Acids and Bases
Arrhenius Acids and Bases Definition
An Arrhenius acid is a substance that when added to water increases the
concentration of H1+ ions present.
An Arrhenius base is a substance that when added to water increases the
concentration of OH1- ions present. HCl is an example of an Arrhenius acid
and NaOH is an example of an Arrhenius base.
Bronsted-Lowry acids and bases
A Bronsted-Lowry (BL) acid is defined as any substance that can donate a
hydrogen ion (proton) and a Bronsted-Lowry base is any substance that can
accept a hydrogen ion (proton).
BL acids and bases must come in what is called conjugate pairs. For
example, consider acetic acid dissolved in water:
What are the conjugate acid/base pairs?
Similarly when ammonia is dissolved in water, one has
What are the conjugate acid/base pairs
Another important advantage of the BL definition is that we are not limited
to water as the solvent. Consider the reaction that occurs when HCl is
dissolved in ammonia:
Here, HCl acts as a BL acid with Cl
as a BL base with NH
as its conjugate base. Also, NH
acts
as its conjugate acid.
Some species can act either as a BL acid or a BL base. Such beasts are
called amphoteric. An example is the hydrogen carbonate ion, HCO .
When dissolved in water, two posible reaction can occur:
or
H3O+ is a hydronium ion. A water molecules has accepted a proton.
In the first of these, HCO
base, while
acts as a BL acid with CO
as its conjugate
In the second it acts as a BL base with H CO as its conjugate acid.
Acid base reactions are reversible and therefore equilibrium expressions can
be written
A strong acid is one in which the equilibrium lies far to the right.
A strong acid produces a weak conjugate base.
A weak acid is one in which the equilibrium lies far to the left. Not much
acid has dissociated.
A weak acid produces a stronger conjugate base.
pH?
An acidic pH is
A basic pH is
A neutral pH is
pH = -log[H+]
pH of strong acids
Because a strong acid is completely ionized the concentration of the acid is
the concentration of the hydrogen ions.
For a 0.1 M H+ solution
pH = -log(0.1) = -log (10-1) = 1
For a 0.01 M H+ solution
pH = -log(0.01) = -log(10-2) = 2
What is the pH of 0.567 M nitric acid?
What is the hydrogen ion concentration [H+] for a solution that has a pH of
3.2?
1.0 * 103
Table of Acid and Base Strength
Acid
Name
Formula
Formula
Perchloric acid
HClO4
ClO4 Hydroiodic acid
HI
IHydrobromic
HBr
Bracid
Hydrochloric
HCl
Clacid
Sulfuric acid
H2SO4
HSO4 -
2.4 * 101
-------5.4 * 10-2
Nitric acid
Hydronium ion
Oxalic acid
1.3 * 10-2
Sulfurous acid
1.0 * 10-2
Ka
Large
3.2 * 109
1.0 * 109
1.3 * 106
7.1 * 10-3
7.2 * 10-4
6.6 * 10-4
1.8 * 10-4
6.3 * 10-5
5.4 * 10-5
Chloride
Hydrogen
sulfate ion
Nitrate ion
Water
Hydrogen
oxalate ion
Hydrogen
sulfite ion
Sulfate ion
NO3 H2O
HO2C2O2-
Hydrogen sulfate
ion
Phosphoric acid
HSO4 -
SO4 2-
H3PO4
H2PO4 -
Nitrous acid
Hydrofluoric
acid
Methanoic acid
Benzoic acid
HNO2
HF
NO3 F-
Dihydrogen
phosphate ion
Nitrite ion
Fluoride ion
HCO2H
C6H5COO
H
HO2C2O2-
