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Unit 11

Acids & Bases

Definitions · pH/pOH · Strong & Weak · Ka/Kb · Titration

Concept Overview

11.1–11.3 Properties of Acids and Bases

Acids: taste sour, react with metals to produce H₂ gas, turn blue litmus red, pH < 7, contain H⁺ in solution. Examples: vinegar (acetic acid), lemon juice (citric acid), stomach acid (HCl).

Bases: taste bitter, feel slippery, turn red litmus blue, pH > 7, contain OH⁻ in solution. Examples: baking soda (NaHCO₃), ammonia (NH₃), soap.

11.4 Strength vs. Concentration

Strong/Weak refers to the degree of dissociation (how much the acid/base breaks apart). Concentrated/Dilute refers to the amount of solute dissolved. A weak acid can be concentrated; a strong acid can be dilute — these are independent properties.

  • Strong acid: essentially 100% dissociation (single arrow →)
  • Weak acid: partial dissociation, equilibrium mixture (double arrow ⇌)

11.5–11.7 Arrhenius vs. Brønsted-Lowry Definitions

DefinitionAcidBase
ArrheniusProduces H⁺ in waterProduces OH⁻ in water
Brønsted-LowryProton (H⁺) donorProton (H⁺) acceptor

Note: Water can act as either acid or base (amphoteric). Brønsted-Lowry is broader — it includes reactions not in water and includes species like NH₃ that have no OH⁻ but clearly act as bases.

11.8 Conjugate Acid/Base Pairs

In a Brønsted-Lowry reaction, when an acid donates H⁺ it becomes its conjugate base; when a base accepts H⁺ it becomes its conjugate acid. Conjugate pairs differ by exactly one H⁺.

HCl + H₂O → H₃O⁺ + Cl⁻ acid base conj.acid conj.base NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ base acid conj.acid conj.base

11.10–11.11 Strong Acids and Bases — MEMORIZE

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

Strong Bases (hydroxides of): alkali metals (Li, Na, K, Rb, Cs) and Ba, Sr, Ca. Everything else is weak.

Common Weak Acids: HC₂H₃O₂ (acetic acid), H₃PO₄ (phosphoric acid), HF, HCN, H₂CO₃

11.12–11.13 Self-Ionization of Water and Kw

H₂O + H₂O ⇌ H₃O⁺ + OH⁻ Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ (at 25°C)

Water always has both H⁺ and OH⁻ present. If you know one, calculate the other:

[OH⁻] = Kw / [H⁺] [H⁺] = Kw / [OH⁻]

11.14–11.16 pH and pOH Calculations

pH = −log[H⁺] [H⁺] = 10^(−pH) pOH = −log[OH⁻] [OH⁻] = 10^(−pOH) pH + pOH = 14
Solution[H⁺] vs [OH⁻]pH
Acidic[H⁺] > [OH⁻]pH < 7
Neutral[H⁺] = [OH⁻]pH = 7
Basic[H⁺] < [OH⁻]pH > 7

11.19 Strong Acid/Base Ion Concentrations

For strong acids/bases, dissociation is complete:

0.025 M HCl → [H⁺] = 0.025 M → pH = −log(0.025) = 1.60 0.010 M NaOH → [OH⁻] = 0.010 M → pOH = 2.00 → pH = 12.00

H₂SO₄ gives 2× [H⁺] since it has 2 ionizable protons.

11.20–11.22 Ka, Kb, and ICE Charts for Weak Acids/Bases

Ka = acid dissociation constant (larger Ka → stronger acid)

HA ⇌ H⁺ + A⁻ Ka = [H⁺][A⁻] / [HA] ICE Chart: HA H⁺ A⁻ I: 0.10 M 0 0 C: −x +x +x E: 0.10−x x x Ka = x² / (0.10 − x) ≈ x² / 0.10 (if x is small)

% Dissociation = (x / initial concentration) × 100%

Larger Ka or Kb → more dissociation → stronger acid/base

11.24–11.28 Neutralization and Titration

Neutralization: acid + base → salt + water

HCl + NaOH → NaCl + H₂O H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O

Titration equation (complete):

N_a × C_a × V_a = N_b × C_b × V_b (or for monoprotic: M_a × V_a = M_b × V_b)
  • Equivalence point: moles acid = moles base (stoichiometric amount added)
  • Endpoint: indicator changes color (ideally matches equivalence point)
  • Incomplete titration: limiting reactant used up → find excess [H⁺] or [OH⁻] → calculate pH

Common errors: overshooting endpoint (too much base added) → calculated concentration of unknown is too high; rinsing buret with water → dilutes solution → results in lower-than-actual concentration.

