ACIDS and BASES

Properties of ACIDS Sour to taste React with some metals to form

Hydrogen gas Turn Litmus RED Phenolphthalein stays colorless Electrolytes (conduct) Form H+ (H3O+ when attached to

water molecules)

Properties of BASES Bitter to taste Slippery to touch Turn Litmus BLUE Phenolphthalein turns MAGENTA Electrolytes Many form OH- in water

Definitions of Acids and Bases

Arrhenius (traditional) Acids: produce H3O+ in water Bases: produce OH- in water

Examples HCl (g) → H+

(aq) + Cl-(aq) NaOH (s) → OH-

(aq) + Na+(aq)

Most common definition

Definitions of Acids and Bases

Bronsted-Lowry Acids: H+ donor (proton donor) Bases: H+ acceptor (proton acceptor)

Examples HCl → Cl- (donates H+) NH3 → NH4

+ (accepts H+)

Bronsted-Lowry cont. When acids and bases donate or

accept hydrogen ions, conjugate acids and bases are formed.

Conjugate Acid: particle formed when a Base gains a H+

Conjugate Base: particle formed when an Acid donates a H+

Bronsted-Lowry cont. Conjugate acid-base pair: two

substances related by the loss or gain of a single H+

Always paired with an acid and base Examples- Label acid/base and conjugate

acid/base NH3 + H2O → NH4

+ + OH-

HCl + H2O → H3O+ + Cl- Hint: Acid/Base is always Reactant; Conjugate

Acid/Base is always Product

Terms Amphoteric substances can behave

as acids or bases (water) Monoprotic acids donate only one

H+

Ex. HCl, HNO3

Polyprotic acids donate more than one H+

H2SO4, H2CO3

Relative Strengths of Acids and Bases

• The stronger the acid, the weaker the conjugate base.

• H+ is the strongest acid that can exist in equilibrium in aqueous solution.

• OH- is the strongest base that can exist in equilibrium in aqueous solution.

Relative Strengths of Acids and Bases

Conjugate Acids and Bases Any acid or base that is stronger than

H+ or OH- simply reacts stoichiometrically to produce H+ and OH-.

The conjugate base of a strong acid (e.g. Cl-) has negligible acid-base properties.

Similarly, the conjugate acid of a strong base has negligible acid-base properties.

Acid-Base Equilibria In pure water, the following equilibrium is

established:

This is referred to as the autoionization of water.

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

14-3

-3

22

22

-3

100.1]OH][OH[

]OH][OH[]OH[

]OH[

]OH][OH[

w

eq

eq

K

K

K

Types of Acids and Bases Weak acids and bases

Weak electrolytes Partially ionize in water (much less than

100%) Establish equilibria

Weak Acids H3PO4, HC2H3O2, H2CO3

Weak Bases NH3, low [OH-]

pH Pouvoir hydrogene: “Hydrogen Power”

Measure of the acidity of a solution Uses [H3O+] or [H+]

Concentrations usually expressed as powers of 10

pH = -log [H+]

pH scale 0-7 acid, 7 neutral, 7-14 base Can be lower than 0 or higher than 14

pH Scale

pOH “Hydroxide Power”

Measures alkalinity of a solution Uses [OH-] pOH = -log [OH-]

pOH scale 0-7 base, 7 neutral, 7-14 acid Can be lower than 0 or higher than 14

pH + pOH = 14 From the autoionization of water

Calculating Ka from pH• Weak acids are simply equilibrium calculations.• The pH gives the equilibrium concentration of H+.• Using Ka, the concentration of H+ (and hence the pH) can

be calculated.– Write the balanced chemical equation clearly showing

the equilibrium.– Write the equilibrium expression. Find the value for Ka.– Write down the initial and equilibrium concentrations

for everything except pure water. We usually assume that the change in concentration of H+ is x.

– Substitute into the equilibrium constant expression and solve. Remember to turn x into pH if necessary

• Percent ionization is another method to assess acid strength.

• For the reactionHA(aq) + H2O(l) H3O+(aq) + A-(aq)

100]HA[]OH[

ionization %0

3 eqm

• Percent ionization relates the equilibrium H+ concentration, [H+]eq, to the initial HA concentration, [HA]0.

• The higher percent ionization, the stronger the acid.

• Percent ionization of a weak acid decreases as the molarity of the solution increases.

Polyprotic acids• Polyprotic acids have more than one ionizable

proton.• The protons are removed in steps not all at

once:

• It is always easier to remove the first proton in a polyprotic acid than the second.

• Therefore, Ka1 > Ka2 > Ka3 etc.

H2SO3(aq) H+(aq) + HSO3-(aq) Ka1 = 1.7 x 10-2

HSO3-(aq) H+(aq) + SO3

2-(aq) Ka2 = 6.4 x 10-8

Relationship between Ka and Kb

• For a conjugate acid-base pair

• Therefore, the larger the Ka, the smaller the Kb. That is, the stronger the acid, the weaker the conjugate base.

• Taking negative logarithms:

baw KKK

baw pKpKpK

Strong Acids and Bases Strong electrolytes Ionize (separate) 100% in water

Strong Acids: HCl, HNO3, H2SO4, HClO3, HClO4, HI, HBr

Usually only source of H+, so pH can be calculated from the molarity of the acid (unless < 10-6)

Strong Bases: NaOH, KOH, Ca(OH)2 Most Group 1 and 2 metal hydroxides are strong

bases Ionic metal oxides, hydrides, and nitrides

Salt Solutions Most salts are strong electrolytes (ionize

in solution) Acid-Base properties result from their

ions. Many ions react with water to form H+ and

OH- (hydrolysis) Anions from weak acids are basic. Anions from strong acids are neutral. All cations (except alkali/alkaline earth

metals) are weak acids

Salt Solutions SA + SB = Neutral Salt SA + WB = Acidic Salt WA + SB = Basic Salt WA + WB = Acidic or Basic Salt

Based on relative strength of Ka and Kb Ka > Kb = acidic Ka < Kb = basic

Acid-Base Behavior and Chemical Structure

For compound H-X, If H is partially positive, then it is an acid If H is partially negative, then it is a base

Bond Strength and Polarity affects acid/base strength

Bond strength used to determine strength in a group; Bond polarity used to determine strength in a period

Acid strength tends to increase down a group; Base strength tends to decrease down a group

Acid strength tends to increase L to R; Base strength tends to decrease L to R

Acid-Base Behavior and Chemical Structure

Oxyacids Acids that contain OH groups bound to

the central atom ( Y – O – H) Strength depends on Y and the atoms

attached to Y Increasing electronegativity of Y =

increasing acidity Increasing the number of O atoms attached

to Y increases polarity, which increases strength

Ex. HClO < HClO2 < HClO3 < HClO4

Carboxylic Acids Contain –COOH Additional oxygen atom on the

carboxyl group increase the polarity of the O-H bond and stabilizes the conjugate base

Lewis Acids and Bases Lewis Acid – electron pair acceptor Lewis Base – electron pair donor

Do not need to contain protons – most general definition of acids/bases

Many Lewis acids have an incomplete octet (BF3) Transition metal ions are usually Lewis acids Lewis acids must have an empty orbital Compounds with multiple bonds can be Lewis

acids

Amino Acids Amphoteric

Contains carboxyl group (acid) and ammine group (basic)

Proton of the carboxyl group is transferred to the basic nitrogen of the ammine

Results in a zwitterion or dipolar ion