chapter 16 acid-base equilibria. dissociation of water autoionization or autoprotolysis ion-product...

Chapter 16

Acid-Base Equilibria

Dissociation of water

Autoionization or autoprotolysis

Ion-product constantAutoprotolysis constant

-(aq)(aq)(l)2 OHHOH

]][OHH[K

O]H[

]][OH[HK

-w

2

-

constant

Kw = [H+][OH-] = 1.0x10-14

When [H+] = [OH-] neutral. Doesn’t usually happen.

As one increases, the other decreases; the product must equal 1.0x10-14.

When

[H+] > [OH-] acidic

[OH-] > [H+] basic

H+ is a proton with no electrons.In water:

H

HOH

Hydronium ion

Bronstead-Lowry Acid-Base

Acid - Can donate a proton

Base - Can accept a proton

*Doesn’t have to be in H2O. Can be in other solvents.

Conjugate Acid-Base Pairs

(aq)3-(aq)(l)2(aq) OHAOHHA

conj base

conj base

conj acid

conj acid

(aq)(aq)(l)23(aq) OHNH4OHNH

(aq)3-2(aq)(l)22(aq) OHNOOHHNO

The stronger an acid, the weaker its conjugate base.

The weaker an acid, the stronger its conjugate base.

pH scale

pH = -log [H+]

Remember

Kw = (1x10-7)(1x10-7) = 1.0x10-14

pH = -log [H+] = -log (1x10-7)

pH = 7 (neutral)

[H+] pH

acidic > 1.0x10-7 < 7.00

basic < 1.0x10-7 > 7.00

You can also speak in terms of [OH-]

pOH = -log [OH-]

= 14 - pH

Because

pH + pOH = -log Kw = 14

Measure pH by

pH meter

Acid-base indicators

Litmus

red = pH < 5

blue = pH > 8

Figure 16.7 shows several acid-base indicators and their ranges

Strong Acids and Bases

Strong electrolytes

Completely ionize

HA + H2O A- + H3O+

Bases form hydroxides in solvent

In H2O, Alkali metal hydroxides

Alkaline earth metal

Hydroxides (except Be)

Many are insoluble

Also, substances that will abstract a H+ from H2O.

O2- + H2O 2OH-

Na2O or CaO would do this. O2-, H-, N3- bases that would do this.

Weak acids

-(aq)(aq)(aq) AH HA

Only partially ionize

[HA]

]][A[HK

-

a

Acid dissociation constant

Larger Ka means stronger acid. ex.

N

O

C - O - H

O=

0.020M solutionpH = 3.26? Ka

pH = -log [H+] = 3.26[H+] = 5.50x10-4

N

O

C - OH

O=

N

O

C - O

O=

+ H+

HA A- H+

1:1

a

24-5-

4-

need stillthis

4-4-

a

-

a

K(0.0195)

)(5.5x101.55x10

[HA]0.0195MM)(5.5x10-M020.0

? [HA]

]][5.5x10[5.5x10K

[HA]

]][A[HK

Can calculate pH in same manner if you have Ka and concentration of solution.

Let’s use niacin again.

N

O

C - OH

O=

N

O

C - O

O=+ H+

HA A- H+

x)-(0.010

[HA]][H

(x)

][A

(x)

1.5x10

pH ?solution 0.010M1.5x10K

-5-

-5a

x)-(0.010

x1.5x10

25

** Simplifying Assumption **

x is very very small compared to 0.010M

sooooooooo,

ignore x in denominator

4-7-2

7-5-2

5-2

3.9x10x1.5x10x

1.5x10)(0.010)(1.5x10x

x105.10.010

x

pH = -log [H+]

x = [H+] = 3.9x10-4

pH = 3.41

What percent of niacin molecules ionized?

3.9%100x 0.010

x109.3 -4

Polyprotic Acids

ex. H2SO4 H3PO4 H2SeO4

H2SO4 H+ + HSO4-

Ka1 = 1.7x10-2

HSO4- H+ + SO4

2-

Ka2 = 6.4x10-8

Ka1 always larger than Ka2

If Ka1/ Ka2 103, can estimate pH by Ka1 only.

