Acids and Bases
12.1 Properties of Acids and Bases
12.2 Definitions of Acids and Bases
12.3 Water as an Acid; Water as a Base
12.4 Strong Acids and Bases
12.5 pH and pOH Scales
12.6 Weak Acids and Bases
12.7 Acid–Base Titrations
12.8 Buffers
12.1 Arrhenius Acids and Bases
- Acids: Produce H^+ when dissolved in water.
- Examples: HF, HCl, HBr, H2SO4, H3PO4, H2CO3, HClO_4, and more.
- Bases: Produce OH^- when dissolved in water.
- Examples: NaOH, KOH, Ca(OH)2, Ba(OH)2, and more.
- Neutralization Reaction: Reaction between an acid and a base.
12.2 Brønsted Acids and Bases
- Brønsted Acid: Proton donors.
- Brønsted Base: Proton acceptors.
- Proton: H^+ = hydrogen atom without its electron.
- Example:
- Ammonia gas plus hydrogen chloride gas:
- HCl donates H^+ to NH_3
12.2 Definitions of Acids and Bases—Conjugate Acid-Base Pairs
- Brønsted Acid: Donates a proton and becomes a conjugate base.
- Brønsted Base: Accepts a proton and becomes a conjugate acid.
- Example:
Conjugate Bases of Common Species
- Acetic Acid → Acetate (Took away H^+)
- Nitrous Acid → Nitrite (Took away H^+)
- Sulfuric Acid → Hydrogen sulfate ion (Took away H^+)
Conjugate Acids of Common Species
- Ammonia → Ammonium cation (Added H^+)
- Water → Hydronium cation (Added H^+)
- Hydroxide → Water (Added H^+)
- Urea → Uronium cation (Added H^+)
Sample Problems 12.1
- What is the conjugate acid of O^{2-}? OH^-
- What is the conjugate base of HSO4^−? SO4^{2-}
- What is the conjugate acid of HCO3^−? H2CO_3
Sample Problem 12.2
- Label each species as an acid, base, conjugate base, or conjugate acid:
- (a) HF (aq) + NH3 (aq) ←→ F^− (aq) + NH4^+ (aq)
- HF loses a proton and becomes F^−: HF = acid, F^− = conjugate base
- NH3 gains a proton and becomes NH4^+: NH3 = base, NH4^+ = conjugate acid
- (b) CH3COO^− (aq) + H2O (l) ←→ CH_3COOH (aq) + OH^− (aq)
- CH3COO^− gains a proton and becomes CH3COOH: CH3COO^− = conjugate base, CH3COOH = acid
- H2O loses a proton to become OH^−: H2O = acid, OH^−=conjugate base
- (a) HF (aq) + NH3 (aq) ←→ F^− (aq) + NH4^+ (aq)
12.3 Water as an Acid; Water as a Base—Amphoteric
- Amphoteric: A species that can behave either as a Brønsted acid or a Brønsted base.
- Autoionization of water:
- Pure water has a low concentration of OH^− and H_3O^+.
12.3 Water Autoionization
- PURE WATER: H_3O^+ and OH^− are ALWAYS present
- If [H_3O^+]↑, then [OH^−]↓
- If [H_3O^+]↓, then [OH^−]↑
Sample Problem 12.3
- The concentration of hydronium ions in stomach acid is 0.10 M. Calculate the concentration of hydroxide ions in stomach acid at 25°C.
Sample Problem 12.3 - Solution
- Calculate [OH^−] when [H3O^+] = 0.10 M Kw = [H3O^+][OH^−] = 1.0 × 10^{-14} at 25°C [OH^−] = \frac{1.0 × 10^{-14}}{[H3O^+]} = \frac{1.0 × 10^{-14}}{0.10} = 1.0 × 10^{-13} M
12.4 Strong Acids and Bases
- Strong acids ionize completely in aqueous solution.
