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Unit 8 · Acids & Bases Unit Hub Flashcards Cheat Sheet Essentials Visual Review MC Practice SAQ Practice

AP Chemistry Unit 8 Visual Review

A topic-by-topic visual walkthrough of Acids & Bases — pH, weak-acid equilibria, titrations, buffers, and the Henderson-Hasselbalch equation.

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TOPIC 8.1 Introduction to Acids and Bases Arrhenius model Acid → produces H⁺ (H₃O⁺) in water. Base → produces OH⁻ in water. Example: HCl → H⁺ + Cl⁻ NaOH → Na⁺ + OH⁻ Brønsted–Lowry model (broader) Acid = proton (H⁺) DONOR. Base = proton (H⁺) ACCEPTOR. Every acid has a conjugate base; every base a conjugate acid. Conjugate acid–base pairs differ by one H⁺ HA + H₂O ⇌ A⁻ + H₃O⁺. Here HA / A⁻ is one conjugate pair and H₂O / H₃O⁺ is the other. Water is AMPHOTERIC — it can act as an acid or a base. The stronger the acid, the weaker its conjugate base. Autoionization: 2 H₂O ⇌ H₃O⁺ + OH⁻, with K_w = [H₃O⁺][OH⁻] = 1.0×10⁻¹⁴ at 25 °C. Brønsted–Lowry: acids donate H⁺, bases accept H⁺ — conjugate pairs differ by exactly one proton. The Review Hub · AP Chemistry Unit 8 TOPIC 8.2 pH and pOH of Strong Acids and Bases THE pH SCALE (25 °C) 0 acidic 7 neutral 14 basic Core formulas pH = −log[H₃O⁺] pOH = −log[OH⁻] pH + pOH = 14 [H₃O⁺][OH⁻] = 1.0×10⁻¹⁴ (the relations at 25 °C) Strong = 100% ionized Strong acids/bases fully dissociate, so [H₃O⁺] or [OH⁻] = the concentration. Strong acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄ Worked example 0.010 M HCl (strong): [H₃O⁺] = 0.010 M pH = −log(0.010) = 2.0 0.010 M NaOH → pOH = 2, so pH = 14 − 2 = 12. For strong acids/bases, [H₃O⁺] or [OH⁻] equals the concentration — then take −log. The Review Hub · AP Chemistry Unit 8 TOPIC 8.3 Weak Acid and Base Equilibria Weak acids/bases only PARTIALLY ionize HA + H₂O ⇌ H₃O⁺ + A⁻ reaches equilibrium. K_a = [H₃O⁺][A⁻] / [HA]. For a weak base: B + H₂O ⇌ BH⁺ + OH⁻, K_b = [BH⁺][OH⁻] / [B]. A SMALLER K means a WEAKER acid/base. ICE + approximation Set up an ICE table for [HA]₀. K_a = x² / ([HA]₀ − x). If K_a is small, [HA]₀ − x ≈ [HA]₀, so x = √(K_a·[HA]₀) = [H₃O⁺]. Valid when ionization < ~5%. K_a · K_b = K_w For a conjugate acid–base pair, K_a × K_b = K_w = 1.0×10⁻¹⁴. Also pK_a + pK_b = 14. A strong acid ⇒ its conjugate base is very weak (tiny K_b). Weak acids need an ICE table & K_a; the small-x approximation gives [H₃O⁺] = √(K_a·[HA]₀). The Review Hub · AP Chemistry Unit 8 TOPIC 8.4 Acid–Base Reactions and Buffers Neutralization Acid + base → salt + water. Net ionic (strong + strong): H⁺ + OH⁻ → H₂O Strong+strong goes essentially to completion. What makes a buffer A buffer contains a weak acid AND its conjugate base (or a weak base + conjugate acid) in comparable amounts. e.g. CH₃COOH / CH₃COO⁻ or NH₃ / NH₄⁺. How a buffer resists pH change Add ACID → the conjugate base neutralizes it: A⁻ + H₃O⁺ → HA + H₂O. Add BASE → the weak acid neutralizes it: HA + OH⁻ → A⁻ + H₂O. Ratio shifts only slightly, so pH barely moves. The buffer works because both partners are present to absorb added H⁺ or OH⁻. A buffer = weak acid + its conjugate base — each partner neutralizes added base or acid. The Review Hub · AP Chemistry Unit 8 TOPIC 8.5 Acid–Base Titrations STRONG ACID + STRONG BASE equivalence point (pH 7) pH volume of titrant added → Equivalence point moles acid = moles base added. Found at the STEEP vertical jump. Strong+strong: pH = 7. Weak acid + strong base: pH > 7 (basic salt). Half-equivalence point Halfway to equivalence, [HA] = [A⁻], so pH = pK_a. This is the point of maximum buffering. Great way to read K_a straight off the curve. Equivalence = moles matched (steep jump); half-equivalence: pH = pK_a. The Review Hub · AP Chemistry Unit 8 TOPIC 8.6 Molecular Structure of Acids and Bases Bond strength & polarity Acid strength depends on how easily the H–X bond breaks to release H⁺. A WEAKER, more polar H–X bond gives a STRONGER acid. Binary: HF < HCl < HBr < HI (bond weakens down). Oxyacids & conjugate stability More electronegative atoms / more O atoms pull electron density away, so the O–H bond is more polar/acidic. HClO₄ > HClO₃ > HClO₂ > HClO. A more