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

Prentice Hall ©2004

Chapter 14Chapter 14

Aqueous Equilibria: Acids and Bases


Acid–Base Concepts 01Acid–Base Concepts 01

Arrhenius Acid: A substance which dissociates in water to form hydrogen ions (H+) in solution.

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



• 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.


• A Lewis Acid is an electron-pair acceptor. These are generally cations and neutral molecules with vacant valence orbitals, such as Al3+, Cu2+, H+, BF3.

• A Lewis Base is an electron-pair donor. These are generally anions and neutral molecules with available pairs of electrons, such as H2O, NH3, O2–.

• The bond formed is called a coordinate bond.




• Problems 14.1,14.2• Write balanced equations for the dissociation of each of the following Brønsted–Lowry acids.(a) H2SO4 (b) HSO4

– (c) H3O+

• Problems 14.27• Identify the Lewis acid and Lewis base in each of the following reactions:

(a) AlCl3(s) + Cl–(aq) æ AlCl4–(aq)

(b) SO2 (aq) + OH–(aq) æ HSO3–(aq)

(c) Ag+(aq) + 2 NH3(aq) æ Ag(NH3)2+(aq)


Dissociation of Water 01Dissociation of Water 01

• Water can act as an acid or as a base.

H2O(l) æ H+(aq) + OH–(aq)

• This is called the autoionization of water.

H2O(l) + H2O(l) æ H3O+(aq) + OH–(aq)



• This equilibrium gives us the ion product constant for water.

Kw = Kc = [H3O+][OH–] = 1.0 x 10–14

If we know either [H3O+] or [OH–] then we can

determine the other quantity.



• Problem 14.6 The concentration of OH– in a sample of seawater is 5.0 × 10-6 M. Calculate the concentration of H3O+ ions, and classify the solution as acidic, neutral, or basic.

• Problem 14.7 At 50°C the value of Kw is 5.5 × 10-14.

What are the concentrations of H3O+ and OH– in a

neutral solution at 50°C?


pH – A Measure of Acidity 01pH – A Measure of Acidity 01

• The pH of a solution is the negative logarithm of the

hydrogen ion concentration (in mol/L).

pH = –log [H3O+]

pOH = –log [OH– ]

pH + pOH = 14

Acidic solutions: [H+] > 1.0 x 10–7 M, pH < 7.00

Basic solutions: [H+] < 1.0 x 10–7 M, pH > 7.00

Neutral solutions: [H+] = 1.0 x 10–7 M, pH = 7.00



• Problem 14.8 Calculate the pH of each of the following

Problem 14.9 Calculate the concentrations of H3O+ and OH– in each of the following solutions:

(a) Human blood (pH 7.40)(b) A cola beverage (pH 2.8)

Problem 14.10 Calculate the pH of(a)0.050 M HClO4 (b) 6.0 M HCl (c) 4.0 M KOH (d) 0.010 M Ba(OH)2

Problem 14.11 Calculate the pH of a solution prepared by dissolving 0.25 g of BaO in enough water to make 0.500 L of

solution


Strength of Acids and Bases01Strength 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.



HClO4

HI

HBr

HCl

H2SO4

HNO3

H3O+

HSO4–

HSO4–

HF

HNO2

HCOOH

NH4+

HCN

H2O

NH3

ClO4–

I–

Br –

Cl –

HSO4 –

NO3 –

H2O

SO42–

SO42–

F –

NO2 –

HCOO –

NH3

CN –

OH –

NH2 –

ACID CONJ. BASE ACID CONJ. BASE

Incr

easi

ng A

cid

Str

engt

h

Incr

easi

ng A

cid

Str

engt

h



• Stronger acid + stronger base

weaker acid + weaker base

• Problems 14.4 Predict the direction of the following:

HF(aq) + NO3– (aq) æ F–(aq) + HNO3(aq)

NH4+(aq) + CO3

–2 (aq) æ HCO3– (aq) + NH3(aq)


Acid Ionization Constants 01Acid Ionization Constants 01

• Acid Ionization Constant: the equilibrium constant for the ionization of an acid.

