AQUEOUS EQUILIBRIAACIDS AND BASES

Chapter 4.4 & 18.1-5 Silberberg

Understanding these Concepts

See the Learning Objectives on page 821.

Understand these Concepts: 18.1-20, 23-25.

Master these Skills: 18.1-5, 6-10, 12-14, 17.

Table 18.1 Some Common Acids and Bases and their Household Uses.

Theories of Acids and Bases

The Savante Arrhenius Theory (1884) Acid-ionizes in aqueous solution to produce H+

ions. Base-ionizes in aqueous solution to produce

OH- ions.

Svante August Arrhenius

Svante Arrhenius was born on February 19, 1859 and he died October 2, 1927. He was a Swedish physical chemist best known for the development of his acid-base theory. In 1903 he was awarded the Nobel Prize for Chemistry.

Strong and Weak Acids

A strong acid dissociates completely into ions in water:HA(g or l) + H2O(l) → H3O+(aq) + A-(aq)

A dilute solution of a strong acid contains no HA molecules.

A weak acid dissociates slightly to form ions in water:

HA(aq) + H2O(l) H3O+(aq) + A-(aq)

In a dilute solution of a weak acid, most HA molecules are undissociated.

[H3O+][A-]

[HA][H2O]Kc = has a very small value.

Classifying the Relative Strengths of Acids

Strong acids include the hydrohalic acids (HCl, HBr, and HI) and

oxoacids in which the number of O atoms exceeds the number of ionizable protons by two or more (eg., HNO3, H2SO4, HClO4.)

Weak acids include the hydrohalic acid HF,

acids in which H is not bonded to O or to a halogen (eg., HCN),

oxoacids in which the number of O atoms equals or exceeds the number of ionizable protons by one (eg., HClO, HNO2), and

carboxylic acids, which have the general formula RCOOH (eg., CH3COOH and C6H5COOH.)

Arrhenius Acids

HBr + H2O ------> H3O+ + Br -1 HBr is a strong acid--it completely dissociates.

HF + H2O = H3O+ + F HF is a weak acid--it only partially ionizes.

The Hydronium Ion

H+ ions are bare protons. In aqueous solutions these protons are

hydrated H(H2O)n+ where n is a small

number. The hydrated hydrogen ion is normally

represented as H3O+. In many reactions when it is obvious

that aqueous solutions are involved, H+ will be used to represent H3O+.

Strong Acids

HCl HBr HI HNO3

HClO3

HClO4

H2SO4

Classifying the Relative Strengths of Bases

Strong bases include water-soluble compounds containing O2- or OH- ions.

The cations are usually those of the most active metals:

M2O or MOH, where M = Group 1A(1) metal (Li, Na, K, Rb, Cs)

MO or M(OH)2 where M = group 2A(2) metal (Ca, Sr, Ba).

Weak bases include ammonia (NH3),

amines, which have the general formula

The common structural feature is an N atom with a lone electron pair.

RNH2, R2NH, or R3N

Arrhenius Bases

NaOH ------> Na+ + OH-1 NaOH is a strong base--it completely

dissociates.

NH3 + H2O = NH4+ + OH-1

NH3 is a weak base-- it only partially ionizes to produce OH- ions.

Strong Bases

LiOH NaOH KOH RbOH CsOH Ca(OH)2

Ba(OH)2

Sr(OH)2

Neutralization

Neutralization is the combination of hydrogen ions(acid) with hydroxide ions(base) to produce water(neutral).

The net ionic equation for the reaction of a strong acid with a strong base would be

H+ + OH- = H2O

The Bronsted-Lowry Theory (1923)

Acid- a proton (H+) donor Base- a proton(H+) acceptor An acid-base reaction is the transfer of a

proton from an acid to a base. NH3 + H2O = NH4

+ + OH-1 base acid conjugate

conjugate acid base

Johannes Nicolaus Bronsted

Johannes Brønsted was born on February 22, 1879 and died on December 17, 1947. He was a Danish physical chemist known for a widely applicable acid-base concept introduced in 1923. His work was independent of Lowry.

