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Acids and Bases
Earliest Characterizations Æ Acids have a sour taste; bases taste bitter, feel slippery
Modern Definitions
An Arrhenius acid liberates H+ (hydrogen ions, protons) in solution
An Arrhenius base liberates OH- (hydroxide ions) in solution. The classic Arrhenius bases are the Group IA
hydroxides (other metal hydroxides are mostly insoluble in aqueous solution)
A Bronsted-Lowry acid is a proton (H+) donor
A Bronsted-Lowry base is a proton acceptor (any molecule with a lone pair or pairs of electrons
is a potential Bronsted base, e.g. NH3, H2O,
CH3NH2, etc.)
Acid-Base Terminology
We represent a general acid as HA and a general base (B-L) as B:
Consider the equilibrium
HA(aq) +
acid
H2O(l)
base
H3O+ (aq)
conjugate
acid
+
A-(aq)
conjugate
base
The equilibrium expression for this reaction is written as
Ka =
[H 3 O + ][A − ]
[HA]
where Ka is the acid dissociation constant
We often abbreviate acid equilibrium as follows:
for which Ka is written as
For bases,
B: (aq)
+
Ka =
[H + ][A − ]
[HA]
H2O(l)
H+(aq)
HA(aq)
+
A-(aq)
but it is understood that H+(aq) is really H3O+(aq)
BH+(aq)
+
OH-(aq)
[BH + ][OH − ]
for which Kb =
[B :]
Acid Strength
A strong acid has Ka >> 1 (i.e. the equilibrium lies far to the right; the acid is “completely
dissociated (ionized)”)
Common strong acids in water: HCl, HBr, HI, HNO3, HClO4, H2SO4 (first dissociation)
A weak acid has Ka < 1
(i.e. the equilibrium lies far to the left; the acid is mostly undissociated)
Autoionization of Water
Water is amphoteric: it can act as an acid or a base Æ H2O(l) + H2O(l)
For this equilibrium we write
Kw = [H3O+][OH−]
***** For any aqueous solution at 25oC Æ
H3O+(aq) + OH-(aq)
At 25oC, Kw = 1.0 x 10-14
Kw = [H+][OH−] = 1.0 x 10-14 *******
The pH Scale
Whenever you see a “p” in front of a quantity, it means to take the negative base 10 log of that quantity
pH = −log[H+]
pOH = −log[OH-]
pKa = −log Ka
pKw = −log Kw = 14.00 at 25oC
************************************************************************************
Solution Equilibrium -- Words to Live By
Follow these steps to make your "solutions" come out correctly!
1) Write down the ions and molecules that are floating around in the solution before any reaction
(equilibrium or otherwise) occurs.
Fine point: Remember that soluble compounds (strong acids, group I salts, nitrates, etc.) dissociate
"completely" in water. They exist as ions in solution: you should write them down as ions. (For example, you
would write sodium nitrite not as NaNO2, but as Na+ and NO2-)
2) Check to see if any reactions other than equilibrium are possible. Use stoichiometry to take care of such
reactions.
Fine point: A reaction that is "other than equilibrium" is one that goes essentially to completion. A good
example in aqueous solution is H+ + OH- Æ H2O, which would occur if you added a strong base to a strong
or weak acid. You need to use stoichiometry to see which, if any, of the reacting species are left in solution.
3) Write the equilibrium equations for any possible significant equilibrium. Use the ICE table to solve the
equilibrium problem.
Fine point: A significant equilibrium reaction is one with pK appreciably less than that for the dissociation of
the solvent. In aqueous solution, the pK for water dissociation equilibrium is 14, so any reaction with pK ≤ 13
or so would be a significant equilibrium in water.
A final word about ionic salts: Remember -- when you dissolve an ionic salt in water, you need to consider
whether the cation is acidic or the anion is basic (or both). The anion will be basic if it is the conjugate base of a
weak acid. The cation will be acidic if it is the conjugate acid of a weak base or if it is has a small atomic radius
and high charge, in which case it facilitates water hydrolysis when solvated.
For a conjugate acid/base pair ⇒ KaKb = Kw
(really important)
Acid-Base Class Example Problems
Strong Acid/Strong Base (Key Point: complete dissociation; no ICE chart necessary)
1) What is the pH of a 2.5 x 10-4 M HCl solution? Of a 2.5 x 10-4 M NaOH solution?