HCO2 C6H5COO-
Methanoate ion
Benzoate ion
O2C2O2 2-
Oxalate ion
CH3COO
H
CO3 2-
CH3COO
H2S
HS-
H2PO4 -
HPO4 2-
HS-
S2-
Ethanoate
(acetate) ion
Hydrogen
carbonate ion
Hydrogen
sulfide ion
Hydrogen
phosphate ion
Sulfite ion
1.8 * 10-5
4.4 * 10-7
Carbonic acid
1.1 * 10-7
Hydrosulfuric
acid
Dihydrogen
phosphate ion
Hydrogen sulfite
ion
6.2 * 10-8
Name
Perchlorate ion
Iodide
Bromide
HNO3
H3O+
HO2C2O2
H
H2SO3
Hydrogen
oxalate ion
Ethanoic acid
6.3 * 10-8
Base
HSO3 -
HCO3 -
2.9 * 10-8
6.2 * 10-10
5.8 * 10-10
5.8 * 10-10
4.7 * 10-11
4.2 * 10-13
1.8 * 10-13
1.3 * 10-13
1.6 * 10-14
---------
Hypochlorous
acid
Hydrocyanic
acid
Ammonium ion
Boric acid
Hydrogen
carbonate ion
Hydrogen
phosphate ion
Dihydrogen
borate ion
Hydrogen
sulfide ion
Hydrogen borate
ion
water
HClO
ClO-
HCN
CN-
NH4 +
H3BO3
NH3
H2BO3 -
HCO3 -
CO3 2-
Ammonia
Dihydrogen
carbonate ion
Carbonate ion
HPO4 2-
PO4 3-
Phosphate ion
H2BO3-
HBO3 2-
HS-
S 2-
Hydrogen
borate ion
Sulfide ion
HBO3 2-
BO3 3-
Borate ion
H2O
OH-
Hydroxide
Hypochlorite
ion
Cyanide ion
1. Strong acid are listed at the top left hand corner of the table and have Ka
values >1
2. Acid with values less than one are considered weak.
3. The strong bases are listed at the bottom right of the table and get weaker
as we move to the top of the table.
Calculations with weak acids
Rice tables are used to determine the concentrations in weak acids.
Solving Equilibrium Problems Involving Weak Acids
Example: Consider the process by which we would calculate the H 3O+,
OAc-, and HOAc concentrations and pH of the equilibrium in an 0.10 M
solution of acetic acid in water.
Calculate the H3O+ concentration in an 0.0001 M solution of hydrocyanic
acid (HCN).
HCN(aq) + H2O(l) D H3O+(aq) + CN-(aq) Ka = 6 x 10-10
Calculate the pH of an 0.023 M solution of saccharin (HSc), if Ka is 2.1 x 1012
for this artificial sweetener.
In a solution of acetic acid, the equilibrium concentrations are found to be
[CH3COOH] = 1.000; [CH3COO-] = 0.0042. Evaluate the pH of this solution
and the equilibrium constant of ionization of acetic acid.
Its acidity constant Ka is 1.85x 10-5. What is the pH of a concentrated
vinegar, which is a 1.0 M acetic acid solution?
What is the pH of a 0.010 M HCl solution?
What is the pH of a 1.0x 10-4 acetic acid solution (Ka = 1.85x 10-5)?
What is the pH of a 1.0 x 10-3 M chloroacetic acid solution (Ka = 1.4 10-3)?
What is the pH of a 1.0 x 10-6 M chloroacetic acid solution
(Ka = 1.4 x 10-3)?
• What is the pH of a 1.0 x 10-7 M chloroacetic acid solution
(Ka = 1.4 x 10-3)?
Autoionization of Water
Water is an amphoteric substance. It can act as an acid or a base.
Water ionizes:
2H2O(l) D H3O+(aq) + OH-(aq)
giving the following equilibrium expression
Kw = [H3O+] [OH-] = [H+] [OH-]
where Kw is the dissociation constant for water.
At 25 oC Kw = 1.0 x 10-14 M
The following expression is valid in any acid base equilibrium
Kw = [H+] [OH-] = 1.0 x 10-14 M
And pH + pOH = 14
For the acid/base reaction:
HA + H2O à A- + H3O+
Ka x Kb = Kw
What is the [OH-] for a 0.004 M HCl solution?
