Acids, Bases, and Acid-Base
Equilibria
Brown and LeMay
Chapters 16 and 17
16.1, 16.2 – Acid-Base Theories and Relative
Strengths
Arrhenius Theory of acids and bases
• acid – produces H+ ions
• base – produces OH- ions
• Strong acids and bases ionize completely
• Problems with this theory: It’s restricted to aqueous
solutions and it doesn’t include bases like NH3 which
don’t directly ionize to yield OH-
• Bronsted-Lowry Theory of Acids and Bases
• Acids are proton (H+) donors (A Brønsted–Lowry acid
must have at least one removable (acidic) proton (H+)
to donate)
• Bases are proton (H+) acceptors (A Brønsted–Lowry
base must have at least one nonbonding pair of
electrons to accept a proton (H+))
• Conjugate acid: formed when a base accepts a proton
• Conjugate base: formed when an acid donates a proton
• Conjugate acid-base pair: an acid and a base that differ by
only one H+
• On each side of the equation, there’s an acid and a base
• See sample exercises 16.1 and 16.2
• Amphiprotic – a substance that can act as an acid or a base
• Ex. H2O
• As a base →
• As an acid →
• The relative strengths of acids and bases:
– The stronger an acid is, the weaker its conjugate base is
– The stronger a base is, the weaker its conjugate acid is
• In every acid-base
reaction, equilibrium
favors transfer of the
proton from the stronger
acid to the stronger base
to form the weaker acid
and the weaker base.
• See sample exercise 16.3
16.3 – The Self-ionization of Water
• Kw (ion product constant of water) = H3O+xOH- = 1 x 10-14
at 25oC
•
H3O+ = OH- = 1 x 10-7 for pure water (at 25°C)
• If a solution is neutral, [H+] = [OH–].
• If a solution is acidic, [H+] > [OH–].
• If a solution is basic, [H+] < [OH–].
• To calculate the concentration of H+ or OH- when you only
know one of them, use the equation H3O+xOH- = 1 x 10-14
• See sample exercises 16.4 and 16.5
16.4 – The pH Scale
• pH: it’s a logarithmic scale:
•
pH = -log H3O+
•
pOH = -log OH-
•
•
pKw = -log Kw = 14 (at 25°C)
•
pKw = pH + pOH = 14
•
•
•
•
•
Neutral solutions have a pH = 7
Acidic solutions have a pH<7
Basic solutions have a pH>7
Note: when the pH changes by 1, the H3O+ changes 10 fold.
See sample exercises 16.6 and 16.7
• Measuring pH: 2 main ways of measuring pH –
1) with a pH meter – has electrodes that indicate small changes
in voltage to detect pH.
2) With acid-base indicators (substances that can exist in either
an acid or a base form, and are different colors at different
pHs)
sample rxn: HIn + H2O
H3O+ + In• Works according to
LeChatelier’s Principle
16.5 – Strong Acids and Bases
• Strong Acids
• Remember that the seven strong acids are HCl, HBr, HI, HNO3,
H2SO4, HClO3, and HClO4.
• These are strong electrolytes and exist totally as ions in
aqueous solution; e.g.,
HA + H2O → H3O+ + A–
• So, for the monoprotic strong acids,
•
[H3O+] = [acid]
• So the pH of a strong monoprotic acid = -log[acid]
• See sample exercise 16.8
• Strong Bases
• Strong bases are the soluble hydroxides, which are the alkali metal and
heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+).
• Again, these substances dissociate completely in aqueous solution; e.g.,
MOH(aq) → M+(aq) + OH–(aq) or
M(OH)2(aq) → M2+(aq) + 2 OH–(aq)
• For the alkali metal hyroxides:
•
pOH = -log[OH-] = -log[base]
•
pH = 14 – pOH
• For the alkaline earth metal hydroxides:
•
pOH = -log[OH-] = -log{2x[base]}
• See sample exercise 16.9
16.6 – Weak Acids
• For a weak acid, the equation for its dissociation is
HA(aq) + H2O(l) ⇌ H3O+(aq) + A–(aq)
• Since it is an equilibrium, there is an equilibrium constant
related to it, called the acid-dissociation constant, Ka:
•
•
• The greater the value of Ka,
the stronger the acid is.
• Calculating Ka from pH:
• To find Ka , use the pH to calculate the equilibrium quantity of
the H3O+, then use an ICE chart
• see sample exercise 16.10
• Calculating Percent Ionization:
• Percent ionization is a measure of acid strength (the stronger
the acid, the greater the % ionization)
• Percent ionization =
 100
• See sample exercise 16.11
• Calculating pH from Ka:
• in these problems, you will be given the Ka and the initial
concentration of the acid
• Set up an ICE chart using “x” for both the [H3O] and the [A-].
