Chapters 14 and 15

Acids and Bases

• Aqueous solutions have a sour taste.• Acids change the color of acid-base indicators.• Some acids react with active metals to release

hydrogen gas, H2.• Acids react with bases to produce salts and

water (i.e. a neutralization reaction)• Some acids conduct electric current.

Properties of acids

Properties of bases• Aqueous solutions taste bitter.• Bases change the color of acid-base indicators.• Dilute aqueous solutions of bases feel slippery

(ex. soap).• Bases react with acids to produce salts and

water. (i.e. neutralization reactions)• Bases conduct electric current.

Theories of acids and basesTheory Acid

propertiesBase properties

Arrhenius H+ ion produced

OH- ion produced

Brønsted-Lowry

Proton donor Proton acceptor

Lewis e- pair acceptor

e- pair donor

• Strong acids and bases are those that ionize completely in aqueous solution. They are also typically strong electrolytes.

• Weak acids and bases are those that do not ionize completely in aqueous solution. They are also typically weak electrolytes.

Strong vs. weak

Strong acids Weak acidsH2SO4 + H2OH3O+ + HSO4

- HSO4- + H2O H3O+ + SO4

2-

HClO4 + H2O H3O+ + ClO4- H3PO4

+ H2OH3O++ H2PO4-

HClO3 + H2O H3O+ + ClO3- HF + H2O H3O+ + F-

HCl + H2O H3O+ + Cl- CH3COOH +H2OH3O+ + CH3COO-

HNO3 + H2O H3O+ + NO3- H2CO3

+H2OH3O+ + HCO3-

HBr + H2O H3O+ + Br- HCO3- + H2OH3O+ + CO3

2-

HI + H2O H3O+ + I- H2S + H2O H3O+ + HS-

HCN + H2O H3O+ + CN-

Page 474 of your textbook

Common strong and weak acids

Strong bases Weak basesCa(OH)2 Ca2+ +

2OH-

NH3 + H2O NH4+ + OH-

Sr(OH)2 Sr2+ + 2OH- C6H5NH2+ H2O C6H5NH3+ +

OH-

Ba(OH)2 Ba2+ + 2OH-

NaOH Na+ + OH-

KOH K+ + OH-

RbOH Rb+ + OH-

CsOH Cs+ + OH-

Common strong and weak bases

p. 475 of text

Conjugate acid-base reactions

HF (aq) + H2O (l) F- (aq) + H3O+ (aq)

c.a.1 c.b.2 c.b.1 c.a.2

Amphoteric substances are

those thatcan react as either an

acid or a base.

Molecules containing –OH groups (hydroxyl groups) can be acidic or amphoteric. The more oxygen atoms, more polar more acidic

Oxyacids

Oxyacids of chlorine

Ka for weak acids

Acid-dissociation equilibrium constant (Ka) - A measure of the relative strength of an acid. For the generic acid dissociation reaction with water,

HA (aq) + H2O (l) H3O+ (aq) + A- (aq)

Ka = [H3O+][A- ]

[HA]As the Ka value of an acid increases, so does the strength of the acid. By definition:

strong acid: Ka > 1

weak acid: Ka < 1

strong acid + water weaker acid + weaker base (water acts as a strong base)

weak acid + water ATTEMPTS TO CONVERT TO stronger acid + stronger base(water acts as a weak base) BUT the reaction cannot naturally proceed in this direction. This is why strong acids dissociate nearly completely whereas weak acids dissociate only slightly.

The larger the value of Ka, the stronger is the acid.

Ka is a better measure of the strength of an acid than pH because adding more water to the acid solution will not change the value of the equilibrium constant Ka, but it will change the H+ ion concentration on which pH depends.

Kb for weak basesBase dissociation constant or equilibrium constant, Kb

For the reaction in which the Arrhenius base, BOH, dissociates to form the ions OH- and B+:

BOH OH- + B+ For a Brønsted-Lowry base:

B + H2O BH+ + OH-

Kb = [BH+][OH-]

[B]

Kb provides a measure of the strength of a base1.if Kb is large, the base is largely

dissociated so the base is strong 2.if Kb is small, very little of the

base is dissociated so the base is weak.