Weak Bases

ex. Amines

“an organic substituted ammonia”

ammoniaNH3

N

H

HH NH

CH3

H

methyl amine

N

H

CH3 + H2O H N

H

CH3 + OH-H

H

ClO- + H2O HClO + OH-

Kb = 3.3x10-7

][ClO

][HClO][OHx103.3

-

-7-

Can use this in the same manner in which you used Ka.

Anions of weak acids

Ka and Kb

How are they related?

][NH

]][OH[NHK

][NH

]][H[NHK

OHNHOHNH

HNHNH

3

-4

b4

3a

-(aq)4(aq)(l)23(aq)

(aq)3(aq)4(aq)

-(aq)(aq)(l)2

-(aq)4(aq)(l)23(aq)

(aq)3(aq)4(aq)

OHHOH

OHNHOHNH

HNHNH

1)

2)

3)

When two reactions are added together, the equilibrium constant for the third reaction is given by the product of equilibrium constants of equations 1 and 2.

K1 x K2 = K3

rxn 1 rxn 2 rxn 3

w-

3

-4

4

3ba

K]][OH[H

][NH

]][OH[NH

][NH

]][H[NHKK

Special Case

Ka x Kb = Kw

For conjugate acid-base pairs.

Bond polarity and Bond strength effect on Acid-base behavior: In binary acids

polarity(across a row) acidity

bond strength(in a group) acidity

stability of conj. base acidity

Metal hydrides are basic or show no acid/base properties in H2O.

Nonmetal hydrides are acidic or show no acid/base properties in H2O (except NH3)

Acidity increases moving down a group.

Oxyacids

H O S

O

O

O

HHave unprotonated and protonated oxygens.

Y O H H3PO4

• As electronegativity of Y increases, acidity increases.

• As number of unprotonated oxygens increases, acidity increases (effect of formal charge and oxidation number)

•Ex. HClO, HClO2, HClO3, HClO4

Carboxylic Acids

RC

OH

O COOH = Carboxyl group

R = H or an organic group.

The more electron withdrawing R is, the greater the acidity (this stabilizes anion and weakens O-H bond)

ex.

CH C

H

H O

OH

Acetic acidKa = 1.8x10-5

CF C

F

F O

OH

Trifluoroacetic acidKa = 5.0x10-1

Lewis Acids and Bases

This is a completely different definition for acid/base chemistry than what you have seen thus far!!!

Lewis acid = electron pair acceptor

Lewis base = electron pair ‘donor’

Not giving them away, just has them available to ‘share’.

H+ Bronstead-Lowry acid

also a Lewis acid

H+ electron pair acceptor

OH -

Electron pair donorLewis basealso Bronstead-Lowry base

B

H H

H

BH3 not a Bronstead-Lowry acid, but it’s a Lewis acid

Incomplete Octet

N

H

H

H

Lewis Basehas an electron pair available to attack an area that is e- deficient

Transition metal ions are often Lewis Acids. They have vacant d orbitals. (s and p also)

3

6

_3

NCFe

NC6Fe

H

HO

O = C = O Can be a Lewis Acid because e- density around the C is bound in just 2 directions.

H

HO

=

=

O

O

C

H

HO

=

=

O

O

C

H

H O

=

=

O

O

CCarbonic acid

Hydrolysis of metal ions

Metal ions have positive charge so they attract the lone e- pair on H2O molecules

6 ofthese

HO

H Fe3+H

OH

HO

H

HO

H

HO

H

H

O

H

H

O

H

Fe

3+

Because the metal is (+), e- density of H2O moves toward the metal. When this happens, there is less e- density in water’s O-H bonds, so H+ can come off easier… pH will drop.

The higher the charge density of the metal ion, the greater the acidity of its aqua complex.

OHin acidity stronger radius ionicsmaller

charge )(greater soooo

radius ionic

chargedensity charge

2