11.17–11.18 Particle Diagrams: Strong vs. Weak

A strong acid particle diagram shows mostly ions (H⁺ and A⁻), very few or no intact HA molecules. A weak acid particle diagram shows mostly intact HA molecules with only a few H⁺ and A⁻ ions present.

Two acids at the same initial concentration: the strong acid has a lower pH (higher [H⁺]) because it dissociates completely.

Flashcards

Click any card to flip and reveal the answer.

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Arrhenius Acid
A substance that produces H⁺ ions (protons) when dissolved in water. Example: HCl → H⁺ + Cl⁻
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Arrhenius Base
A substance that produces OH⁻ ions when dissolved in water. Example: NaOH → Na⁺ + OH⁻
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Brønsted-Lowry Acid
A proton (H⁺) donor. Can operate in non-aqueous environments. Broader definition than Arrhenius.
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Brønsted-Lowry Base
A proton (H⁺) acceptor. Example: NH₃ is a base because it accepts H⁺ from water to form NH₄⁺.
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Conjugate Acid/Base Pair
Two species that differ by exactly one H⁺. The acid donates H⁺ to become its conjugate base; the base accepts H⁺ to become its conjugate acid.
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6 Strong Acids
HCl, H₂SO₄, HClO₄, HNO₃, HI, HBr. These dissociate 100% in water. All others are weak.
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Strong vs. Weak
Strong = 100% dissociation (single arrow). Weak = partial dissociation, equilibrium (double arrow ⇌). Independent of concentration (concentrated/dilute).
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Kw
Water dissociation constant = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C. Applies to ALL aqueous solutions at 25°C.
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pH Formula
pH = −log[H⁺] and [H⁺] = 10^(−pH) Acidic: pH < 7 | Neutral: pH = 7 | Basic: pH > 7
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pH + pOH Relationship
pH + pOH = 14 (at 25°C). If you know one, subtract from 14 to get the other.
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Ka Expression
For HA ⇌ H⁺ + A⁻ : Ka = [H⁺][A⁻] / [HA] Larger Ka = stronger acid (more dissociated at equilibrium).
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ICE Chart (Weak Acid)
I = Initial, C = Change (±x), E = Equilibrium. Plug E row into Ka expression, solve for x = [H⁺], then find pH.
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% Dissociation
% dissociation = ([H⁺]eq / [HA]initial) × 100% Higher % = stronger acid. Strong acids are ~100%.
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Neutralization Reaction
Acid + Base → Salt + Water Example: HCl + NaOH → NaCl + H₂O Always balance the equation first!
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Equivalence Point
The point in a titration where moles of acid exactly equal moles of base (stoichiometrically). NOT necessarily pH = 7 unless strong acid + strong base.
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Endpoint (Titration)
The point where the indicator permanently changes color. Ideally matches the equivalence point. Determined visually by the experimenter.
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Titration Formula
M_a × V_a = M_b × V_b (monoprotic acids/bases) or N_a × C_a × V_a = N_b × C_b × V_b (using normality)
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Incomplete Titration
Unequal moles acid and base mixed. The limiting reactant is consumed; the excess determines pH. Find moles excess, divide by total volume → [H⁺] or [OH⁻] → pH.
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Amphoteric / Amphiprotic
A substance that can act as either an acid or a base. Water is the classic example: it donates H⁺ (acid) or accepts H⁺ (base).
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Self-Ionization of Water
H₂O + H₂O ⇌ H₃O⁺ + OH⁻ This occurs because water molecules can transfer protons to each other. Kw = 1.0×10⁻¹⁴ at 25°C.