Thomas Martin Lowry

Thomas Lowry was born on October 26, 1874 and died on September 2, 1936.

He was an English chemist widely - known for an acid-base concept identical to that of Johannes N. Bronsted.

Brønsted-Lowry Acid-Base Definition

An acid is a proton donor, any species that donates an H+ ion.• An acid must contain H in its formula.

A base is a proton acceptor, any species that accepts an H+ ion.• A base must contain a lone pair of electrons to bond to H+.

An acid-base reaction is a proton-transfer process.

Conjugate Acid-Base Pairs

H2S + NH3 HS- + NH4+

NH3 accepts a H+ to form NH4+.

In the forward reaction:

In the reverse reaction:

H2S + NH3 HS- + NH4+

HS- accepts a H+ to form H2S.

H2S donates a H+ to form HS-.

NH4+ donates a H+ to form NH3.

Conjugate Acid-Base Pairs

H2S + NH3 HS- + NH4+

H2S and HS- are a conjugate acid-base pair:HS- is the conjugate base of the acid H2S.

A Brønsted-Lowry acid-base reaction occurs when an acid and a base react to form their conjugate base and conjugate acid, respectively.

NH3 and NH4+ are a conjugate acid-base pair:

NH4+ is the conjugate acid of the base NH3.

acid1 + base2 base1 + acid2

Base Acid+Acid Base+

Conjugate Pair

Conjugate Pair

Table 18.4 The Conjugate Pairs in some Acid-Base Reactions

Reaction 4 H2PO4- OH-+

Reaction 5 H2SO4 N2H5++

Reaction 6 HPO42- SO3

2-+

Reaction 1 HF H2O+ F- H3O++

Reaction 3 NH4+ CO3

2-+

Reaction 2 HCOOH CN-+ HCOO- HCN+

NH3 HCO3-+

HPO42- H2O+

HSO4- N2H6

2++

PO43- HSO3

-+

Sample Problem 18.4 Identifying Conjugate Acid-Base Pairs

PLAN: To find the conjugate pairs, we find the species that donated an H+ (acid) and the species that accepted it (base). The acid donates an H+ to becomes its conjugate base, and the base accepts an H+ to becomes it conjugate acid.

PROBLEM: The following reactions are important environmental processes. Identify the conjugate acid-base pairs.

(b) H2O(l) + SO32-(aq) OH-(aq) + HSO3

-(aq)

(a) H2PO4-(aq) + CO3

2-(aq) HPO42-(aq) + HCO3

-(aq)

SOLUTION:

(a) H2PO4-(aq) + CO3

2-(aq) HPO42-(aq) + HCO3

-(aq)acid1 base1 base2 acid2

The conjugate acid-base pairs are H2PO4-/HPO4

2- and CO32-/HCO3

-.

Sample Problem 18.4

(b) H2O(l) + SO32-(aq) OH-(aq) + HSO3

-(aq)

acid1 base1 acid2 base2

The conjugate acid-base pairs are H2O/OH- and SO32-/HSO3

-.

Figure 18.8 Strengths of conjugate acid-base pairs.

The stronger the acid is, the weaker its conjugate base. When an acid reacts with a base that is farther down the list, the reaction proceeds to the right (Kc > 1).

The Auto Ionization of Water

Water can act as both an acid and a base. Water is said to be amphoteric (amphiprotic).

H2O + H2O = H3O+ + OH -1 base acid conjugate

conjugate acid base

Autoionization of Water

Water dissociates very slightly into ions in an equilibrium process known as autoionization or self-ionization.

2H2O (l) H3O+ (aq) + OH- (aq)

The Gilbert Lewis Theory(1923)

Acid--species that accepts a share in a pair of electrons--an electron pair acceptor

Base--species that donates a share in a pair of electrons--an electron pair donor

:NH3 + H2O = H:NH3+ + OH-1

base acid conj.acid conj.base

Gilbert Newton Lewis

G. N. Lewis was born on October 23, 1875 and died on March 23, 1946. He was a famous American physical chemist, who in 1923, developed the electron-pair theory of acid-base chemical reactions. He is also known for the creation of Lewis structures for drawing chemical molecules.

