1. Aqueous Acid-Base Equilibria
Acids and bases
Acid-base properties of water (Kw)
pH scale
Strength of Acids and Bases
Weak acid (Ka)
Weak base (Kb)
Relation between conjugate acid-base ionization constants (Ka and Kb)
Chemical structure and acidity
Acid-base properties of salt solutions

2. Arrhenius Acids and Bases
Arrhenius acid - a compound that increases [H+] in water
another example
HNO3(aq) + H2O(l)  H3O+(aq) + NO3-(aq)
Arrhenius Base : a compound that increases [OH-] in water
NH3(aq) + H2O (l) NH4+(aq) + OH-(aq)

3. Arrhenius A. Arrhenius acid - a compound that increases [H+] in water
Chemists often use the notation H+(aq) for the H3O+(aq) ion, and call it the hydrogen ion.
Remember, however, that the aqueous hydrogen ion is actually chemically bonded to water, that is, H3O+.
Even that is a shorthand, as each H3O+ ion coordinates with the water molecules around it!
Examples of hydrogen ion structures from Martin Chaplin under CC License

4. Limitations of the Arrhenius Concept
The Arrhenius concept is limited in that it looks at acids and bases in aqueous solutions only.
All Arrhenius acids must be protic
In addition, it singles out the OH- ion as the source of base character, when other species can play a similar role.
Broader definitions of acids and bases are discussed in the next sections.

5. Brønsted-Lowry Acids and Bases
Brønsted-Lowry acid : a proton donor
Brønsted-Lowry base : a proton acceptor
NH3(aq) + H2O (l) NH4+(aq) OH-(aq)
base
acid
acid
base
base
acid
conjugate acid
conjugate base
Conjugate acid-base pair : acid-base pair that are related by a loss or gain of a proton H+

6. Conjugate Acid Base Pairs
Conjugate acid-base pair : acid-base pair that are related by a loss or gain of a proton H+
All acid base reactions have two conjugate acid base pairs

7. Further Examples of Conjugate Acid Base Pairs
Conjugate acid-base pair : acid-base pair that are related by a loss or gain of a proton H+

8. Brønsted-Lowry vs Arrhenius
Bronsted-Lowry
Arrhenius
Aqueous
Doesn’t have to be
Yes
Base
Proton acceptor
OH-
Acid
Protic

9. Lewis Acid-Base Concept
A Lewis acid is a substance that can accept a pair of electrons (electron pair acceptor)
A Lewis base is a substance that can donate a pair of electrons (electron pair donor)
base • • •
H OH-  H – O – H
acid
N H • H
N H H +
base H+ +
acid 
F B F
N H • H
F B F
N H H
base 
acid

10. Brønsted-Lowry vs Arrhenius vs Lewis
Bronsted-Lowry
Arrhenius
Aqueous
Doesn’t have to be
Must be
Base
Electron Pair Donor
Proton acceptor
OH-
Acid
Electron Pair Acceptor
Protic

11. Acid-Base Properties of Water
Water is amphoteric – it can act as either an acid
or a base
Since water is amphoteric it can autoionize yielding both OH-(aq) and H3O+ (aq) ions

12. Autoionization of water
We can write the equilibrium expression for the autoionization of water As
For pure water (density g/ml at 25 oC)

13. Kw
The equilibrium concentrations of OH-(aq) and H3O+(aq) at 25 oC are x 10-7 M so
and numerically
For most purposes (e.g. tests unless told otherwise) we will approximate Kw as 1.0 x and in neutral water [H3O+(aq)] = [OH-(aq)] = 1.0 x 10-7 M
Depending on context we will also use either H3O+(aq) or H+(aq)

14. pH pOH pK
Because concentrations of ions can vary over many orders of magnitude we use a logarithmic scale. For example the concentration of H3O+(aq) is usually transformed to
where pX indicates minus the base 10 logarithm of X. OH-(aq) concentrations are transformed to
A little bit of manipulation shows that pKw= -log Kw

15. pH Scale Acids and bases 7 is neutral on the pH scale.
Anything above 7 is basic &
anything below 7 is acidic
If the pH of a solution decreases one says it has been acidified or become more acidic.
OTOH if the pH of a solution increases one says it has become more basic.

16. Acid Equilibrium Ionization Constant
Acids that (almost) completely ionize in water
are called strong acids and those that only partially ionize are called weak acids. The quantitative amount of ionization is described by the acid ionization equilibrium constant Ka
We can also define pKa = -log10Ka
The larger pKa the stronger the acid and visa versa

17. Acid Equilibrium Ionization Constant
Similarly for bases we can define Kb
and pKb= -log10 Kb
There is a simple relationship between Ka, Kb and Kw
Ka Kb=Kw
And
pKa + pKb =pKw = 14.00

18. Aqueous Acid-Base Equilibria
When an acid ionizes in water it produces positive hydrogen ions and negative conjugate base ions
OTOH, if we mixed a soluble salt of the negative (CN-) ion into water the following reaction would occur
If we add these two reaction up the net is
and
KaKb= = [H+][OH-]=Kw

19. Range of Ka For Acid/Base Conjugate PairsHA
K a
pKa A−
K b
pKb
hydroiodic acid HI
2 × 109
−9.3 I−
5.5 × 10−24
23.26
sulfuric acid (1)*
H2SO4
1 × 102
−2.0
HSO4−
1 × 10−16
16.0
nitric acid
HNO3
2.3 × 101
−1.37
NO3−
4.3 × 10−16
15.37
hydronium ion
H3O+
1.0
0.00
H2O
1.0 × 10−14
14.00
sulfuric acid (2)*
1.0 × 10−2
1.99
SO42−
9.8 × 10−13
12.01
hydrofluoric acid HF
6.3 × 10−4
3.20 F−
1.6 × 10−11
10.80
nitrous acid
HNO2
5.6 × 10−4
3.25
NO2−
1.8 × 10−11
10.75
formic acid
HCO2H
1.78 × 10−4
3.750
HCO2−
5.6 × 10−11
10.25
benzoic acid
C6H5CO2H
6.3 × 10−5
4.20
C6H5CO2−
1.6 × 10−10
9.80
acetic acid
CH3CO2H
1.7 × 10−5
4.76
CH3CO2−
5.8 × 10−10
9.24
pyridinium ion
C5H5NH+
5.9 × 10−6
5.23
C5H5N
1.7 × 10−9
8.77
hypochlorous acid
HOCl
4.0 × 10−8
7.40
OCl−
2.5 × 10−7
6.60
hydrocyanic acid
HCN
6.2 × 10−10
9.21
CN−
1.6 × 10−5
4.79
ammonium ion
NH4+
5.6 × 10−10
9.25
NH3
1.8 × 10−5
4.75
water
OH−
1.00
acetylene
C2H2
1 × 10−26
26.0
HC2−
1 × 1012
−12.0
ammonia
1 × 10−35
35.0
NH2−
1 × 1021
−21.0
*The number in parentheses indicates the ionization step referred to for a polyprotic acid.

20. HCl(aq) +H2O(l)  H3O+(aq) + Cl-(aq)
Direction of Acid-Base Reactions
We know that if we mix HCl in water the hydrochloric acid will almost completely ionize
HCl(aq) +H2O(l)  H3O+(aq) + Cl-(aq)
In the language of conjugate acid base pairs the HCl(aq) is a strong acid and the conjugate Cl-(aq) base is a weak base.
The table to the left shows the relative strengths of different acid base pairs