Of a 7.7 x 10-11 M HCl solution?
Weak Acid (Key Point: Small Ka, so ICE chart necessary)
2) What is the pH of 1.5 M solution of HNO2? (Ka = 4.0 x 10-4)
3) A 0.100 M HF solution has 8.1% dissociation. What is Ka for HF?
Polyprotic Acid (Key Point: H+ from successive dissociations is usually insignificant, but becomes
more significant the smaller the initial acid concentration)
4) What is the pH of a 0.500 M H2SO3 solution? (Ka,1 = 1.5 x 10-2; Ka,2 = 1.0 x 10-7)
Weak Base (Key Point: Small Kb, so ICE chart necessary)
5) What is the pH of a 0.0050 M hydrazine solution? (Kb = 3.0 x 10-6)
Acid-Base Properties of Salts
When a salt (ionic compound) dissolves in water:
Æ you need to consider if the cation is the conjugate acid of a weak base
Æ you also need to consider if the anion is the conjugate base of a weak acid
*** Cations and anions that are conjugates of weak bases or acids will affect the solution pH ****
If the anion is the conjugate base of a weak acid, the solution will be weakly basic
If the cation is the conjugate acid of a weak base, the solution is weakly acidic
(If cation is acidic and anion is basic, larger K value determines acidity or basicity)
Super important equations:
For an acid/conjugate base pair Æ
For a base/conjugate acid pair Æ
Ka x Kb,conj. = Kw
Kb x Ka,conj. = Kw
*** These equations allow you to calculate the Ka of an acidic anion or the Kb of a basic cation ***
Examples
6) What is the pH of a solution with [NaNO3] = 0.125 M? How about [NaNO2] = 0.125 M?
7) What is the pH of the solution when 1.00 g of anilinium chloride (C6H5NH3Cl; 129.60 g/mol) is dissolved in
enough water to make 100.0 mL of solution?
8) Rank the following 1.0 M solutions in order of increasing pH:
NaNO3
KF
LiOH
HBr
9) Would a solution of ammonium nitrite be acidic or basic?
NH4Cl
Acidic Cations
Small, highly charged metal cations (e.g. Al3+, Bi3+, Fe3+, etc) facilitate water hydrolysis and result in acidic
solutions. In solution, cations that are small and highly charged act as Lewis acids. A Lewis acid is an electron
pair acceptor. When an H2O molecule, acting as a Lewis base (an electron pair donor), sticks to a small cation, the
H2O molecule is more easily able to lose H+, because the remaining OH− is stabilized by the positive cation
charge.
10) The equilibrium Fe(H2O)63+ (aq)
the pH of a 0.25 M Fe3+ solution?
Fe(H2O)5(OH)2+(aq) +
H+(aq) has Ka = 6.5 x 10-3. What is
Acid-Base Properties of Oxides
Æ When a covalent oxide (oxygen bonded to a nonmetal) dissolves in water, an acidic solution results
Examples: CO2, NO2, SO3, etc.
Æ When an ionic oxide (oxygen bonded to a metal) dissolves in water, a basic solution results
Examples: Na2O, CaO, etc.
Effects of Structure on Acid-Base Properties
Binary Acids:
Acid strength increases down a group (Why? Even though bond polarity is decreasing, the bond
strength is also decreasing)
Acid strength increases left-to-right across a period
Oxyacids:
Acid strength increases as:
1) the number of oxygens increases (e.g. Ka HClO3 > Ka HClO2)
2) the oxidation state of the nonmetal atom increases, given the same number
of oxygens (e.g. Ka HClO4 > Ka H2SO4)
3) the electronegativity of the nonmetal atom increases (only for the hypohalous
acids HOX, where X is a halogen)
Rank the following 1.0 M solutions in order of increasing pH (10 pts)
a) HClO2
HIO
b) H2O
NO2
c) NaOCl
NH4Cl
KI
KOH
HClO
NaClO2
HNO3
Na2O
HCN
(Ka HOCl = 3.5 x 10-8; Kb NH3 = 1.8 x 10-5; Ka HCN = 6.2 x 10-10)