- All HCl molecules come apart: H^+ and Cl^−
- Strong bases dissociate completely in aqueous solution
- All NaOH dissociates to Na^+ and OH^−
Sample Problem 12.4
- Calculate the [OH^−] when [H_3O^+] of an aqueous solution that is:
- (a) 0.0311 M HNO_3
- (b) 4.51 x 10^{-5} M HClO_4
- (c) 8.74 x 10^{-6} M HI
Sample Problem 12.4 - Solution
- HNO3, HClO4, and HI are strong acids
- [H_3O^+] = concentration of the strong acid
- (a) 0.0311 M HNO3: [H3O^+] = 0.0311 M, [OH^−] = 3.22 x 10^{-13} M
- (b) 4.51 x 10^{-5} M HClO4: [H3O^+] = 4.51 x 10^{-5} M, [OH^−] = 2.22 x 10^{-10} M
- (c) 8.74 x 10^{-6} M HI: [H_3O^+] = 8.74 x 10^{-6} M, [OH^−] = 1.14 x 10^{-9} M
- Solve for [OH^−]
Sample Problem 12.5
- Calculate the [H_3O^+] of an aqueous solution that is
- (a) 0.0311 M LiOH
- (b) 4.51 x 10^{-5} M Ca(OH)_2
- (c) 8.74 x 10^{-6} M KOH
Sample Problem 12.5: Solution - OH^− conc.
- LiOH, KOH, and Ca(OH)_2 are strong bases
- Ca(OH)2: two moles of OH^− per mole of Ca(OH)2
- (a) 0.0311 M LiOH: [OH^−] = 0.0311 M
- (b) 4.51 x 10^{-5} M Ca(OH)_2: [OH^−] = 2 x (4.15 x 10^{-5} M) = 8.30 x 10^{-5} M
- (c) 8.74 x 10^{-6} M KOH: [OH^−] = 8.74 x 10^{-6} M
Sample Problem 12.5 - Solution - [H_3O^+]
- (a) 0.0311 M LiOH: [OH^−] = 0.0311 M
- (b) 4.51 x 10^{-5} M Ca(OH)_2: [OH^−] = 2 x (4.15 x 10^{-5} M) = 8.30 x 10^{-5} M
- (c) 8.74 x 10^{-6} M KOH: [OH^−] = 8.74 x 10^{-6} M
12.5 pH and pOH—Overview
- Acidity depends on [H_3O^+]
- pH scale.
- The pH of a solution is defined as the negative base-10 logarithm of the hydronium ion concentration in mol/L.
- pH = -log[H_3O^+]
- The pOH is defined as the negative base-10 logarithm of the hydroxide concentration in mol/L.
- pOH = -log[OH^-]
Benchmark pH Values
- Range of Hydronium Ion Concentrations at 25°C
- [H_3O^+] (M) pH
- 0.10 1.00 (Acidic)
- 1.0 × 10^{-7} 7.00 (Neutral)
- 1.0 × 10^{-14} 14.00 (Basic)
pH Values of Some Common Fluids
- Fluid pH
- Stomach acid 1.5
- Lemon juice 2.0
- Vinegar 3.0
- Grapefruit juice 3.2
- Orange juice 3.5
- Urine 4.8-7.5
- Rainwater 5.5
- Saliva 6.4-6.9
- Milk 6.5
- Pure water 7.0
- Blood 7.35-7.45
- Tears 7.4
- Milk of magnesia 10.6
- Household ammonia 11.5
Sample Problem 12.6
- Determine the pH of a solution at 25°C in which the hydronium ion concentration is
- (a) 3.5 x 10^{-4} M [H_3O^+]
- (b) 1.7 x 10^{-7} M [H_3O^+]
- (c) 8.8 x 10^{-11} M [H_3O^+]
Sample Problem 12.6 - Solution
- (a) 3.5 x 10^{-4} M [H_3O^+] pH = 3.46
- (b) 1.7 x 10^{-7} M [H_3O^+] pH = 6.77
- (c) 8.8 x 10^{-11} M [H_3O^+] pH = 10.06
Using pH to Calculate Hydronium Ion Concentration
- Equation 12.3: [H_3O^+] = 10^{-pH}
Sample Problem 12.7
- Calculate the hydronium ion concentration in a solution at 25°C in which the pH is:
- (a) 4.76
- (b) 11.95
- (c) 8.01
Sample Problem 12.7 - Solution
- Calculate the hydronium ion concentration in a solution at 25°C in which the pH is:
- (a) 4.76: [H_3O^+] = 10^{-pH} = 10^{-4.76} = 1.7 × 10^{-5}
- (b) 11.95: [H_3O^+] = 10^{-pH} = 10^{-11.95} = 1.1 × 10^{-12}
- (c) 8.01: [H_3O^+] = 10^{-pH} = 10^{-8.01} = 9.8 × 10^{-9}
12.5 pH and pOH—REPEAT
Acidity depends on [H_3O^+]
pH scale
The pH of a solution is defined as the negative base-10 logarithm of the hydronium ion concentration in mol/L