• Shortcut (to avoid quadratic): If Macid /Ka > 100, then ignore
the x in (M-x) concentrations at equilibrium.
• Since “x” is the [H3O] at equilibrium, you get the pH by taking
the -log [H3O]
• See sample exercises 16.12 and 16.13
• Polyprotic Acids:
• Monoprotic acid – 1 ionizable H
• Polyprotic acid – more than 1 ionizable H. The ionizations
occur separately. Each step has its own Ka. The first Ka is the
largest. If the Ka values for the first and second dissociation
differ by a factor of 103 or more, the pH generally depends
only on the first dissociation.
• See sample exercise 16.14
16.7 – Weak Bases
• Like weak acids, weak bases have an equilibrium constant
called the base dissociation constant (Kb).
• Equilibrium calculations work the same as for acids, using the
base dissociation constant instead.
• Since “x” in this case is the [OH-] at equilibrium, taking the
-log of it gives you the pOH.
• See sample exercise 16.15
• Types of Weak Bases:
• Two main categories
1) Neutral substances with an atom
that has a nonbonding pair of electrons
that can accept H+ (like ammonia
and the amines)
2) Anions of weak acids
For example, ClO- is the conjugate base of the weak acid
HClO. So ClO- is a weak base.
• See sample exercise 16.16
16.8 – Relationship between Ka and Kb
For a conjugate acid–base pair, Ka and Kb are related in this way:
Ka × Kb = Kw
Therefore, if you know one of them, you can calculate the other.
Also, pKa + pKb = pKw = 14 for a conjugate acid-base pair
See sample exercise 16.17
16.9 – Acid-Base Properties of Salt Solutions
• Many ions react with water to create H+ or OH–. The reaction
with water is often called hydrolysis.
• To determine whether a salt is an acid or a base, you need to
look at its cation and anion separately.
• Anions: An anion, A- can be considered to be the conjugate
base of an acid.
• If the acid, HA is a strong acid, then its anion will be too weak
to react with water to produce OH- ions and therefore does
not affect the pH.
• If the acid is a weak acid, then its conjugate base is a weak
base, and it can react with water to produce OH- ions which
increases the pH of the solution
• Cations:
• Group 1 or 2 metal ions do not hydrolyze and are therefore
neutral
• Polyatomic cations are the conjugate acids of a weak base;
e.g., NH4+. They can react with water to form H3O+ ions.
• Some small, highly charged cations (like Fe3+, Al3+, and Cr3+)
because they become hydrated, weakening the O-H bond of
water, and enabling the transfer of a H+ to another water,
forming H3O+
• Combined effect of cation and anion – possible combinations:
1) Salts containing ions from strong acids and strong bases form
neutral solutions. (ex. NaCl, KNO3, BaI2)
2) Salts containing ions from weak acids and strong bases form
basic solutions. (ex. Na2CO3, KNO2, NaCH3COO)
3) Salts containing ions from strong acids and weak bases form
acidic solns. (ex. NH4Cl, NH4NO3)
4) Salts containing ions from weak acids and weak bases (or
containing ions like Fe3+) –depends on the relative acid and
base strength (ex. NH4CN, NH4NO2, CrF3)
• See sample exercise 16.18
16.10 – Acid-Base Behavior and Chemical
Structure
• Factors that Affect Acid Strength:
1) The H—A bond must be polar with δ+ on the H atom and δ–
on the A atom
2) Bond strength: Weaker bonds can be broken more easily,
making the acid stronger.
3) Stability of A–: More stable anion means stronger acid.
The strength of an acid is often a combination of all three
factors.
• Binary Acids:
• Binary acids consist of
H and one other element.
• Within a group, H—A bond
strength is generally the
most important factor (acid
strength increases going
down a group).
• Within a period, bond
polarity is the most
important factor to
determine acid strength
(acid strength increases
going across a period)
• Oxyacids: acids in which OH groups and possible additional
oxygen atoms are bound to a central nonmetal atom.
• As the electronegativity of the nonmetal increases, the O-H
bond gets weaker, and the acid gets stronger (see figure 16.19
on page 707)
• Also, as additional atoms high in electronegativity bond to the
center atom, the acid strength increases:
• Carboxylic Acids: have a –COOH (carboxyl) group. The
electronegativity of the R group attached to the carboxyl
group determines the strength. The greater the
electronegativity, the stronger the acid.
•
Ka = 1.6 x 10-5
Ka = 2.3 x 10-1