Acid and Base Equilibria

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)


• 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.


• Arrhenius Bases • Compounds that produce the

hydroxide ion(s) (OH-) in water are called Arrhenius Bases.

• NaOH(aq)


• 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)


• 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)


• The Ion-Product of WaterH2O(l) H3O+

(aq) + OH-(aq)

Keq = Kw =

Kw = 1.0 x 10-14


• 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?


• What would be the concentration of [OH-] in a solution where [H3O+] = 0.025 M?


• 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.


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.


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


-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


-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.

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Acid and Base Equilibria