Acids & Bases Memory Tricks | Chemistry

🧪 Acids & Bases — 9 Memory Tricks · Click any card to expand · Save favorites · Switch to Flashcard or Quiz mode below

pH Scale

pH 7 = Pure water = Perfect neutral

pH < 7 = acidic · pH = 7 = neutral · pH > 7 = basic (alkaline)

The pH scale runs 0–14. Pure water sits at exactly 7 — perfect neutral. Every unit below 7 is 10× more acidic; every unit above is 10× more basic. Battery acid sits near 0; drain cleaner near 14. Stomach acid is about pH 2; blood must stay between 7.35–7.45.

Difficulty: Beginner

pH formula

pH = -log[H⁺]. If [H⁺] = 10⁻³ M, then pH = 3. Each unit change = 10× change in [H⁺].

pOH

pOH = -log[OH⁻]. At 25°C: pH + pOH = 14. Find one, calculate the other instantly.

Common pH values

Lemon juice ≈ 2 · Vinegar ≈ 3 · Coffee ≈ 5 · Blood ≈ 7.4 · Baking soda ≈ 9 · Bleach ≈ 13.

Logarithmic scale

pH 3 is 10× more acidic than pH 4, and 100× more acidic than pH 5. This trips up students who treat pH as linear.

Strong Acids

HCl HBr HI — "Have Big Ions"

Plus HNO₃ · H₂SO₄ · HClO₄ — the 6 strong acids that dissociate 100%

The 6 strong acids dissociate completely in water — no equilibrium, no Ka needed. Have Big Ions covers HCl, HBr, HI (the hydrohalic acids except HF). Add HNO₃ (nitric), H₂SO₄ (sulfuric), and HClO₄ (perchloric) to complete the set. Everything else is a weak acid.

Difficulty: Beginner

Why not HF?

HF is a weak acid — the H-F bond is extremely strong (F is small, holds tightly). Despite F being the most electronegative element, the bond strength prevents full dissociation.

H₂SO₄ is diprotic

First dissociation is complete (strong). Second dissociation is weak (HSO₄⁻ ⇌ H⁺ + SO₄²⁻). Only the first proton counts as "strong."

Strong acid calculation

For strong acids, [H⁺] = molarity of acid. 0.01 M HCl → [H⁺] = 0.01 M → pH = 2. No Ka needed.

Strong bases

Group 1 hydroxides (NaOH, KOH, LiOH) and Group 2 hydroxides (Ca(OH)₂, Ba(OH)₂) are the strong bases — dissociate completely.

Acid-Base Theory

BL (Brønsted-Lowry): Proton Donor/Acceptor · Lewis: Electron Pair

Brønsted-Lowry = H⁺ transfer · Lewis = electron pair donation

Two theories, two definitions. Brønsted-Lowry: acid = proton (H⁺) donor, base = proton acceptor. Lewis: acid = electron pair acceptor, base = electron pair donor. Lewis is broader — it includes reactions with no proton transfer at all (like BF₃ + NH₃). Every BL acid is also a Lewis acid, but not vice versa.

Difficulty: Intermediate

Conjugate pairs

In BL theory: HA ⇌ H⁺ + A⁻. HA and A⁻ are a conjugate acid-base pair. Strong acid → very weak conjugate base. Weak acid → stronger conjugate base.

Amphoteric species

Water is amphoteric — it can act as both acid and base (donates OR accepts H⁺). HSO₄⁻ is also amphoteric.

Lewis examples

BF₃ = Lewis acid (accepts electron pair into empty orbital). NH₃ = Lewis base (donates lone pair). No proton transfer occurs — BL theory can't describe this reaction.

Arrhenius (original)

Oldest definition: acid produces H⁺ in water, base produces OH⁻. Most limited — only works in aqueous solution. BL and Lewis are more general.

Buffers

Buffer = Weak Acid + Its Salt (conjugate base)

Resists pH change by absorbing added H⁺ or OH⁻

A buffer contains a weak acid and its conjugate base (the salt). When acid is added — the base component neutralizes it. When base is added — the acid component neutralizes it. The buffer resists large pH changes. Blood (carbonic acid/bicarbonate) is the most important biological buffer system in your body.

Difficulty: Intermediate

Henderson-Hasselbalch

pH = pKa + log([A⁻]/[HA]). When [A⁻] = [HA], pH = pKa exactly. Best buffering capacity is within 1 pH unit of pKa.

How it works (add acid)

Add H⁺ → A⁻ + H⁺ → HA. The conjugate base absorbs the extra proton. pH barely changes because [H⁺] increase is minimal.

How it works (add base)

Add OH⁻ → HA + OH⁻ → A⁻ + H₂O. The weak acid donates a proton to neutralize the base. Again, pH barely changes.