H2O(l) + H2O(l) H3O+ (aq) + OH- (aq)

Kw = [H3O+] [OH-] = 1 x 10-14

Measurements at 25˚C show that the[H3O+] and [OH-] are each 1 x 10-7 M

Self-ionization of water

When something is acidic [H3O+] > [OH-];

[H3O+] >1 x 10-7 M and [OH-] < 1 x 10-7 M

When something is basic [H3O+] < [OH-];

[H3O+] <1 x 10-7 M and [OH-] > 1 x 10-7 M

[H3O+] and [OH-] are inversely proportional

pH means proportion of H+ ion

scale from 0 – 14pH = -log[H3O+] pOH = -log[OH-]

As Kw = 1 x 10-14 we can conclude that

pH + pOH = 14

pH [H3O+] pOH [OH-]

0 1 14 1 x 10-14 M

1 1 x 10-1 M 13 1 x 10-13 M

2 1 x 10-2 M 12 1 x 10-12 M

3 1 x 10-3 M 11 1 x 10-11 M

4 1 x 10-4 M 10 1 x 10-10 M

5 1 x 10-5 M 9 1 x 10-9 M

6 1 x 10-6 M 8 1 x 10-8 M

7 1 x 10-7 M 7 1 x 10-7 M

8 1 x 10-8 M 6 1 x 10-6 M

9 1 x 10-9 M 5 1 x 10-5 M

10 1 x 10-10 M 4 1 x 10-4 M

11 1 x 10-11 M 3 1 x 10-3 M

12 1 x 10-12 M 2 1 x 10-2 M

13 1 x 10-13 M 1 1 x 10-1 M

14 1 x 10-14 M 0 1

Acidic pH 0-7[H3O+] > [OH-]; [H3O+] > 1 x 10-7 M [OH-] < 1 x 10-

7 M

pH = 7 NEUTRAL[H3O+] = [OH-] 1 x 10-7 M

Basic pH 7-14 [H3O+] < [OH-]; [H3O+] < 1 x 10-7 M [OH-] > 1 x 10-7 M

pH of common items

compounds whose colors are sensitive to pH. Indicators change colors because they are eitherweak acids or weak bases.

In basic solution HIn H+ + In-

In acidic solution

Acid – base indicators

Indicators come in many different colors. The pH range over which an indicator changes color is called its transition interval.

Blue litmus paper turns red under acidic conditions and red litmus paper turns blue under basic (i.e. alkaline) conditions

A pH meter determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution.

DR between strong acids and bases always form water and a salt (an ionic compound formed from the cation of the base and the anion of the acid). The net ionic equation is:

H3O+ (aq) + OH- (aq) 2H2O (l)

Neutralization reactions

Titration is the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration.

The point at which the two solutions used in atitration are present in chemically equivalentamounts is the equivalence point. The point in atitration at which an indicator changes color iscalled the end point.

The solution that contains the precisely known concentration of a solute is known as a standard solution. A primary standard is a highly purified solid compound used to check the concentration of the known solution in a titration.

Burettes for acid-base titrations

Titration curve strong acid + strong base

Titration curve weak acid + strong base

Titration curves

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• solutions composed of a weak acid or base and its salt (ex. CH3COOH and NaCH3COO) and can withstand large changes in pH.

• Buffer capacity is the amount of acid or base a buffer solution can absorb without a significant change in pH. This has very important biological implications.

BUFFERS

Buffer examplesEx. Blood maintains a pH of 7.4. The conjugate acid-base pair that acts as a buffer system is H2CO3 and HCO3

-.

When there is excess acid presentH3O+ + HCO3

- H2O + H2CO3

When there is excess base presentOH- + H2CO3 H2O + HCO3

-