - 11, 13, 15, 17, 19, 21– Conjugate Acids/Bases; pH

• HW #2 - 25, 27, 29, 31, 33, 37, 89– Equilibrium Constants and pKa

• HW#3 - 39, 41, 43, 45, 47, 49– pKb, Types of Reactions

• HW#4 - 51, 53, 55, 93– pH to calculate K values

• HW#5 - 57, 59, 63, 65, 97, 101– Equilibrium problems

17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT

• A. Definitions– 1. Arrhenius definition of acid

• Substance that contains H+

• Ex. HCl; H2SO4

– 2. Arrhenius definition of base• Substance that contains OH-

• Ex. NaOH, Ca(OH)2


• B. Properties of acids• Ionize when put into water• React with active metals (Group I, II) to

produce Hydrogen gas• Neutralized with bases• Have a sour taste


• C. Properties of bases• Ionize when put into water• Neutralized with acids• Have a bitter taste• Feel Slippery


• D. Strong acids - definition and examples– Strong acids dissociate completely in

aqueous solution:– Strong Acids

» HCl, HBr, HI, HNO3, HClO4, H2SO4

» “BrIClSONOClO 4 3 4”

– EX. HCl


E. Weak acids - definition • Weak acids are partly dissociated in

soln.• Treated as an equilibrium

• Ex. CH3COOH

– Species• H+

• H3O+

• [insert H-containing cation here]+


• F. Example 17.1 Write equations for the aqueous dissociation of the following weak acids: HPO4

-2, Fe(H2O)63+, HCN


• G. Strong bases - definition and examples– Strong bases completely dissociate in

aqueous solution:• Group I & MOST Group II Hydroxides

– LiOH, NaOH, KOH, RbOH, CsOH

– Ca(OH)2, Sr(OH)2, Ba(OH)2,

– Ex. NaOH


• H. Weak bases - definition and species– Two types of substances act like

weak bases in aqueous solution: • Nitrogen-containing compounds

– Ex. NH3

• Anions of acids– Ex. HCO3

-


• I. Example 17.2 Write equations for the water dissociation of the following weak bases: NO2

1-, CO32-, C2H5NH2

17.2 BRONSTED CONCEPT OF ACIDS AND BASES

• A. Arrhenius definitions – problems– Requires that the solution be aqueous

in order to work– Constrains acids and bases to contain

H+ and OH-


• B. Bronsted-Lowry acid - definition and example

• Acid = Proton Donor• “For every acid, there must be a base”

• Ex. HNO3


• C. Bronsted-Lowry base - definition and example– Base= Proton acceptor


• D. Conjugate acids and bases– Conjugate acid: Bronsted Base with an

extra proton (H+) attached• Every base has a conjugate acid

– Conjugate base: Bronsted acid minus a proton (H+)

• Every acid has a conjugate base


• E. Amphiprotic substances

• Substances that can ionize as either an acid or a base depending on the properties of the other species in soln.

• Ex. H2O– can act as a conjugate acid or a conjugate

base


• F. Example 17.3 Identify the Bronsted Lowry acids and bases and their conjugates in the following equations:

– NH4+(aq) + OH1-(aq) NH3(aq) +

HOH(l)

– N2H4(aq) + HOH(l) N2H51+(aq) +

OH1-(aq)


• G. Example 17.4 Write equations to show that HSO4

1- can act like an

amphiprotic substance.


• H. Polyprotic acids– Acids that give off more than 1 H+

when put into water• 1st Proton usually given off rapidly• Subsequent protons are given off with

increasing difficulty (stronger bases at each step)

17.3WATER AND THE pH SCALE

• A. Water - self-ionization, and the water dissociation constant, Kw

H2O + H2O --- H3O+ + OH-

Acid(1) Base(2) Acid(2) Base(1)

– For any sample of water molecules:• 2 H2O (l) <-----> H3O+

(aq) + OH- (aq)

•Kw = = 1.0 x 10-14


• B. Relationships for all aqueous solutions – In a neutral soln:

• [H3O+] = [OH-] = 1.0 x 10-7 M

– In an acidic soln: [H3O+] > [OH-] • [H3O+] > 1.0 x 10-7 M

• [OH-] < 1.0 x 10-7 M

– In a basic soln: [H3O+] < [OH-] • [H3O+] < 1.0 x 10-7 M

• [OH-] > 1.0 x 10-7 M


• C. The pH scale• Ranges from 0 to 14• 0~7 = ACID• 7~14 = BASE• 7 = Neutral


• D. Definition of pH and pOH• pH = -log [H+]

• [H3O+] = 10-pH

• Eqns. for pOH?

