Acid Base Definitions

Originally recognized by properties like taste,
feel, reactions with indicators
–
–
Acids taste sour and turn blue litmus red
Bases are bitter, feel slippery, and turn red litmus
blue
Arrenhius Definition


Very limited
Very few substances could be classified by this
definition.
–
–
Acids produce excess Hydrogen ions when added
to water
Bases produce excess Hydroxide ions when added
to water
Bronsted-Lowry Definition

Most frequently used
–
–
Acids are proton donors (give off an H+)
Bases are proton acceptors
Strong Acids



Acid that dissociates completely in water
100% of the sample breaks apart into ions
Seven strong acids
–
–
–
–
–
–
–
a.
b.
c.
d.
e.
f.
g.
HClO4
HClO3
H2SO4
HNO3
HCl
HBr
HI
- Perchloric acid
- Chloric acid
- Sulfuric acid
- Nitric acid
- Hydrochloric acid
- Hydrobromic acid
- Hydroiodic acid
Strong Base



Base that dissociates completely in water
100% of the sample breaks apart into ions
Strong bases
–
Hydroxides of the metals in group 1A and 2A
Be or Mg)
(not
Acid-Base Equilibrium

A system at equilibrium can be described by
its equilibrium constant.
–
–
For acidic systems, we call the equilibrium
constant the acid dissociation constant and is
symbolized as Ka.
For basic systems, the equilibrium constant is
known as the base dissociation constant and is
symbolized as Kb.
Acid-Base Equilibrium

Let’s consider water in equilibrium with
hydroxide and hydronium ions
–
The balanced equation is
2H2O (l)  H3O1+ (aq) + OH1- (aq)
–
The equilibrium constant would be written as
Keq = [H3O1+] [OH1-]
Acid-Base Equilibrium

The equilibrium constant for water has a set
value, 1.0 x 10-14, and is referred to Kw
–
–
–
In pure water, the concentration of H3O1+ and OH1are equal, and the product of the two
concentrations is always equal to 10-14.
Thus in pure water, [H3O1+] = [OH1-] = 1.0 x 10-7
If the concentration of one of these two ions
changes, the concentration of the other ion must
also change. For example, if the [H3O1+] = 10-3,
then the [OH1-] = 10-11.
Acid-Base Equilibrium

It is important to remember that when you are working
with a system involving water, the water always
contributes to the [H3O1+] and [OH1-]
–
–
The contribution of the water can be ignored when working with
strong acids or if the [H3O1+] from the water is < or = to 1% of
the contribution of the weak acid.
For strong acids, the [H3O1+] will be equal to the concentration
of the acid.


For example, a 0.1 M solution of HCl would have [H3O1+] = 0.1 M
The same is true for strong bases and the [OH1-]
Summary

Water
–
–
–
–
–
2H2O  H3O1+ + OH1Kw = [H3O1+] [OH1-] = 1.0 x 10-14
[H3O1+] = [OH1-] = 1.0 x 10-7 (pure water)
If [H3O1+] increases, [OH1-] will decrease
You must remember that water contributes to the
[H3O1+] for equilibrium

Can be ignored for strong acids
Weak acid or weak base




Do NOT dissociate completely in water
In water, establishes equilibrium between the
molecular form and ionic form
Any acid or base that is not a strong acid or
base is weak
Example:
Acetic acid
HC2H3O2 + H2O  C2H3O21- + H3O1+
Conjugate base


Ion that is formed when an acid donates a
Hydrogen ion (proton)
Examples:
Acid
Conjugate base
H2SO4
HSO41HNO3
NO31HC2H3O2 C2H3O21-
Conjugate base

If original acid is a strong acid, the conjugate
base is so weak that is does not behave as a
base.
–


It behaves as a neutral species
If original acid is a weak acid, the conjugate
base behaves as a weak base
(Important for equilibrium considerations later)
Conjugate acid


Ion that is formed when a base accepts a
Hydrogen ion
Examples:
Base
Conjugate acid
NaOH
H2O
NH3
NH41+
Conjugate acid

If original base is a strong base, the conjugate
acid is so weak that is does not behave as an
acid.
–

It behaves as a neutral species
If original base is a weak base, the conjugate
acid behaves as a weak acid
Direction of acid-base reactions

The stronger acid and base will react with each
other to yield the weaker acid and base
Example: HClO4 + H2O
stronger acid
stronger base
 ClO41- + H3O1+
weaker base
weaker acid
Acidic Hydrogen


Hydrogen that will be donated or removed
Not all hydrogen’s in an acid can be donated
–

In oxy acids, the hydrogen attached to an
oxygen is the acidic hydrogen
–


This is especially true for organic acids
Example: Acetic acid
Most acids are monoprotic (donate one H)
Some acids are polyprotic (i.e. H2SO4)
Autoionization of water