The pOH is defined as the negative base-10 logarithm of the hydroxide concentration in mol/L
Sample Problem 12.8
Determine the pOH of a solution at 25°C in which the hydroxide ion concentration is
(a) 3.7 x 10-5 M H3O^+
(b) 4.1 x 10-7 M H3O^+
(c) 8.3 x 10-2 M H_3O^+
Sample Problem 12.8—Problem
Determine the p O H of a solution at 25°C in which the hydroxide ion concentration is
(a) 3.7 x 10^{-5} M OH^−: pOH = -log(3.7 x 10^{-5}) = 4.43
(b) 4.1 x 10^{-7} M OH^−: pOH = -log(4.1 x 10^{-7}) = 6.39
(c) 8.3 x 10^{-2} M OH^−: pOH = -log(8.3 x 10^{-2}) = 1.08
Using pOH to Calculate Hydroxide Ion Concentration
POH = -log [OH-]
[OH-] = 10-POH
[OH] = 10-POH
Sample Problem 12.9
Calculate the hydroxide ion concentration in a solution at 25°C in which the p OH is
(a) 4.91 [OH-] = 10-pOH
(b) 9.03 [OH-] = 10-pOH
(c) 10.55 [OH-] = 10-pOH
Sample Problem 12.9—Problem
Calculate the hydroxide ion concentration in a solution at 25°C in which the p OH is
(a) 4.91 [OH^−] = 10^{-pOH} = 10^{-4.91} = 1.2 × 10^{-5} M
(b) 9.03 [OH^−] = 10^{-pOH} = 10^{-9.03} = 9.3 × 10^{-10} M
(c) 10.55 [OH^−] = 10^{-pOH} = 10^{-10.55} = 2.8 × 10^{-11} M
pH + pOH = 14 At 25 °C (room temperature)
-log(Kw) = -log([H3O+][OH-]) = -log(1.0 x 10-14)
K w = [H3O+][OH-] = 1.0 x 10-14
p(Kw) = p([H3O+][OH-]) = p(1.0 x 10-14)
p(Kw) = p[H3O+] + p[OH-]) = p(1.0 x 10-14)
12.6 Weak Acids and Bases—Weak Acids
Weak acids ionize PARTIALLY!
We use the double arrow to indicate partial reaction
12.6 Common Weak Acids
Name of Acid Formula
Hydrofluoric acid HF
Nitrous acid HNO2
Formic acid HCOOH
Benzoic acid C6H5COOH
Acetic acid CH3COOH
There are too many weak acids to memorize, so we memorize the short list of strong acids. Other acids are weak.
HCl HBr HI HN O3 HCl O3 HCl O4 H2 S O4
Sample Problem 12.10
Determine the pH value that a 0.10-M solution of HC2H3O2 must be above
This question is asking you to think about what the pH would be if ALL THE ACETIC ACID dissociated into H+ and the acetate anion. In other words, pretent the acid is STRONG. What would the pH be if this were a strong acid?
[H3O+] = concentration of the strong acid pH = -log([H3O+])
Sample Problem 12.10
Determine the pH value that a 0.10-M solution of HC2H3O2 must be above
This question is asking you to think about what the pH would be if ALL THE ACETIC ACID dissociated into H+ and the acetate anion. In other words, pretent the acid is STRONG. What would the pH be if this were a strong acid?
[H3O+] = concentration of the strong acid pH = -log([H3O+])
pH = -log([H3O+]) = -log(0.1) = 1.00 Acetic acid is WEAK, so we know the pH must be higher than 1.00
12.6 Weak Acids and Bases—Weak Bases
Weak bases ionize PARTIALLY
NH3 increases [OH-] by accepting a protons from water
12.6 Weak Acids and Bases—Common Weak Bases
Name of Base Formula
Ethylamine C2H5NH2
Methylamine CH3NH2
Ammonia NH3
Pyridine C5H5N
Aniline C6H5NH2
Dimethyl amine C2H7N
Too many weak bases to memorize, Memorize the short list of strong bases. Any other base is weak. NaOH KOH RbOH CsOH LiOH Ca(OH)2 Sr(OH)2 Ba(OH) 2*
Sample Problem 12.11-Problem
Determine the pH value that a 0.089-M solution of NH 3 must be below.
