Acid and Base EquilibriaThe concept of acidic and basic solutions is perhaps one of the most important topics in chemistry. Acids and bases affect the
properties of foods, biochemical reactions, pharmaceuticals, and industrial materials.
Acid and Base Equilibria
• Properties of Acids• Sour or tart taste.• Corrosive (deteriorate).• Electrolytes• Electrolytes are able to conduct an
electrical current because of the presence of ions in aqueous solutions.
Acid and Base Equilibria• Properties of Acids (cont.)• Will react with most metals to form
hydrogen gas.• Some acids are ‘stronger’ than others• All acids contain a hydrogen that they
can give away.
HA(aq) + H2O(l) H3O+(aq) + A-
(aq)
Acid and Base Equilibria• Properties of Bases• Bitter tasting• Slippery• Caustic – They will degrade biological
tissue. Chemical burns from strong bases are nasty.
• Bases form the hydroxide ion (OH-) in water.
Acid and Base Equilibria• Arrhenius Acids • Svante Arrhenius (1900) defined an acid.• Acids are hydrogen containing compounds
that yield a hydrogen ion (H+) in water.• An Arrhenius acid donate an H+ ion.
HCl(aq) + H2O(l) H3O+(aq) + Cl-
(aq)
Acid and Base Equilibria• Arrhenius Acids • HNO3(aq) + H2O(l) H3O+
(aq) + NO3-(aq)
• H2SO4(aq) + H2O(l)
• H2O(l) + H2O(l)
Acid and Base Equilibria
• Arrhenius Acids • Acids that have one hydrogen ion to
donate are called monoprotic.• Acids that have 2 hydrogen ions to
donate are called diprotic.• Acids that have 3 hydrogen ions to
donate are called ____protic.
Acid and Base Equilibria
• Arrhenius Bases • Compounds that produce the
hydroxide ion(s) (OH-) in water are called Arrhenius Bases.
• NaOH(aq)
Acid and Base Equilibria
• Arrhenius Acids and Bases • Problem with the Arrhenius
definition; Some bases can form OH- ions in solution but not have an OH- ion in the chemical formula.
NH3(g) + H2O(l) NH4+
(aq) + OH-(aq)
Acid and Base Equilibria• Bronsted-Lowry Acids and Bases• A better definition of an acid and a base• Bronsted-Lowry Acid – A molecule that donates
an H+ to another molecule.• Bronsted-Lowry Base – A molecule that accepts
an H+ from the B-L acid.
Acid and Base Equilibria• Bronsted-Lowry Acids and Bases• Identify the Bronsted-Lowry acid and base;
NH3(g) + H2O(l) NH4+
(aq) + OH-(aq)
Acid and Base Equilibria
• Bronsted-Lowry Acids and Bases• Conjugate Acid – becomes the H+ donor in
the reverse reaction.• Conjugate Base – becomes the H+
acceptor in the reverse reaction.• Identify the BL acid, base, conjugate acid
and conjugate base;HClO2(aq) + H2O(l) H3O+1
(aq) + ClO2-(aq)
Acid and Base Equilibria
• The Ion-Product of WaterH2O(l) H3O+
(aq) + OH-(aq)
Keq = Kw =
Kw = 1.0 x 10-14
Acid and Base Equilibria
• The Ion-Product of WaterH2O(l) H3O+
(aq) + OH-(aq)
Keq = Kw = [H3O+] [OH-] = 1.0 x 10-14
• What would be the concentration of [H3O+] [OH-] in pure water?
Acid and Base Equilibria
• What would be the concentration of [OH-] in a solution where [H3O+] = 0.025 M?
Acid and Base Equilibria
• The pH Scale – Measures the concentration of H+ ion in an aqueous solution.
Acid and Base Equilibria• The pH Scale • Remember that a water molecules
ionizes;H2O(l) H+
(aq) + OH-(aq)
• In pure water, the concentration of H+ and OH- each is 1.0 x 10-7M
Acid and Base Equilibria• The pH Scale
H2O(l) H+(aq) + OH-
(aq)
• Therefore, the product of [H+] and [OH-] must be equal to 1.0 x 10-14M2.
• An aqueous solution will always have the concentration of H+ and OH- equal to 1.0x10-14M2.
Acid and Base Equilibria• The pH Scale
H2O(l) H+(aq) + OH-
(aq)
Keq = [H+] x [OH-]
Kw = [1.0x10-7] x [1.0x10-7] = 1.0x10-14
Kw is called the autoionization constant for water.