HA(aq) + H2O(l) æ H3O+(aq) + A–(aq)

[HA]

]][A[H3O

aK



7.1 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 –10

1.3 x 10 –10

HF

HNO2

C9H8O4 (aspirin)

HCO2H (formic)

C6H8O6 (ascorbic)

C6H5CO2H (benzoic)

CH3CO2H (acetic)

HCN

C6H5OH (phenol)

F–

NO2 –

C9H7O4 –

HCO2 –

C6H7O6 –

C6H5CO2 –

CH3CO2 –

CN –

C6H5O –

ACID Ka CONJ. BASE Kb

1.4 x 10 –11

2.2 x 10 –11

3.3 x 10 –11

5.9 x 10 –11

1.3 x 10 –10

1.5 x 10 –10

5.6 x 10 –10

2.0 x 10 –5

7.7 x 10 –5


CalculatingEquilibriumConcentrationinSolutions ofWeak Acids


HA æ H+ + A

(M): 0.50 0.00 0.00 (M): –x +x +x

Equilib (M): 0.50 –x x x


• Initial Change Equilibrium Table: Determine the pH

of 0.50 M HA solution at 25°C. Ka = 7.1 x 10–4.

InitialChange

(aq) (aq)-(aq)



• pH of a Weak Acid (Cont’d):

1. Substitute new values into equilibrium expression.

2. If Ka 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 Ka is not significantly smaller than [HA] the quadratic

equation must be used to solve for x and pH.



• 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.

aacbb

x2

42

02 cbxax



• 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.

100%[HA]

][H3OonDissociati %


Base Ionization Constants 01Base Ionization Constants 01

• 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) + H2O(l) æ BH+(aq) + OH–(aq)

• Calculations follow the same procedure as used for

a weak acid but [OH–] is calculated, not [H+].


Base Ionization Constants 02Base Ionization Constants 02

5.6 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 –10

1.5 x 10 –14

C2H5NH2 (ethylamine)

CH3NH2 (methylamine)

C8H10N4O2 (caffeine)

NH3 (ammonia)

C5H5N (pyridine)

C6H5NH2 (aniline)

NH2CONH2 (urea)

C2H5NH3+

CH3NH3+

C8H11N4O2+

NH4+

C5H6N+

C6H5NH3+

NH2CONH3+

BASE Kb CONJ. ACID Ka

1.8 x 10 –11

2.3 x 10 –11

2.4 x 10 –11

5.6 x 10 –10

5.9 x 10 –6

2.6 x 10 –5

0.67

Note that the positive charge sits on the nitrogen.


Diprotic & Polyprotic Acids 01Diprotic & Polyprotic Acids 01

• 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.


Diprotic & Polyprotic Acids 02Diprotic & Polyprotic Acids 02

Very Large1.3 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 –11

9.5 x 10 –8

1 x 10 –19

7.5 x 10 –3

6.2 x 10 –8

4.8 x 10 –13

H2SO4

HSO4–

C2H2O4

C2HO4–

H2SO3

HSO3–

H2CO3

HCO3–

H2SHS–

H3PO4

H2PO4–

HPO42–

ACID Ka CONJ. BASE Kb

HSO4 –

SO4 2–

C2HO4–

C2O42–

HSO3 –

SO3 2–

HCO3–

CO3 2–

HS–

S 2–

H2PO4–

HPO42–

PO43–

Very Small7.7 x 10 –13

1.5 x 10 –13

1.6 x 10 –10

7.7 x 10 –13

1.6 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 –12

1.6 x 10 –7

2.1 x 10 –2


Molecular Structure and Acid Strength 01Molecular Structure and Acid Strength 01

• 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



• 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.



• 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.



• 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.



• 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.



• 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.



• Oxoacids of Chlorine:



• Predict the relative strengths of the following groups of oxoacids:

a) HClO, HBrO, and HIO.

b) HNO3 and HNO2.

c) H3PO3 and H3PO4.


Acid–Base Properties of Salts 01Acid–Base Properties of Salts 01

• 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.

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

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