The Lewis Acid-Base Definition

A Lewis base is any species that donates an electron pair to form a bond.

A Lewis acid is any species that accepts an electron pair to form a bond.

The Lewis definition views an acid-base reaction as the donation and acceptance of an electron pair to form a covalent bond.

B + H+ B H+

base acid adduct

Lewis Acids and Bases

The Lewis definition expands the classes of acids.

A Lewis base must have a lone pair of electrons to donate.Any substance that is a Brønsted-Lowry base is also a Lewis base.

A Lewis acid must have a vacant orbital (or be able to rearrange its bonds to form one) to accept a lone pair and form a new bond.Many substances that are not Brønsted-Lowry acids are Lewis acids.

Electron-Deficient Molecules as Lewis Acids

B and Al often form electron-deficient molecules, and these atoms have an unoccupied p orbital that can accept a pair of electrons:

BF3 accepts an electron pair from ammonia to form a covalent bond.

Lewis Acids with Polar Multiple Bonds

Molecules that contain a polar multiple bond often function as Lewis acids:

The O atom of an H2O molecule donates a lone pair to the S of SO2, forming a new S‒O σ bond and breaking one of the S‒O p bonds.

Metal Cations as Lewis Acids

A metal cation acts as a Lewis acid when it dissolves in water to form a hydrated ion:

The O atom of an H2O molecule donates a lone pair to an available orbital on the metal cation.

Sample Problem 18.15 Identifying Lewis Acids and Bases

SOLUTION:

PLAN: We examine the formulas to see which species accepts the electron pair (Lewis acid) and which donates it (Lewis base) in forming the adduct.

PROBLEM: Identify the Lewis acids and Lewis bases in the following reactions:

(a) H+ + OH- H2O (b) Cl- + BCl3 BCl4- (c) K+ + 6H2O K(H2O6)+

(a) The H+ ion accepts the electron pair from OH-. H+ is the Lewis acid and OH- is the Lewis base.

(b) BCl3 accepts an electron pair from Cl-. Cl- is the Lewis base and BCl3 is the Lewis acid.

(c) An O atom from each H2O molecule donates an electron pair to K+. H2O is therefore the Lewis base, and K+ is the Lewis acid.

Strong Electrolytes

ionize or dissociate completely Strong acids

HCl 100% H+ + Cl-1

Strong bases NaOH 100% Na+ + OH-1

Soluble salts NaCl 100% Na+ + Cl-1

Calculating Ion Concentrations

Determine the ion concentrations in 0.050M nitric acid solution.

Determine the concentration of all ions in a 0.020M solution of calcium hydroxide, Ca(OH)2.

The Auto-ionization of Water

H2O = H+ + OH-1 KW = [H+][OH-1] At 25OC KW = 1.0x10-14

Calculate the [H+] and [OH-1 ] in 0.050M HCl solution.

The pH Scale

pH = -log[H3O+]

The pH of a solution indicates its relative acidity:

In an acidic solution, pH < 7.00In a neutral solution, pH = 7.00In basic solution, pH > 7.00

The higher the pH, the lower the [H3O+] and the less acidic the solution.

The pH Scale a convenient way of expressing the

acidity and basicity of dilute aqueous solutions.

pH = -log[H+] This applies to other ion

concentrations as well pOH = -log[OH -1 ] Another useful relationship pH + pOH = 14

Figure 18.7 Methods for measuring the pH of an aqueous solution.

pH (indicator) paper

pH meter

Figure 18.5

The pH values of some familiar

aqueous solutions.

pH = -log [H3O+]

Table 18.3 The Relationship between Ka and pKa

Acid Name (Formula) Ka at 25°C pKa

Hydrogen sulfate ion (HSO4-) 1.0x10-2 1.99

Nitrous acid (HNO2) 7.1x10-4 3.15

Acetic acid (CH3COOH) 1.8x10-5 4.75

Hypobromous acid (HBrO) 2.3x10-9 8.64

Phenol (C6H5OH) 1.0x10-10 10.00

pKa = -logKa

A low pKa corresponds to a high Ka.

