Chapter 16
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Chapter 16. Acid-Base Equilibria. constant. Dissociation of water. Autoionization or autoprotolysis. Ion-product constant Autoprotolysis constant. K w = [H + ][OH - ] = 1.0x10 -14 When [H + ] = [OH - ] neutral. Doesn’t usually happen.

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Chapter 16

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Chapter 16

Chapter 16

Acid-Base Equilibria


Chapter 16

constant

Dissociation of water

Autoionization or autoprotolysis

Ion-product constant

Autoprotolysis constant


Chapter 16

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


Chapter 16

H+ is a proton with no electrons.

In water:

+

é

ù

H

O

H

ê

ú

H

ë

û

Hydronium ion


Chapter 16

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.


Chapter 16

conj base

conj acid

conj acid

conj base

Conjugate Acid-Base Pairs


Chapter 16

The stronger an acid, the weaker its conjugate base.

The weaker an acid, the stronger its conjugate base.


Chapter 16

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


Chapter 16

You can also speak in terms of [OH-]

pOH = -log [OH-]

= 14 - pH

Because

pH + pOH = -log Kw = 14


Chapter 16

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


Chapter 16

Strong Acids and Bases

Strong electrolytes

Completely ionize

HA + H2O  A- + H3O+

Bases form hydroxides in solvent


Chapter 16

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.


Chapter 16

Weak acids

Only partially ionize

Acid dissociation constant


Chapter 16

O

=

C - O - H

O

N

Larger Ka means stronger acid.

ex.

0.020M solution

pH = 3.26

? Ka

pH = -log [H+] = 3.26

[H+] = 5.50x10-4


Chapter 16

O

O

=

=

C - OH

C - O

O

O

N

N

+ H+

A-

H+

HA

1:1


Chapter 16

O

O

=

=

C - OH

C - O

O

O

N

N

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

Let’s use niacin again.

+ H+

A-

H+

HA


Chapter 16

** Simplifying Assumption **

x is very very small compared to 0.010M

sooooooooo,

ignore x in denominator


Chapter 16

pH = -log [H+]

x = [H+] = 3.9x10-4

pH = 3.41

What percent of niacin molecules ionized?


Chapter 16

Polyprotic Acids

ex. H2SO4 H3PO4H2SeO4

H2SO4 H+ + HSO4-

Ka1 = 1.7x10-2

HSO4- H+ + SO42-

Ka2 = 6.4x10-8

Ka1 always larger than Ka2

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


Chapter 16

Weak Bases

ex. Amines

“an organic substituted ammonia”

methyl amine

ammonia

NH3

CH3

H

N

H

N

H

H

H


Chapter 16

H

CH3 + H2O 

H

N

H

N

CH3 + OH-

H

H

Anions of weak acids

ClO- + H2O  HClO + OH-

Kb = 3.3x10-7

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


Chapter 16

Ka and Kb

How are they related?


Chapter 16

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.


Chapter 16

K1 x K2 = K3

rxn 1 rxn 2 rxn 3


Chapter 16

Special Case

Ka x Kb = Kw

For conjugate acid-base pairs.


Chapter 16

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


Chapter 16

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.


Chapter 16

Oxyacids

O

Have unprotonated and

protonated oxygens.

H

O

S

O

H

O

H3PO4

Y

O

H

  • 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


Chapter 16

Carboxylic Acids

O

COOH = Carboxyl group

C

R

OH

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.

H

F

O

O

H

C

C

F

C

C

O

O

H

H

F

H

Acetic acid

Ka = 1.8x10-5

Trifluoroacetic acid

Ka = 5.0x10-1


Chapter 16

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’.


Chapter 16

H+Bronstead-Lowry acid

also a Lewis acid

H+electron pair acceptor

OH-

Electron pair donor

Lewis base

also Bronstead-Lowry base


Chapter 16

H

H

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

B

H

Incomplete Octet

H

Lewis Base

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

N

H

H


Chapter 16

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

H

H

O

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


Chapter 16

O

H

H

O

H

C

= =

H

O

O

H

O

O

H

C

= =

O

C

= =

O

O

Carbonic acid

Hydrolysis of metal ions

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


Chapter 16

3+

H

H

H

O

O

O

Fe3+

H

H

H

Fe

H

H

O

O

O

O

H

H

H

H

H

H

6 of

these

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.


Chapter 16

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


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