Buffer capacity

Higher concentration of buffer components = greater capacity. A buffer is "exhausted" when one component is used up entirely — then pH changes rapidly.

Ka & Kb

Ka × Kb = Kw = 1.0 × 10⁻¹⁴

Conjugate pair Ka and Kb always multiply to Kw at 25°C

For any conjugate acid-base pair, Ka × Kb = Kw (1.0 × 10⁻¹⁴ at 25°C). This means: strong acid (large Ka) → very weak conjugate base (tiny Kb). Know Ka for an acid and you can instantly calculate Kb for its conjugate base — no table lookup needed.

Difficulty: Intermediate

pKa + pKb = 14

Taking -log of both sides: pKa + pKb = pKw = 14 at 25°C. If pKa = 4.74 (acetic acid), then pKb of acetate = 14 - 4.74 = 9.26.

Strong acids

HCl Ka ≈ 10⁷ (essentially infinite). Conjugate base Cl⁻ has Kb ≈ 10⁻²¹ — essentially zero. Cl⁻ is such a weak base it doesn't affect pH at all.

Weak acid calculation

Acetic acid Ka = 1.8×10⁻⁵. pH of 0.1M solution: use ICE table or approximate [H⁺] = √(Ka × C) = √(1.8×10⁻⁶) ≈ 0.00134 M → pH ≈ 2.87.

Kw itself

Kw = [H⁺][OH⁻] = 1.0×10⁻¹⁴ at 25°C. In pure water: [H⁺] = [OH⁻] = 1.0×10⁻⁷ M → pH = 7. Kw increases with temperature — pure water at 37°C has pH slightly below 7.

Titration

Equivalence Point ≠ Neutral — depends on the pair

Strong/strong = pH 7 · Weak acid/strong base = pH > 7 · Strong acid/weak base = pH < 7

The equivalence point is where moles of acid = moles of base — NOT necessarily pH 7. Strong acid + strong base = pH 7 (neutral salt). Weak acid + strong base = pH > 7 (basic, because conjugate base hydrolyzes). Strong acid + weak base = pH < 7 (acidic). This is one of the most tested concepts in acid-base chemistry.

Difficulty: Intermediate

Half-equivalence point

Halfway to equivalence: [HA] = [A⁻] → pH = pKa. This is the best point for buffer activity and is how you determine pKa experimentally.

Titration curve shape

Strong/strong: steep S-curve, sharp equivalence point. Weak acid/strong base: flatter S-curve, buffer region visible, equivalence above pH 7.

Indicator choice

Indicator must change color at or near the equivalence point. Phenolphthalein (8–10) = weak acid titrations. Methyl orange (3–4) = weak base titrations.

Calculation shortcut

M₁V₁ = M₂V₂ at equivalence point. If 25 mL of 0.1M NaOH neutralizes HCl: 25×0.1 = V₂×M₂. Solve for unknown concentration or volume.

Polyprotic

H₂SO₄ loses H⁺ in steps — first strong, second weak

H₂CO₃ · H₂SO₄ · H₃PO₄ — each proton has its own Ka, getting smaller

Polyprotic acids lose protons one at a time. Each successive Ka is much smaller than the last — it gets harder to remove each proton because the ion becomes more negative. For most calculations, only Ka₁ matters because Ka₂ is so much smaller. H₂CO₃ (carbonic acid) is the most tested polyprotic acid.

Difficulty: Advanced

Carbonic acid system

H₂CO₃ ⇌ H⁺ + HCO₃⁻ (Ka₁ = 4.3×10⁻⁷) · HCO₃⁻ ⇌ H⁺ + CO₃²⁻ (Ka₂ = 4.8×10⁻¹¹). Critical for blood pH buffering — CO₂ drives Ka₁ equilibrium.

Phosphoric acid

H₃PO₄ has 3 protons. Ka₁ = 7.5×10⁻³ · Ka₂ = 6.2×10⁻⁸ · Ka₃ = 2.2×10⁻¹³. Each is about 10,000× smaller — each proton much harder to remove.

Intermediate species

HCO₃⁻ and H₂PO₄⁻ are amphoteric — they can donate OR accept a proton. pH of intermediate = (pKa₁ + pKa₂)/2 — a useful shortcut.

Simplification rule

If Ka₁/Ka₂ > 1000, treat as monoprotic using only Ka₁. The second dissociation contributes negligibly to [H⁺].

ICE Tables

ICE (I=Initial concentrations, C=Change using stoichiometry, E=Equilibrium concentrations) — Initial, Change, Equilibrium

Set up every weak acid/base equilibrium problem the same way

ICE tables are the universal tool for weak acid/base problems. Initial — write starting concentrations. Change — use +x and -x based on stoichiometry. Equilibrium — add Initial + Change. Then plug into Ka or Kb expression and solve for x. The 5% rule: if x/initial < 5%, the approximation x ≈ √(Ka × C) is valid.

