Unit 13Unit 13

Acids and BasesAcids and Bases

D. Finding the pH of SolutionsSelf- ionization of water – the simple dissociation of water

H2O H+ + OH-

Concentration of each ion in pure water: [H+] = 1.0 x 10-7M + [OH-] = 1.0 x 10-7M

Ion-product constant for water (Kw), Where KWhere Kww = 1.0 x 10 = 1.0 x 10-14-14

KKww = [H = [H++] [OH] [OH--]]

Acid [H+] > [OH-] Base [H+] < [OH-]

Neutral [H+] = [OH-]

[OH-] pOH pH [H+]

1 x 10-14 14 0 1 x 100

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

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

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

1 x 10-10 10 Increasing acidity 4 1 x 10-4

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

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

1 x 10-7 7 Neutral 7 1 x 10-7

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

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

1 x 10-4 4 Increasing basicity 10 1 x 10-10

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

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

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

1 x 100 0 14 1 x 10-14

Calculating [H+] and [OH-]

• reverse the pH equation

• The pH of a solution is 7.52. Find the [H+] and [OH-] and determine whether it is acidic, basic, or neutral.

– basic

pOHpH 10][OH and 10][H

M 101010][OH

M1010][H6-48.6)52.714(

-852.7

Example1. If the [H+] in a solution is 1.0 x 10-5M, is the

solution acidic, basic or neutral?

1.0 x 10-5 M

What is the concentration of the [OH-]?Use the ion-product constant for water (Kw):

Kw = [H+] [OH-] 1.0 x 10-14 = [1.0 x 10-5] [OH-] 1.0 x 10-14 = [OH-] 1.0 x 10-5

1.0 x 10-(14-5)

pH 5 = acidic

1.0 x 10-9 OH-

Examples2. If the pH is 9, what is the concentration of

the hydroxide ion?

Kw = [H+] [OH-]

1.0 x 10-14 = [1.0 x 10-9] [OH-]

1.0 x 10-5 = [OH-]

3. If the pOH is 4, what is the concentration of the hydrogen ion?

Kw = [H+] [OH-]

1.0 x 10-14 = [H+] [1.0 x 10-4]

1.0 x 10-10 = [H+]

14 = pH + pOH

14 = 9 + pOH

5 = pOH

14 = pH + pOH

14 = pH + 4

10 = pH

Example

• A solution has a pH of 4. Calculate the pOH, [H+] and [OH-]. Is it acidic, basic, or neutral?

14 = pH + pOH

14 = 4 + pOH

10 = pOH

– acidic

M101][OH 10

M4101][H

Practice Problems:

Classify each solution as acidic, basic or neutral.

1. [H+] = 1.0 x 10-10

2. [H+] = 0.001

3. [OH-] = 1.0 x 10-7

4. [OH-] = 1.0 x 10-4

Basic pH 101.0 x 10-3 acid pH 3

Neutral14 – 4 = 10 base pH 10

Fill in the chart.

[OH-] pOH pH [H+]

8

1x 10-12

10

1 x 10-3

5

1 × 10-1

1.0 X 10 -8

1.0 X 10 -2

1.0 X 10 -4

1.0 X 10 -6

1.0 X 10 -10

1.0 X 10 -11

6

3 11

4

2 12

9

113

1.0 X 10 -5 1.0 X 10 -9

1.0 X 10 -13

pH = -log[H+]

E. pH Scale

0

7INCREASING

ACIDITY NEUTRALINCREASING

BASICITY

14

pouvoir hydrogène (Fr.)“hydrogen power”

pH is the negative logarithm of the hydrogen ion concentration

E. pH Scale

pH = -log[H+]

pOH = -log[OH-]

pH + pOH = 14

E. The pH Scale

E. pH Scale

pH of Common SubstancespH of Common SubstancespH of Common SubstancespH of Common Substances

F. Neutralization

• Chemical reaction between an acid and a base.

• Products are a salt (ionic compound) and water.

F. Neutralization

ACID + BASE ACID + BASE SALT + WATER SALT + WATER

HCl + NaOH HCl + NaOH NaCl + H NaCl + H22OO

HCHC22HH33OO22 + NaOH + NaOH NaC NaC22HH33OO22 + H + H22OO

– Salts can be neutral, acidic, or basic.

– Neutralization does not mean pH = 7.

weak

strong strong

strong

neutral

basic

G. Titration

• TitrationTitration– Analytical method in

which a standard solution is used to determine the concentration of an unknown solution.

standard solution

unknown solution

• Equivalence point Equivalence point (endpoint)(endpoint)– Point at which equal amounts of H+

and OH- have been added.– Determined by…

• indicator color change

G. Titration

• dramatic change in pH

G. Titration

moles H+ = moles OH-

MV n = MV n

M: MolarityV: volumen: # of H+ ions in the acid

or OH- ions in the base

G. Titration

• 42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4.

H3O+

M = ?V = 50.0 mLn = 2

OH-

M = 1.3MV = 42.5 mLn = 1

MV# = MV#M(50.0mL)(2)=(1.3M)(42.5mL)(1)

M = 0.55M H2SO4

Naming Acids

• Binary acids

– Contains 2 different elements: H and another

– Always has “hydro-” prefix

– Root of other element’s name

– Ending “-ic”

– Examples: HI, H2S, HBr, HCl

Naming Acids

• Ternary Acids - Oxyacids

– Contains 3 different elements: H, O, and another

– No prefix

– Name of polyatomic ion

– Ending “–ic” for “-ate” and “–ous” for “-ite”

– Examples: HClO33, H3PO4, HNO2

Practice

• H2SO3

– Sulfurous acid• HF

– Hydrofluoric acid

• H2Se

– Hydroselenic acid

• Perchloric acid

– HClO4

• Carbonic acid

– H2CO3

• Hydrobromic acid– HBr

Definitions of Acids and Bases

• Arrhenius – Most specific/exclusive definition– Created by Svante Arrhenius, Swedish– Acid : compound that creates H+ in an

aqueous solution– Base : compound that creates OH- in an

aqueous solution

– HNO3 H+ + NO3-

– NaOH Na+ + OH-

Definitions of Acids and Bases• Bronsted-Lowry

– A bit more general than Arrhenius definition– Most commonly used definition– Created by two scientists around the same time

(1923)– Acid: Molecule or ion that is a proton (H+) donor– Base: Molecule or ion that is a proton (H+) acceptor

– HCl + H2O H3O+ + Cl-

– NH3 + H2O ↔ NH4+ + OH-

Definitions of Acids and Bases• Lewis

– The most general definition– Defined by electrons and bonding instead of H+

– Created by same scientist who electron-dot diagrams are named after

– Acid: atom, ion, or molecule that accepts electron pair to form covalent bond

– Base: atom, ion or molecule that donates and electron pair to form covalent bond

– NH3 + Ag+ [Ag(NH3)2]1+

– BF3 + F- BF4-