STABLE conjugate base ⇒ stronger acid. Strength vs. concentration — don't confuse them STRENGTH = degree of ionization (an intrinsic property). CONCENTRATION = amount of acid per liter. A dilute strong acid can have a higher pH than a concentrated weak acid — the two ideas are independent. "Strong" is about how completely it dissociates, not how much is dissolved. Weaker/more-polar H–X and a more stable conjugate base = stronger acid. The Review Hub · AP Chemistry Unit 8 TOPIC 8.7 pH and pKa pK_a = −log K_a Smaller pK_a → larger K_a → STRONGER acid. Each pK_a unit = a factor of 10 in acid strength. Comparing pH and pK_a pH < pK_a → mostly protonated (HA). pH > pK_a → mostly deprotonated (A⁻). pH = pK_a → [HA] = [A⁻] exactly. This is the logic behind titration curves. Why pK_a is convenient K_a values span many orders of magnitude (e.g. 1.8×10⁻⁵ for acetic acid). Taking −log turns them into a friendly linear scale: acetic acid pK_a ≈ 4.74. Compare acids directly by their pK_a values. Lower pK_a = stronger acid, holds a proton less tightly. Lower pK_a = stronger acid; comparing pH to pK_a tells you if HA or A⁻ dominates. The Review Hub · AP Chemistry Unit 8 TOPIC 8.8 Properties of Buffers A buffer's pH depends on the RATIO, not the total amount pH is set by [A⁻]/[HA] and pK_a. Diluting a buffer barely changes its pH, because both concentrations drop by the same factor — the ratio (and therefore pH) stays essentially the same. Choosing a buffer Pick a weak acid whose pK_a is CLOSE to the target pH. Effective range ≈ pK_a ± 1. Within this range [A⁻]/[HA] runs from 1:10 to 10:1. Best resistance at 1:1 A buffer resists change BEST when [A⁻] = [HA] (pH = pK_a), because both partners are equally abundant. Far from a 1:1 ratio, one partner runs low and the buffer weakens. Buffer pH follows the [A⁻]/[HA] ratio; pick pK_a near the target, best at 1:1. The Review Hub · AP Chemistry Unit 8 TOPIC 8.9 Henderson–Hasselbalch Equation pH = pK_a + log ( [A⁻] / [HA] ) Reading it [A⁻] = [HA] → log 1 = 0 → pH = pK_a. More base (A⁻) → log > 0 → pH ↑. More acid (HA) → log < 0 → pH ↓. Because it's a ratio, you can use MOLES instead of molarity (same volume cancels). Worked example Acetic acid / acetate, pK_a = 4.74. 0.10 M CH₃COOH + 0.10 M CH₃COO⁻: pH = 4.74 + log(1) = 4.74. Double the acetate → 0.20 / 0.10: pH = 4.74 + log(2) = 5.04. pH = pK_a + log([A⁻]/[HA]) — the fast route to any buffer's pH. The Review Hub · AP Chemistry Unit 8 TOPIC 8.10 Buffer Capacity Capacity = how much acid/base a buffer can absorb Buffer capacity is the amount of strong acid or base a buffer can neutralize before its pH changes significantly. It's set by the TOTAL concentration of the buffer components — not by the pH. Bigger capacity when… • Concentrations of HA and A⁻ are both HIGH (more to react). • The ratio [A⁻]/[HA] is near 1:1 (balanced against acid AND base). Note: capacity is independent of pH itself. Capacity vs. pH — separate ideas The RATIO sets the pH; the TOTAL amount sets the capacity. A dilute and a concentrated buffer can share a pH but differ hugely in how much they can absorb. Capacity grows with total concentration — the ratio sets pH, the amount sets capacity. The Review Hub · AP Chemistry Unit 8 TOPIC 8.11 pH and Solubility pH shifts solubility when the anion is a base If a salt's anion is the conjugate base of a WEAK acid (F⁻, OH⁻, CO₃²⁻, S²⁻…), adding H⁺ removes it from solution. By Le Châtelier, the dissolving equilibrium shifts right → solubility INCREASES. Example: CaF₂ CaF₂ ⇌ Ca²⁺ + 2 F⁻ Add acid: H⁺ + F⁻ → HF removes F⁻. Equilibrium shifts right, dissolving more CaF₂ → MORE soluble in acid. Same idea for CaCO₃, Mg(OH)₂, metal sulfides. When pH does NOT matter If the anion is the conjugate base of a STRONG acid (Cl⁻, Br⁻, NO₃⁻), it won't react with added H⁺. So salts like AgCl show little solubility change with pH. Basic anions dissolve more in acid; anions of strong acids are unaffected by pH. The Review Hub · AP Chemistry Unit 8
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This is great for review the night before the exam — fast, visual, and covers everything you need to remember about Unit 8's acid-base content.