Acid and Base Equilibria• The pH Scale
Keq = [H+] x [OH-]
[1.0x10-7] x [1.0x10-7] = 1.0x10-14
• If the addition of an acid makes the [H+] increase, then the [OH-] will decrease.
• If the addition of a base makes the [OH-] increase, the [H+] will decrease.
Acid and Base Equilibria• The pH Scale • Calculate the [H+] if 0.05 moles of HCl
is added to 1.0 L of water.
Acid and Base Equilibria• The pH Scale • pH is a measure of the concentration
of H+.
pH = -log[H+]
Acid and Base Equilibria• The pH Scale • What is the pH of an aqueous solution
where [H+] = 1.0 x 10-7?
• What is the pH of an aqueous solution where [H+] = 1.0 x 10-2?
Acid and Base Equilibria• The pH Scale • What is the pH of an aqueous solution
where [H+] = 2.7 x 10-1?
• What is the pH of an aqueous solution where [H+] = 8.0 x 10-12?
Acid and Base Equilibria• The pH Scale • Since Kw = 1 x 10-14 = [H+] x [OH-]
14 = pH + pOH
Acid and Base Equilibria• The pH Scale • Calculate the pOH of an aqueous
solution that has an [H+] = 1.0 x 109.
Acid and Base Equilibria• The pH Scale • Calculate the [H+] of an acid with a pH
of 4.
• Calculate the [H+] of a base with a pH of 10.8.
Acid and Base Equilibria• The pH Scale • Calculate the [OH-] of an acid with a
pH of 2.
• Calculate the [OH-] of a base with a pH of 12.9.
Acid and Base Equilibria• Strong Versus Weak Acids• What makes some acids ‘strong’ and
some ‘weak’?
Acid and Base Equilibria• Strong Versus Weak Acids• We can quantify the relative strength
of an acid by using it’s equilibrium expression (Ka).
• Ka = [products]x = [H+] x [conj. base] [reactants]y [acid]
Acid and Base Equilibria• Strong Versus Weak Acids• Write the Ka expression for HCl.
• Ka =
Acid and Base Equilibria• Strong Versus Weak Acids• Write the Ka expression for H3PO4.
• Ka =
Acid and Base Equilibria• Strong Versus Weak Acids• As the [H+] increases, Ka increases.
Therefore the greater the value of Ka, the more [H+] present, the stronger the acid.
Acid and Base Equilibria• Strong Versus Weak Acids• Calculate the Ka of a 0.10 M solution
of acetic acid that has a pH of 5.0.
Acid and Base Equilibria• Salt Hydrolysis• Sometimes an ion from a salt can
make an aqueous solution acidic or basic.
• What happens when sodium bicarbonate dissolves in water?
Acid and Base Equilibria• Salt Hydrolysis• NaHCO3(aq) Na+
(aq) + HCO3-(aq)
• HCO3-(aq) + H2O(l) H2CO3(aq) + OH-(aq)
Now there will more OH- than H+ in the solution making it basic.
Acid and Base Equilibria• Salt Hydrolysis• Will a solution of ammonium chloride,
NH4Cl, be acidic or basic?
Acid and Base Equilibria• Buffers• A buffer is a solution that resists
changes in pH when either an acid or a base is added.
• Buffers consist of either a weak acids with one of its salts, or a weak base with one of its salts.
Acid and Base Equilibria• Buffers• For example, if a solution is make by
dissolving carbonic acid (weak acid) and sodium bicarbonate (salt of the acid) we get the following;
H2CO3(aq) + H2O(l) H3O+(aq) + HCO3
-1(aq)
Acid and Base Equilibria• Buffers
H2CO3(aq) + H2O(l) H3O+(aq) + HCO3
-1(aq)
• If we add a base to this buffered solution, the H3O+ will scoop it up.
H3O+(aq) + OH(aq)
- 2H2O(l)
Acid and Base Equilibria• Buffers
H2CO3(aq) + H2O(l) H3O+(aq) + HCO3
-1(aq)
• If we add an acid to this buffered solution, the HCO3
-1 will scoop it up.
H+(aq) + HCO3
-1 (aq)
H2CO3(aq)
Acid and Base Equilibria• Buffers• Write the chemical reaction for the
phosphoric acid – dihydrogen phosphate buffer reaction.