What is the pOH in a solution that has a pH of 8.0
What is Kb for the acetate ion
HOC2H3O2 + H2O D OC2H3O2 + H3O+ Ka = 1.8 x 10-5
Solving Equilibrium Problems Involving Bases
Calculate the pH of a 0.030 M solution of sodium benzoate
(C6H5CO2Na) in water from the value of Ka for benzoic acid
(C6H5CO2H) is 6.3 x 10-5.
Calculate the HOAc, OAc-, and OH- concentrations at
equilibrium in an 0.10 M NaOAc solution. (HOAc: Ka = 1.8 x
10-5)
1.0 * 103
Table of Acid and Base Strength
Acid
Name
Formula
Formula
Perchloric acid
HClO4
ClO4 Hydroiodic acid
HI
IHydrobromic
HBr
Bracid
Hydrochloric
HCl
Clacid
Sulfuric acid
H2SO4
HSO4 -
2.4 * 101
-------5.4 * 10-2
Nitric acid
Hydronium ion
Oxalic acid
1.3 * 10-2
Sulfurous acid
1.0 * 10-2
Large
3.2 * 109
1.0 * 109
1.3 * 106
7.1 * 10-3
7.2 * 10-4
6.6 * 10-4
1.8 * 10-4
6.3 * 10-5
5.4 * 10-5
Chloride
Hydrogen
sulfate ion
Nitrate ion
Water
Hydrogen
oxalate ion
Hydrogen
sulfite ion
Sulfate ion
NO3 H2O
HO2C2O2-
Hydrogen sulfate
ion
Phosphoric acid
HSO4 -
SO4 2-
H3PO4
H2PO4 -
Nitrous acid
Hydrofluoric
acid
Methanoic acid
Benzoic acid
HNO2
HF
NO3 F-
Dihydrogen
phosphate ion
Nitrite ion
Fluoride ion
HCO2H
C6H5COO
H
HO2C2O2-
HCO2 C6H5COO-
Methanoate ion
Benzoate ion
O2C2O2 2-
Oxalate ion
CH3COO
H
CO3 2-
CH3COO
H2S
HS-
H2PO4 -
HPO4 2-
HS-
S2-
Ethanoate
(acetate) ion
Hydrogen
carbonate ion
Hydrogen
sulfide ion
Hydrogen
phosphate ion
Sulfite ion
1.8 * 10-5
4.4 * 10-7
Carbonic acid
1.1 * 10-7
Hydrosulfuric
acid
Dihydrogen
phosphate ion
Hydrogen sulfite
ion
6.2 * 10-8
Name
Perchlorate ion
Iodide
Bromide
HNO3
H3O+
HO2C2O2
H
H2SO3
Hydrogen
oxalate ion
Ethanoic acid
6.3 * 10-8
Base
HSO3 -
HCO3 -
2.9 * 10-8
6.2 * 10-10
5.8 * 10-10
5.8 * 10-10
4.7 * 10-11
4.2 * 10-13
1.8 * 10-13
1.3 * 10-13
1.6 * 10-14
---------
Hypochlorous
acid
Hydrocyanic
acid
Ammonium ion
Boric acid
Hydrogen
carbonate ion
Hydrogen
phosphate ion
Dihydrogen
borate ion
Hydrogen
sulfide ion
Hydrogen borate
ion
water
HClO
ClO-
HCN
CN-
NH4 +
H3BO3
NH3
H2BO3 -
HCO3 -
CO3 2-
Ammonia
Dihydrogen
carbonate ion
Carbonate ion
HPO4 2-
PO4 3-
Phosphate ion
H2BO3-
HBO3 2-
HS-
S 2-
Hydrogen
borate ion
Sulfide ion
HBO3 2-
BO3 3-
Borate ion
H2O
OH-
Hydroxide
Hypochlorite
ion
Cyanide ion
1. Strong acid are listed at the top left hand corner of the table and have Ka
values >1
2. Acid with values less than one are considered weak.
3. The strong bases are listed at the bottom right of the table and get weaker
as we move to the top of the table.