Consider the two examples:
When HCl is dissolved in water, essentially 100% of the H+ ions are transferred to the water to form hydronium ions:
HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq)
In contrast, acetic acid, HC2H3O2, only partially transfers its hydrogen ion to water, so the reaction must be considered an equilibrium:
HC2H3O2(aq) + H2O(l) → ← H3O+(aq) + C2H3O2–(aq)
We distinguish between these two kinds of acids as strong or weak:
Strong acids completely transfer their hydrogen ion to water in aqueous solution. All reactions involving strong acids in aqueous solution go to completion.
Weak acids only partially transfer their hydrogen ion to water in aqueous solution. These reactions are equilbria.
Similar observations and statements can be made for bases.
Because bases are often ionic salts, writing their Bronsted–Lowry acid base reactions can be a little tricky:
NaOH(aq) + HOH(l) → HOH(l) + Na+(aq) + OH–(aq)
The color-coding shows the hydrogen ion transfer, but since water shows up on both sides, it cancels giving the net equation:
NaOH(aq) → Na+(aq) + OH–(aq)
Thus, when we write Brønsted–Lowry for ionic salts, even though the H+ transfer may not be explicitly written, we understand that it does take place.
For bases that are not ionic salts, this problem does not exist:
C6H5NH2(aq) + H2O(l) → ← C6H5NH3+(aq) + OH–(aq)
(Aniline)
Fortunately, there are a very limited number of strong acids and bases; we must memorize these:
Strong Acids
HCl Hydrochloric acid
HBr Hydrobromic acid
HI Hydroiodic acid
HNO3 Nitric acid
HClO4 Perchloric acid
H2SO4 Sulfuric acid (1st hydrogen ion only)
Strong Bases
Group 1 Hydroxides
LiOH Lithium hydroxide
NaOH Sodium hydroxide
KOH Potassium hydroxide
RbOH Rubidium hydroxide
CsOH Cesium hydroxide
Group 2 Hydroxides (except Be(OH)2)
Mg(OH)2 Magnesium hydroxide
Ca(OH)2 Calcium hydroxide
Sr(OH)2 Strontium hydroxide
Ba(OH)2 Barium hydroxide
The group 2 hydroxides are not very water soluble so are not often used as bases.
In any acid–base reaction we write, we must correctly distinguish if the reaction goes to completion or if it is an equilibrium.
Any reaction that uses either a strong acid or a strong base goes to completion.
All other acid–base reactions are equilibria.
Complete and balance and label the acid, base, conjugate acid, and conjugate base:
HBr(aq) + NH3(aq) ?
H3PO4(aq) + F–(aq) ?
In aqueous solution, the consequences of any acid–base chemistry are observed by the changes in the properties of water.
The autoionization of water, in the Brønsted–Lowry context is given by:
H2O(l) + H2O(l) → ← H3O+(aq) + OH–(aq)
What are the concentrations of hydroxide and hydronium ion in pure water at 25 °C?
Now suppose we prepare a 0.010 M solution of nitric acid in water. What are the hydronium ion and hydroxide ion concentrations at 25 °C?
There are now two reactions to consider:
HNO3(aq) + H2O(l) → H3O+(aq) + NO3– (aq)
2 H2O(l) → ← H3O+(aq) + OH– (aq)
The strong acid contributes 0.010 M hydronium ion because the reaction goes to completion.
This causes a shift in the water ionization equilibrium towards reactants so the amount of H3O+ and OH– contributed by the water is less than 1.0×10–7 M. (LeChatelier's Principle)
Since the amount of hydronium ion contributed by the water is so small, we can ignore it:
[H3O+]e = 0.010 + (<1.0×10–7) = 0.010 M
The hydroxide ion concentration is found from Kw:
Kw = [H3O+]e[OH–]e = 1.0×10–14
[0.010][OH–]e = 1.0×10–14
[OH–]e = 1.0×10–12 M
The effect of the strong acid is to increase the concentration of hydronium ion and suppress the concentration of hydroxide ion.
Suppose a 1.0×10–6 M solution of NaOH is prepared. What are the hydronium ion and hydroxide ion concentrations at 25 °C?
Again, we need to consider two reactions:
NaOH(aq) → Na+(aq) + OH–(aq)
2 H2O(l) → ← H3O+(aq) + OH–(aq)
The strong base contributes 1.0×10–6 M hydroxide.
The water contributes less than 1.0×10–7 M hydroxide, so the total is
[OH–] = 1.0×10–6 + <1.0×10–7 ~ 1.0×10–6 M
Kw = [H3O+]e[OH–]e = 1.0×10–14
[H3O+]e[1.0×10–6] = 1.0×10–14
[H3O+]e = 1.0×10–8 M
Suppose that a 1.0×10–7 M solution of KOH were prepared. What are the hydronium ion and hydroxide ion concentrations at 25 °C?
In general, if the concentration of a strong acid or strong base exceeds 1.0×10–6 M, then the effects of the water autoionization equilibrium can be ignored.
If the concentration of the strong acid or strong base is less than 1.0×10–6 M, then a full equilibrium calculation must be done.