Acids and Bases Part I - This is ricksobers.com

Acids and Bases Part I
Dr. Sobers’ Lecture Notes
Robert Boyle
Robert Boyle 1661
Acid Properties:
Sour taste
Corrosive
Cause change in color in
some vegetable dyes.
(makes blue litmus red)
Lose acid properties when
combined with bases (alkalies)
Robert Boyle 1661
Base Properties:
Feel slippery
Cause change in color in
some vegetable dyes.
(makes red litmus blue)
Lose acid properties when
combined with acids
Another property is that they taste bitter
Arrhenius Definition of
Acids and Bases
Arrhenius Definitions
Svante Arrhenius
1884 Definitions
Arrhenius acids are substances that ionizes
in water to give H+ (H3O+) ions.
H2SO4 + H2O → H3O+ + HSO4Arrhenius bases are substances that produce hydroxide ions
in water.
NaOH → Na+ + OHH 2O
Arrhenius Definitions
Svante Arrhenius
Ammonia is often described as forming
hydroxide in solution and so fits this
definition:
NH3(aq) + H2O(l) ⇄ NH4+(aq) + OH-(aq)
At the time ammonia was thought to form this in solution but the
equilibrium actually favors the reactants.
Arrhenius Definitions
Svante Arrhenius
This theory doesn’t explain acid-base
reactions that are not in aqueous solution:
NH3(g) + HCl(g) ⇄ NH4Cl(s)
Brønsted-Lowry
Definitions
Brønsted-Lowry Definitions
Johannes Nicolaus Brønsted
Thomas Martin Lowry
Independently derived theory in 1923
Brønsted-Lowry acids are any species that donate a proton, H+.
Brønsted-Lowry bases are any species that accept a proton, H+.
The definition is not limited to aqueous solutions.
Brønsted-Lowry Definitions
NH3(g) + HCl(g) ⇄ NH4Cl(s)
Ammonia accepts a proton (a base)
Hydrogen chloride donates a proton (an acid)
Brønsted-Lowry Definitions
Conjugate Acids and Bases
Conjugate acid: when a base accepts a proton, it becomes an acid
capable of returning that proton.
Conjugate base: when an acid donates its proton, it becomes capable
of accepting that proton back.
Provide the conjugate acid or base of each formula:
Conjugate Acid
Conjugate Base
HF
F-
H2O
OH-
HCO3-
CO32-
H3O+
H2O
NH3
NH2-
NH4+
NH3
CH3COOH
CH3COO-
They differ by H+
Review: Behavior of
Acids and Bases in Water
Strong and Weak
Acids and Bases in Water
Strong acids ionize in water 100%:
HCl(aq) + H2O(l) → H3O+ (aq) + ClThe conjugate base is very weak; a negligible base.
An acid considered strong, may not be in non-aqueous solvents.
Weak acids ionize in water less than 100%:
HNO2(aq) + H2O(l) ⇄ H3O+ (aq) + NO2-(aq)
An equilibrium exists. The conjugate base is a weak base.
Strong and Weak
Acids and Bases in Water
Strong bases dissociate in water:
NaOH(s) → Na+ (aq) + OH- (aq)
Strong bases are not Brønsted-Lowry bases but the
hydroxide ion produced is. (It can accept a proton)
Dissociation is used to describe this process. The
pure sodium hydroxide is ionic already.
Ionization described a reaction that produces new
ions. Pure HCl is a molecular gas for instance. But
it produces ions in water.
Strong and Weak
Acids and Bases in Water
Strong bases dissociate in water:
NaOH(s) → Na+ (aq) + OH- (aq)
Weak bases ionize in water less than 100%:
NH3(aq) + H2O(l) ⇄ OH- (aq) + NH4+(aq)
An equilibrium exists. The conjugate acid is a weak acid.
Negligible bases do not react. The conjugate acid is very strong.
Cl-(aq) + H2O(l) → no reaction
The Auto-ionization of
Water and the pH Scale
Amphiprotic/Amphoteric
Amphoteric - it may act as an acid or a base.
Amphiprotic - it may accept or donate a proton, H+.
Water and ammonia are examples.
Other examples: hydrogen sulfate ion, hydrogen carbonate ion
The Water Ion Product
Water auto-ionization:
2H2O ⇄ H3O+ + OH-
Equilibrium expression: Kw =
The water ion product:
[H3O+][OH-]/[H2O]
Kw ≈ [H3O+][OH-]
Kw = 1x10-14
Neutral water:
[H3O+] = [OH-] = 1x10-7M
The Water Ion Product
Example: Calculate the concentration of hydronium and
hydroxide ions in a 0.100M HCl solution.
HCl is a strong acid. It ionizes 100%
HCl(aq) + H2O(l) → H3O+ (aq) + ClThe hydronium ion is the concentration of HCl itself:
[H3O+] = 0.100M
Use the water-ion product equation to get [OH-]
1x10-14 = [H3O+][OH-]
The Water Ion Product
Example: Calculate the concentration of hydronium and
hydroxide ions in a 0.100M HCl solution.
Use the water-ion product equation to get [OH-]
1x10-14 = [H3O+][OH-]
1x10-14 = (0.100)[OH-]
[OH-] = 1x10-13 M
The Water Ion Product
For strong acids and bases (assuming no solubility issues), the
concentration of hydronium and hydroxide are easy to find.
The hydronium ion concentration in a 1.000M HBr solution is
1.000M.
The hydroxide ion concentration in a 1.000M NaOH solution is
1.000M.
If the hydronium ion or hydroxide ion concentration is known,
then they are both known:
1x10-14 = [H3O+][OH-]
For weak acids and bases, it is not so simple.
pH and pOH
The pH is the cologarithm of the hydronium ion concentration:
pH = -log([H3O+])
The function “-log( )” is the called the “p” function.
The pOH is defined the same way:
pOH = -log([OH-])
How are pH and pOH mathematically calculated from one another?
pH and pOH
Kw = [H3O+][OH-]
Apply “-log( )” to both sides:
-log(Kw) = -log([H3O+][OH-])
-log(1x10-14) = -log([H3O+]) + -log([OH-])
(14)
(pH)
(pOH)
14 = pH + pOH
Kw = [H3O+][OH-]
[OH-]
[OH-] = 10-pOH
[H3O+] = 10-pH
pH
pOH
14 = pH + pOH
pOH = -log([OH-])
pH = -log([H3O+])
[H3O+]
The pH Scale
pH
[H3O+]
[OH-]
pOH
7
1x10-7M
1x10-7M
7
Acidic Solutions:
<7
> 1x10-7M
< 1x10-7M
>7
Basic Solutions:
>7
< 1x10-7M
> 1x10-7M
<7
Neutral Solutions:
Given the pH, pOH, hydronium or hydroxide concentration,
calculate the rest for each solution.
pH
[H3O+]
[OH-]
pOH
7
1x10-7M
1x10-7M
7
4
1x10-4M
1x10-10M
10
-1
10M
1x10-15M
15
11
1x10-11M
1x10-3M
3
0
1M
1x10-14M
14
9.5
3.1x10-10M
3.2x10-5M
4.5