Sample Problem 12.11
Determine the pH value that a 0.089-M solution of NH3 must be below
Ammonia, NH3, is a weak base and does not ionize completely. pOHMAX = -log(0.089) = 12.95 This is the value if NH3 were STRONG It’s not! so pOH will be LESS than 12.95.
12.7 Acid-Base Titrations—Definitions
Acid–base neutralization: TITRATION
base with known concentration added slowly to acid with unknown concentration
RECORD volume and concentration of base required to neutralize acid sample
Calculate amount of acid neutralized
Calculate original acid concentration
12.7 Acid-Base Titrations—Before the Titration
base of known concentration
acid of unknown concentration
Record volume base added
12.7 Acid-Base Titrations—Equivalence Point
equivalence point → [H3O+] = [OH-]
HOW DO YOU KNOW? INDICATOR
phenolphthalein indicator turns pink
Sample Problem 12.12—HCl
(a) If 25.0 mL of an HCl solution requires 46.3 mL of a 0.203-M NaOH solution to neutralize, what is the concentration of the HCl solution.
Sample Problem 12.12—HCl
(a) If 25.0 mL of an HCl solution requires 46.3 mL of a 0.203-M NaOH solution to neutralize, what is the concentration of the HCl solution.
Molarity of original
Sample Problem 12.12—H2SO4
(b) If 25.0 mL of an H2S O4 solution requires 46.3 mL of a 0.203-M NaOH solution to neutralize, what is the concentration of the H2SO4 solution.
Sample Problem 12.12—H2SO4
(b) If 25.0 mL of an H2S O4 solution requires 46.3 mL of a 0.203-M NaOH solution to neutralize, what is the concentration of the H2SO4 solution.
H2SO4 is a diprotic acid
Molarity of original
12.8 Buffers—Buffers
Buffer prevents drastic pH change upon the addition of acids or bases.
BUFFER = weak acid and its conjugate base [weak acid] ~ [conjugate base] to be effective OR
BUFFER = weak base and its conjugate acid [weak base] ~ [conjugate acid] to be effective
12.8 Non-Buffered Water
Add strong acid -- BIG pH change
pH 7.00 -- before adding one drop of strong acid
pH 3.30 -- after adding one drop of 1 M strong acid
12.8 Buffer Solution
Adda drop of strong acid → smaller pH change
pH 4.74
pH 4.72
Sample Problem 12.13
Indicate which of the following pairs of substances can be used to prepare a buffer:
(a) HCl and NaCl
(b) NaF and KF
(c) HCN and NaCN
Sample Problem 12.13
Indicate which of the following pairs of substances can be used to prepare a buffer:
(a) HCl and NaCl -- strong acid plus Cl- (conjugate base of HCl) strong acid plus conjugate base = TERRIBLE BUFFER Answer is NO
(b) NaF and KF -- no acid present Only have F- , which is a conjugate base, but there’s no acid. Answer is NO
(c) HCN and NaCN -- weak acid and its conjugate base YES -- this one can buffer!!!
Learning Outcomes for Chapter 12
1. Recognize Arrhenius acids and bases
2. Define Arrhenius acid, base and Bronsted acid, base
3. Define hydronium ion
4. Define conjugate acid-base pair
5. Identify a conjugate acid-base pair
6. Define the term amphoteric
7. Write a chemical equation to represent the autoionization of water
8. State the product of hydronium ion concentration and hydroxide ion concentration in water at 25°C
9. Given either hydronium or hydroxide ion concentration, determine the other
10. Identify the strong acids
11. Identify the strong bases
12. Describe the properties of a strong acid using a chemical equation
13. Calculate [OH–] or [H3O+] given the concentration of either a strong acid or strong base
14. State the pH equation – the relationship between pH and [H3O+]
15. Calculate pH given the concentration of either a strong acid or strong base
16. Determine [H3O+] from pH
17. Describe the difference between a strong acid and weak acid
18. Describe the difference between a strong base and weak base
19. Illustrate the behavior of a strong acid in water
20. Illustrate the behavior of a weak acid in water
21. Illustrate the behavior of a strong base in water
22. Illustrate the behavior of a weak base in water
23. Estimate the pH value of a weak acid or weak base solution
24. Define the terms titration and equivalence point
25. Calculate the concentration of an unknown acid or base, given titration data
26. Describe a buffer and how it works
27. Recognize solutions that exhibit buffer properties