The transfer of a hydrogen ion from one water
molecule to another water molecule,
–


results in the formation of a hydroxide ion and a
hydronium ion.
Equation:
2 H2O  H3O1+ + OH1Equal amounts of hydroxide and hydronium
are formed,
–
so water remains neutral.
Amphoteric Substances

A species that can behave as either an acid or
a base

Water is the best example of an amphoteric
substance
Amphoteric Substances

Many aluminum compounds are also
amphoteric:

Base:
Al(OH)3 + 3HCl → AlCl3 + 3H2O

Acid:
Al(OH)3 + NaOH → NaAl(OH)4
pH system

The pH of a system is an indication of the
[H3O1+]. While it is based on the
autoionization of water, it works for all acidbase systems.
pH system




Definitions:
pH = -log [H3O1+]
pOH = -log [OH1-]
pKw = -log Kw
Since Kw = 1.0 x 10-14, pKw = -log (1.0 x 10-14) = 14

Kw = [H3O1+] [OH-1] 
pKw = pH + pOH = 14
pH Strong Acids and Bases

[H3O1+] = initial concentration of acid
–
To find the pH of a strong acid, use the initial
concentration of the acid as the concentration of
H3O1+

[OH1-] = initial concentration of base

pH (strong acid) = -log (initial conc.)
14.5 pH of weak acids and bases




Weak acids/bases do NOT dissociate
completely
To find pH, you must first find the [H3O1+]
This is done by setting up an ICE chart!!!
Weak acid equilibrium
–
–
HA  H1+ + A1Ka = [H+] [A-] / [HA]
pH weak acids and bases

Weak base equilibrium
–
–
B + H20  BH1+ + OH1Kb = [BH+] [OH-] / [B]
Weak acid equilibrium

Percent ionization = Degree of dissociation
–
–

Amount of substance that breaks apart into ions
If an acid has an initial concentration of 0.1 M and a
1 % ionization, [H+] = 1% of 0.1 M
We can calculate the value of the equilibrium
constant, as well as equilibrium concentrations
for weak acid systems. Let’s try some
examples!
Polyprotic Acids

Polyprotic Acids – more than one acidic
hydrogen
–
–
Dissociate in a stepwise fashion
Ka values assigned for each dissociation step



The first step in the dissociation happens completely
before the next step begins.
When solving problems involving polyprotic acids, you
can sometimes ignore the contribution of the subsequent
dissociations because it is ≤ 1% of the dissociation of the
first step.
Let’s try a couple
Acid/Base Properties of Salts

Review 1st Sem. Lab Notes

Ions are often modified when dissolved in
solution. Examine the photos of Fe(III)
salts and solutions. Why do they have
different colors?
Acid/Base Properties of Salts



Fe(NO3)3.6H2O contains pink Fe(H2O)63+
Solutions may hydrolyze to give yellow
Fe(H2O)5OH2+ or even reddish brown
Fe(H2O)3(OH)3
FeCl3.6H2O contains ions such as yellow
Fe(H2O)5Cl2+
Hydrolysis

Hydrolysis is more
important for more
highly charged ions
Hydrolysis

Highly charged metal ions (> +3) cause pH shifts
due to hydrolysis:
Fe(H2O)63+ + H2O ⇌ Fe(H2O)5OH2+ + H3O+

Many other salts also cause pH shifts when
dissolved in water. These salts contain the
conjugate acid or base of a weak base or acid.
NH4Cl: NH4+ + H2O ⇌ NH3 + H3O+ pH<7
NaCH3CO2: CH3CO2- + H2O ⇌ CH3CO2H + OH pH>7

These reactions are called hydrolysis
Hydrolysis

Hydrolysis is not observed with ions derived
from strong acids or bases:
Cations of group I and II (except Be2+)
Anions: Cl-, Br-, I-, NO3-, ClO4-

Hydrolysis is observed for:

Cations with charge > +3
Transition metal +2 ions
Some post-transition metal ions with high charge
Common for Fe3+, Cr3+, Al3+, Zn2+, Cu2+, Bi3+, Pb4+
Hydrolysis

See table below for values of Ka for metal ions.
Na+
Li+
Be2+
Mg2+
Ba2+
Cr3+
Zr4+

95 pm
60 pm
31 pm
65 pm
135 pm
69 pm
78 pm
3.3 x 10-15
1.5 x 10-14
3.2 x 10-7
3.8 x 10-12
1.5 x 10-14
9.8 x 10-5
6.0 x 10-1
Greater values of Ka for ions with larger charge
and smaller size.
Hydrolysis of Salts