• pH + pOH = 14


• E. Example 17.5 Calculate the [H+], [OH1-], pH, and pOH for a. lemon juice with [H+]=5.0 x 10-3M, and b. blood with a pH of 7.4

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• A. Weak acids, Ka

• Equilibrium of an acid with conjugate base

• **REMEMBER** H2O (l) is not included!

HA + H2O <-----> H3O+ + A-

Ka =

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• B. Value

– Less than 1 because it is a weak acid (more reactant compared to product)

– As Ka increases the strength of acid increases (directly related)

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• C. pKa and value

• pKa = - log Ka

• As pKa decreases the strength of the acid increases (inversely related)

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• D. Example 17.6 In a solution prepared by

dissolving 0.100 mole of lactic acid per liter, [H+] = 3.7 x 10-3M. Calculate the Ka for lactic acid.

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• E. Example 17.7 Determine the [H+], pH and %

dissociation of carbonic acid, H2CO3 (Ka = 4.2 x 10-7) in a 0.100 M solution.

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• F. Approximation

– Small Ka values = few protons given off = very little dissociation

– If Ka is smaller than 10-3, “x” is negligible in an IRE where it is “-x”

• For quadratic approximations only!

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• Weak Bases, Kb

• Equilibrium of a base with conjugate acid

• **REMEMBER** H2O (l) is not included!

B- + H2O <-----> OH- + HB

Kb =

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• pKb and Kb values

• pKb = - log Kb

• As pKb decreases the strength of the base increases (inversely related)

• Large pKb = small pKa

• Ka x Kb = Kw = 1.0 x 10-14

• pKa + pKb = pKw = 14.0

• Think of Ka and Kb as opposites

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• I. Example 17.8 Pyruvic acid, HPy, has a Ka of

3.3 x10-3 and Butyric acid, HBut, has a Ka of 2.0 x 10-5. Which is the stronger acid, HPy or HBut? Which is the stronger base, But1- or Py1-?

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES

– J. Example 17.9 Determine the [H+], [OH1-], pH and pOH of a 0.100M solution of a weak base that has a Kb of 1.0 x 10-9.

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• K. Acid –Base Properties of salt

solutions1. salt – definition

– “A compound formed by replacing the hydrogen ions of an acid with metal ions”

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• 2. salt dissociation in water –

hydrolysis reaction– Many of the ionizations we have been

writing are hydrolysis

– Ex. Fe(H2O)6+3 weak acid

17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES• 3. interaction of ions with water

molecules– Anions that are conjugate bases of

strong acids are extremely weak bases– Alkali/Alkaline earth metals have no

effect


– 4. Example 17.10 For the following salts, determine whether the aqueous solution would be acidic or basic or neutral. Then write the net ionic equation for the interaction of each active ion with water. KClO4 Na2HPO4 NH4I

17.6 TYPES OF ACID-BASE REACTIONS

• 1. Strong acid and strong base – neutralization

• 2. weak acid and strong base• 3. weak base and strong acid• 4. weak base and weak acid

• Strong acid and strong base - neutralization

• Weak acid and strong base

• Weak base and strong acid

• Weak base and weak acid


• 5. Example 17.11 Write net ionic equations for the interactions between the following pairs of acids and bases: a. HNO2 and NaOH b. (CH3)2NH and HCl

c. HBr and KOH.


• 5. Acid Base Titrations– Complete dissociations

• No need to worry about Ka or Kb

– Strong Acid + Strong Base --> H2O + Salt

Equation– (#H+) Ma Va = Mb Vb (#OH-)

– Stoich will also work


• B. Equivalence point and general graph– Point in a titration at which one

reactant has been exactly consumed by adding the other reactant

– GRAPH:


– 6. Example 17.12 A chemist synthesizes a sample of nicotinic acid (HC6H4NO2 molar mass 123g/mole). To determine the purity of the sample, a 0.450g portion of the acid is titrated with 0.100M NaOH. A total volume of 35.8mL of NaOH is required to neutralize the nicotinic acid according to the following:

• HC6H4NO2(aq) + OH1-(aq) HOH(l) + C6H4NO21-(aq)

• Determine the percentage of nicotinic acid in the original sample.


– 7. Example 17.13 If 50.00mL of 1.000M HCl is titrated with 1.000M NaOH, find the pH of the solution after the following volumes of NaOH have been added: 25.00mL, 50.00mL, and 50.01mL.

Chapter 17: Chemistry of Acids & Bases

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