pH, pOH, and pKw

pH = -log[H3O+]

pOH = -log[OH-]

pH + pOH = pKw for any aqueous solution at any temperature.

pKw = pH + pOH = 14.00 at 25°C

Kw = [H3O+][OH-] = 1.0x10-14 at 25°C

Since Kw is a constant, the values of pH, pOH, [H3O+], and [OH-] are interrelated:• If [H3O+] increases, [OH-] decreases (and vice versa).• If pH increases, pOH decreases (and vice versa).

Figure 18.5 The relations among [H3O+], pH, [OH-], and pOH.

Sample Problem 18.3 Calculating [H3O+], pH, [OH-], and pOH

PROBLEM: In an art restoration project, a conservator prepares copper-plate etching solutions by diluting concentrated HNO3 to 2.0 M, 0.30 M, and 0.0063 M HNO3. Calculate [H3O+], pH, [OH-], and pOH of the three solutions at 25°C.

SOLUTION:

PLAN: HNO3 is a strong acid so it dissociates completely, and [H3O+] = [HNO3]init. We use the given concentrations and the value of Kw at 25°C to find [H3O+] and [OH-]. We can then calculate pH and pOH.

[H3O+] = 2.0 M

Calculating the values for 2.0 M HNO3:

pH = -log[H3O+] = -log(2.0) = -0.30

Kw

[H3O+][OH-] = =

1.0x10-14

2.0= 5.0x10-15 M

pOH = -log[OH-] = -log(5.0x10-15) = 14.30

Sample Problem 18.3

[H3O+] = 0.30 M

Calculating the values for 0.30 M HNO3:

pH = -log[H3O+] = -log(0.30) = 0.52

Kw

[H3O+][OH-] = =

1.0x10-14

0.30= 3.3x10-14 M

pOH = -log[OH-] = -log(3.3x10-14) = 13.48

[H3O+] = 0.0063 M

Calculating the values for 0.0063 M HNO3:

pH = -log[H3O+] = -log(0.30) = 2.20

Kw

[H3O+][OH-] = =

1.0x10-14

0.0063= 1.6x10-12 M

pOH = -log[OH-] = -log(1.6x10-12) = 11.80

pH Calculations

Calculate the pH of a solution in which [H+] = 0.030M

The pH of a solution is 4.597. Determine the [H+] of this solution.

Determine the [H+], [OH-1 ], pH and pOH for a 0.020M HNO3 solution.

Solving Problems Involving Weak-Acid Equilibria

The notation system

• Molar concentrations are indicated by [ ].

• A bracketed formula with no subscript indicates an equilibrium concentration.

The assumptions

• [H3O+] from the autoionization of H2O is negligible.

• A weak acid has a small Ka and its dissociation is negligible. [HA] ≈ [HA]init.

Ionization Constants for Weak Monoprotic Acids and Bases

Consider the reaction when the weak acid acetic acid is added to water.

CH3COOH + H2O = H3O+ + CH3COO-1

Ka = [H+][CH3COO-1 ] [CH3COOH] Write the equation for the ionization of

HCN in aqueous solution.

Calculation of Ionization Constants

In 0.12M solution, a weak acid HY is 5.0% ionized. Determine the value for the ionization constant for this weak acid.

The pH of a 0.10M solution of a weak monoprotic acid HA is 2.97. Calculate the value for the ionization constant of this weak acid.

Uses of the Ionization Constants

Determine the concentrations of all species in 0.15M acetic acid , CH3COOH, solution.

Ka = 1.8x10-5

Determine the concentrations of all species in 0.15M HCN solution. Ka = 4.0x10-10

Determine the concentrations of all species in 0.15M NH3. Kb=1.8x10-5

Uses of the Ionization Constants

The pH of an aqueous NH3 solution is 11.37. Determine the molarity of this aqueous ammonia solution. Kb=1.8x10-5