Cation/Anion
from:
Strong base,
strong acid
Strong base,
weak acid
Weak base,
strong acid
Weak base,
weak acid
NaCl
no hydrolysis
pH = 7
LiCN
anion hydrolysis
pH > 7
cation hydrolysis
pH < 7
NH4Cl
NH4CN
cation and anion hydrolysis
pH depends on relative Ka and Kb
Hydrolysis of Salts

Calculate the pH of a solution the same as for
any weak acid or weak base, using the
appropriate Ka or Kb for the equilibrium
constant.
Hydrolysis of Salts




Is a solution of NH4OCN acidic
or basic?
NH4+ + H2O ⇌ NH3 + H3O+
Ka = 1.0 x 10-14/1.76 x 10-5
= 5.7 x 10-10
OCN- + H2O ⇌ HOCN + OHKb = 1.0 x 10-14/3.46 x 10- 4 =
2.9 x 10-11
Produces more H3O+ than OH-,
so the solution is acidic.
H 3 O+
OH-
Effect of Structure on Acid-Base
Properties

1. Bond Polarity (with H)
–
–
–
As polarity increases, acid strength increases
The more polar a molecule is, the easier it will be to
remove the acidic hydrogen.
More of the sample will dissociate
Effect of Structure on Acid-Base
Properties

2. Bond Strength
–
–
–
As bond strength increases, acid strength
decreases
It is more difficult to remove the acidic hydrogen
when the bond strength increases.
Less of the sample will dissociate if it is more
difficult to remove the hydrogen, leading to a weaker
acid.
Effect of Structure on Acid-Base
Properties

Atomic radius
–
–
–
Larger atoms form weaker bonds.stronger acids
As mentioned, weaker bonds result in stronger
acids.
The bigger the atom, the weaker the bond

–
The weaker the bond, the easier to remove the H
Acids from the same family


Elements at the bottom of the column are stronger
HI > HBr > HCl > HF
Effect of Structure on Acid-Base
Properties

Electronegativity
–
–
–
As electronegativity increases, bond polarity also
increases
Acids formed with highly electronegative elements
will be stronger than acids formed with elements
that have low electronegativity values.
Acids in the same period


Elements to the right will form stronger acids
HF > H2O > NH3
Effect of Structure on Acid-Base
Properties

Number of acidic hydrogens
–
The Neutral acid is always the strongest
With each hydrogen that is removed, the remaining
acid is much weaker
–
H3PO4 > H2PO41- > HPO42- > PO43-
–
Effect of Structure on Acid-Base
Properties

Oxy acids
–
a) Electronegativity of nonmetal:

The greater the electronegativity, the stronger the acid

when comparing oxy acids of elements in the same family,
the element at the top of the column will form the strongest
oxy acid

HClO > HBrO > HIO
Effect of Structure on Acid-Base
Properties

Oxy acids
–
b) Number of oxygen atoms:

This is true because the addition of oxygen atoms
increases the difference in electronegativity between the
oxygen atoms and the central atom, which in turn causes
the molecule to be more polar


More oxygen atoms = stronger acid
More oxygen atoms causes the molecule to be more polar

HClO4 > HClO3 > HClO2 > HClO
Effect of Structure on Acid-Base
Properties

Oxides that react with water to produce bases
are called basic oxides or base anhydrides.

These are generally metal oxides
Metal (basic) oxides

Sodium oxide reacts with water to produce the
strong soluble base sodium hydroxide:
Na2O (s) + H2O (l)  NaOH (aq)

Calcium oxide ("lime") reacts with water to
produce the insoluble base calcium hydroxide
("slaked lime"):
CaO (s) + H2O (l)  Ca(OH)2 (s)
Acid-base properties of oxides

Oxides that react with water to produce acids
are called acidic oxides or acid anhydrides.

These are generally non-metal oxides.
Non-Metal (Acidic) oxides

Carbon dioxide and sulfur trioxide are two such
compounds. Carbon dioxide reacts with water
to produce carbonic acid:
CO2 (g) + H2O (l) ---> H2CO3 (aq)

Sulfur trioxide reacts with water to produce
sulfuric acid:
SO3 (g) + H2O (l) ---> H2SO4 (aq)
Non-Metal (Acidic) oxides

Most acidic oxides are oxides of nonmetals, or
of metals that are in very high oxidation states.
–
An example of an acidic metal oxide is CrO3, which
reacts with water to produce chromic acid, H2CrO4
Acid-base properties of oxides

Oxides with metalloids and oxygen are
amphoteric

The acidity of oxides increases as you move
up and to the right across the periodic table.
Lewis Acids-Base Model

Lewis Definition – includes the greatest
number of compounds
–
–


Acids are electron pair acceptors
Bases are electron pair donors
An acid and a base come together to form a
covalent bond using the electron pair donated
by the base
Best example: Boron complexes