Quick Review

Key Equations

  • Kw = [H⁺][OH⁻] = 1.0×10⁻¹⁴
  • pH = −log[H⁺]
  • pOH = −log[OH⁻]
  • pH + pOH = 14
  • [H⁺] = 10^(−pH)
  • Ka = [H⁺][A⁻] / [HA]
  • % dissoc = (x / C₀) × 100%
  • M_a·V_a = M_b·V_b (titration)

6 Strong Acids (Memorize!)

  • HCl — hydrochloric acid
  • H₂SO₄ — sulfuric acid
  • HClO₄ — perchloric acid
  • HNO₃ — nitric acid
  • HI — hydroiodic acid
  • HBr — hydrobromic acid

Definitions Quick Hit

  • Arrhenius acid → makes H⁺
  • Arrhenius base → makes OH⁻
  • B-L acid → H⁺ donor
  • B-L base → H⁺ acceptor
  • Conjugate pair → differ by 1 H⁺
  • Amphoteric → acid or base
  • Ka large → stronger acid

pH Reference

  • pH < 7 → acidic
  • pH = 7 → neutral (25°C)
  • pH > 7 → basic
  • Each pH unit = 10× change in [H⁺]
  • Strong base: [OH⁻] = M of base
  • Strong acid: [H⁺] = M of acid
  • pOH = 14 − pH

Titration Tips

  • Equivalence point: mol acid = mol base
  • Endpoint: indicator color change
  • Overshoot → [base] calculated too high
  • Rinsing with water → dilutes → lower M
  • Incomplete titration → find excess moles
  • Total volume = V_acid + V_base

ICE Chart Steps

  • Write equilibrium reaction
  • Fill in Initial concentrations
  • Change: −x for reactants, +x for products
  • Equilibrium = I + C
  • Substitute into Ka = x²/(C₀−x)
  • Solve for x = [H⁺]
  • pH = −log(x)

Practice Quiz

Select an answer for each question, then click Check Answers to see your score.

1 According to the Arrhenius definition, which of the following describes an acid?

2 In the reaction NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, what is the role of NH₃ according to the Brønsted-Lowry definition?

3 In the reaction HF + H₂O ⇌ H₃O⁺ + F⁻, what is the conjugate base of HF?

4 Which of the following is NOT one of the 6 strong acids?

5 A student has a 0.1 M solution of acetic acid (a weak acid) and a 0.1 M solution of hydrochloric acid. How do their pH values compare?

6 What is the Kw for the self-ionization of water at 25°C?

7 If [H⁺] = 1.0 × 10⁻³ M, what is the pH of the solution?

8 A solution has a pOH of 4.0. What is its pH?

9 A solution has [OH⁻] = 1.0 × 10⁻⁵ M. What is the [H⁺]?

10 What is the [H⁺] in a 0.050 M solution of NaOH (a strong base)?

11 Which acid has the LARGEST Ka value and is therefore the strongest?

12 For the weak acid HA with Ka = 1.0 × 10⁻⁵, the equilibrium expression is:

13 A 0.10 M weak acid HA has Ka = 1.0 × 10⁻⁵. Using the approximation x² ≈ Ka × C₀, what is the [H⁺]?

14 What is the percent dissociation of the 0.10 M weak acid from question 13 ([H⁺] = 1.0 × 10⁻³ M)?

15 Which of the following is a balanced neutralization reaction?

16 A 25.0 mL sample of HCl is titrated with 0.200 M NaOH. It takes 30.0 mL of NaOH to reach the equivalence point. What is the concentration of the HCl?

17 During a titration, a student overshoots the endpoint by adding too much NaOH. How will this affect the calculated concentration of the unknown HCl?

18 50 mL of 0.100 M HCl is mixed with 30 mL of 0.100 M NaOH. What is the resulting solution?

19 The particle diagram for a weak acid solution would look MOST like which of the following descriptions?

20 Water acting as a base in the reaction HCl + H₂O → H₃O⁺ + Cl⁻ is an example of water being:

Short Answer Questions

SA1 A student titrates 20.00 mL of an unknown HCl solution with 0.150 M NaOH. The endpoint is reached after adding 36.00 mL of NaOH. (a) Write the balanced neutralization equation. (b) Calculate the molarity of the HCl solution. Show all work.

SA2 A 0.20 M solution of a weak acid HA has Ka = 4.0 × 10⁻⁵. (a) Set up the ICE chart. (b) Calculate [H⁺] (use approximation x² ≈ Ka × C₀). (c) Calculate the pH. (d) Calculate % dissociation.

SA3 40.0 mL of 0.100 M HCl is mixed with 25.0 mL of 0.100 M NaOH (an incomplete titration). (a) Determine which reactant is in excess and by how many moles. (b) Calculate [H⁺] or [OH⁻] of the resulting solution. (c) Calculate the pH.