Chapter 20

Acids and Bases

• Describing Acids and Bases1. Properties of Acids and Bases Acids Bases Contains H+ Contains OH-

Turns blue litmus red Turns red litmus blue Taste sour Taste

bitter Can be electrolytes Can be electrolytes Reacts with bases to Reacts with acids to form water form water

2. Quick review of naming acids

Hydrogen ions and acidity1. Hydrogen Ions from water a. When water molecules lose a hydrogen ion it becomes OH-

Anion ending Example Acid name Example

-ide Cl- chloride Hydro- (stem) –ic acid

-Ite SO3-2 (stem)-ous acid

-ate SO4-2 (stem)-ic acid

b. When water molecules gain a hydrogen ion it becomes H3O+

(called the hydronium ion)

2. Self-ionization a. When two water molecules produce ions

b. H2O (l) H+ (aq) + OH- (aq)

c. [H+] = 1.0 x 10 -7 M

d. [OH-] = 1.0 x 10 -7 M

e. When [H+] and [OH-] are equal it is a neutral solution

f. When they are independent (not equal) [H+] increases, [OH-] decreases [H+] decreases, [OH-] increases

3. Ion-product constant a. kw : product of concentration of H+ and OH-

in water

b. Kw = [H+] [OH-] = 1.0 x 10-14 M2

c. Acidic solution: one where [H+] is greater than [OH-]

[H+] > 1.0 x 10-7 M

d. Basic solution: one where [H+] is less than [OH-]

[H+] < 1.0 x 10-7 M

e. Basic solution also known as Alkaline solution

4. The pH concept a. Better expressed using the pH scale

b. pH + pOH = 14 pH = -log[H+] pOH = -log[OH-]

c. In a neutral solution [H+] = 1.0 x 10-7 M pH = -log (1 x 10-7) = -(log 1 + log 10-7) = -(0.0 + (-7.0)) = 7.0

d.

e.

5. Example problems: a. What is the pH of a solution with a

hydrogen-ion concentration of 1.0 x 10-10M?

b. The pH of an unknown solution is 6.00. What is its hydrogen-ion

concentration?

c. What is the pOH of a solution if [OH-] = 4.0 x 10-11 M?

d. What is [H+] of a solution if the pH = 3.70?

6. Measuring pH a. Acid-base Indicator 1. Indicator (In) is an acid or base that

undergoes dissociation in a known pH range

2. Reaction form: HIn (aq) H+ (aq) + In- (aq) Acid form Base form

3. Types: pH color Thymol blue 1.2-3.0 red yellow 8.0-9.5 yellow blue Bromphenol blue 3.0-4.6 yellow blue Bromcresol green 3.7-5.3 yellow blue methyl red 4.2-6.2 red yellow Alizarin 4.5-6.0 yellow red Bromthymol blue 6.0-7.5 yellow blue Phenol red 6.9-8.2 yellow orange Phenolphthalein 8.0-10.0 colorless pink

alizarin yellow R 8.0 – 12.2 yellow red

4. Useful at room temperature (25 °C)

b. pH meter 1. Useful to make rapid, accurate pH

measurements

2. more practical than liquid indicators

Acid-Base Theories1. Arrhenius Acids and Bases a. Acids are hydrogen containing compounds

that ionize to yield H+ in aq solutions

b. Bases are compounds that ionize to yield OH- in aq solutions

Acids c. Monoprotic acids have one hydrogen HCl

d. Diprotic acids : have two hydrogens H2SO4

e. Triprotic acids: have three hydrogens H3PO4

f. Only very polar bonds will dissociate Hδ+--Clδ- H+ (aq) + Cl- (aq)

g. C-H bonds weakly polar will not dissociate

ex. Ethanoic acid (CH3COOH):

Bases h. NaOH (s) Na+ (aq) + OH- (aq) i. Common bases: KOH, NaOH, Ca(OH)2, Mg(OH)2

2. Bronsted-Lowery Acids and Bases a. Acid is a hydrogen-ion donor

b. Base is a hydrogen-ion acceptor

c. Conjugate acid – particle formed when a base gains a hydrogen ion

d. Conjugate base- particle that remains when an acid has donated a hydrogen ion

e. Conjugate acid-base pair: two substances related by the loss or gain of

a single hydrogen bond f. Examples: 1. NH3(aq) + H2O (l) NH4

+ (aq) + OH- (aq)

acceptor donor (base) (acid) (CA) (CB)

2. HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)

(acid) (base) (CA) (CB)

g. Amphoteric: a substance that can act like both an acid and base

Sample problems• 1. Classify the following as Brønsted acids, bases or

both. a) H2O b) OH- c) NH3 d) NH4

+

• 2. What is the conjugate base of the following acids?

a) HClO4 b) NH4+ c) H2O d) HCO3

-

• 3. What is the conjugate acid of the following bases?

a) CN- b) SO42- c) H2O d) HCO3

-

3. Lewis Acids and Bases a. Acid: a substance that accepts a pair

of electrons to form a covalent bond

b. Base: a substance that donate a pair of electrons to form a covalent

bond c. Examples: 1. H+ + acid base 2.

Strengths of Acids and Bases 1. Strong acids and bases a. Strong acids: completely ionize (dissociate) HCl, HNO3, H2SO4, HBr, HI, HClO4

b. Dissociation constant (Ka): the ratio of the concentration of the

dissociated form of an acid to the concentration of the

undissociated form

***See page 600 Table 20.7

3. Equilibrium-constant expression K = [products] [reactants]

** Remember to raise the concentrations to the coefficient number.

4. Ka = [H+][A-] (Gives the ratio of ions

[HA] vs molecules) Weak acid has Ka <1

Leads to small [H+] and pH of 2-7 5. Kb = [BH+][OH-]

[B] Weak bases has Kb < 1

Leads to small [OH-] and pH of 12-7**Do not use water in the [ ]

6. Examples: a. Calculate the [OH-] of a 0.500 M solution of

aqueous ammonia. The Kb is 1.74 x 10-

5.

NH3 + H2O NH4+ + OH-

Kb = [NH4+][OH-]

[NH3]

b. You have 1.00 M acetic acid (HOAc). Calculate the equilibrium concentrations

of HOAc, H+, OAc-, and the pH. Ka = 1.8 x 10 -5

Step 1 Define equilibrium concentrations . [HOAc] [H+] [OAc-]

Initial: 1.00 0 0Change: -x +x +xEquilib: 1.00-x x x

Step 2: Write the Ka expression

HC2H3O + H20 H+ + C2H3O-

(HOAc) (OAc-) Ka = [H+][OAc-]

[HoAc] 1.8 x 10-5 = (x)(x) = x2

(1.00 –x) (1.00 – x)This is a quadratic. Solve using the quadratic formula. OR

you can make an approximation if x is very small. (Rule of thumb: 10-5 or smaller is OK)

1.8 x 10-5 = x2 1.00

x = [H+] = [OAc -] = 4.2 x 10-3M

pH = -log[4.2 x 10-3] = 2.37

c. You have 0.010 M NH3. Calculate the pH if the Kb = 1.8 x 10-5.

NH3 + H2O NH4+ + OH-

[NH3] [NH4+] [OH-]

Initial 0.010 0 0Change -x x xEquilibrium 0.010-x x x

Kb = [NH4+] [ OH-]

[NH3]

1.8 x 10-5 = (x)(x) 0.010 – x

x = 4.2 x 10-4 M

At equilibrium: 0.010 -4.2 x 10-4 = 0.00958≈0.01

Once you find [OH-], you find the pOH

pH + pOH = 14

pH indicators

1. indicator (In) is an acid or base that dissociates in a known pH range

HIn (aq) acid form

OH-

H+

H+ (aq) + In- (aq) base form

2.

• 3. Types of indicators a. Methyl red: dye that turns red in acids 0-4.4 : red 4.5-6.1: orange 6.2-above: yellow

b. Phenolphthalein: colorless in acids, pink in bases

below pH 8.2: colorless above pH 10: pink

c. Bromothymol blue: used for weak acids/bases

below pH of 6.0 = yellow pH of 7.0 = green above pH of 7.6 = blue

d. Universal indicator: used for acids and bases

0-3 3-6 7 8-11 11-14 red orange/ green blue purple yellow

• Problems with indicators 1. Only work at room temperature (will

change colors at different temp)

2. Salts in the solution may change the dissociation process

pH meter: equipment used to measure pH (best pH measurement)