Difficulty: Intermediate

Example setup

0.1M acetic acid: HA ⇌ H⁺ + A⁻. I: 0.1, 0, 0. C: -x, +x, +x. E: 0.1-x, x, x. Ka = x²/(0.1-x) ≈ x²/0.1 = 1.8×10⁻⁵. x = [H⁺] = 0.00134M. pH = 2.87.

5% approximation rule

If x/initial < 5%, the approximation (0.1-x ≈ 0.1) is valid and saves quadratic math. Check: 0.00134/0.1 = 1.34% — approximation valid here.

When to use quadratic

If x/initial > 5%, use the quadratic formula: x² + Ka·x - Ka·C = 0. Or iterate — plug approximation back in until stable.

Weak bases

Same ICE setup but use Kb. NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. Solve for [OH⁻], then pOH, then pH = 14 - pOH.

Salt Hydrolysis

Salt from strong + strong = neutral · weak + strong = NOT neutral

Salts of weak acids or bases hydrolyze — their solutions are NOT pH 7

When a salt dissolves, its ions may react with water (hydrolyze). NaCl — Na⁺ from strong base, Cl⁻ from strong acid — neither reacts with water → pH 7. CH₃COONa — acetate from weak acid — acetate hydrolyzes (A⁻ + H₂O ⇌ HA + OH⁻) → basic solution. NH₄Cl — ammonium from weak base — hydrolyzes → acidic solution.

Difficulty: Intermediate

Quick prediction rule

Strong acid + strong base salt → neutral (pH 7). Weak acid + strong base salt → basic. Strong acid + weak base salt → acidic. Weak acid + weak base → compare Ka vs Kb to determine.

Acetate example

Sodium acetate (CH₃COONa) in water: CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻. Produces OH⁻ → basic solution. pH of 0.1M solution ≈ 8.87.

Ammonium example

NH₄Cl in water: NH₄⁺ ⇌ NH₃ + H⁺. Produces H⁺ → acidic solution. pH of 0.1M solution ≈ 5.12.

Exam tip

Identify the parent acid and base of the salt. If either parent is weak, the solution won't be neutral. This determines which way pH shifts.

🎓 Common Exam Questions

Q: What are the 6 strong acids and why doesn't HF qualify?

A: The 6 strong acids dissociate 100% in water: HCl (hydrochloric), HBr (hydrobromic), HI (hydroiodic) — remember with 'Have Big Ions' — plus HNO3 (nitric), H2SO4 (sulfuric, first dissociation only), and HClO4 (perchloric). HF is a weak acid because the H-F bond is exceptionally strong — fluorine is so small and electronegative that it holds the bond too tightly for complete dissociation, despite F being the most electronegative element.

Q: Explain the ICE table method — what does each letter stand for and when do you use it?

A: ICE = Initial, Change, Equilibrium. Used whenever a reaction reaches equilibrium and you need to find equilibrium concentrations. Set up: write Initial concentrations (often the starting concentration for reactants, 0 for products). Change row: use stoichiometric ratios with +x or -x. Equilibrium row: Initial + Change. Then substitute into Ka or Kb expression and solve for x. Example: 0.1M acetic acid (Ka = 1.8×10⁻⁵): I=0.1, C=-x, E=0.1-x. Ka = x²/(0.1-x) ≈ x²/0.1. x = √(1.8×10⁻⁶) = 1.34×10⁻³ M = [H⁺]. pH = 2.87.

Q: What makes a good buffer and how does the Henderson-Hasselbalch equation work?

A: A buffer resists pH change — it contains a weak acid and its conjugate base (its salt). Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). When [A⁻] = [HA], pH = pKa exactly. A buffer works best within ±1 pH unit of its pKa. Example: acetic acid/acetate buffer (pKa 4.74) works from pH 3.74–5.74. Buffer capacity is highest at pH = pKa. Blood is buffered at pH 7.4 mainly by the carbonic acid/bicarbonate system (pKa 6.1) — supplemented by phosphate and protein buffers.

Q: Compare Brønsted-Lowry and Lewis acid-base theories.

A: Brønsted-Lowry (BL): acid = proton (H⁺) donor, base = proton (H⁺) acceptor. Requires H⁺ transfer. Every BL acid-base reaction produces conjugate pairs. Lewis: acid = electron pair acceptor, base = electron pair donor. Much broader — includes reactions with no H⁺ at all. Example: BF3 + NH3 → BF3-NH3 (Lewis acid-base, not BL). AlCl3 reacting with Cl⁻ is a Lewis acid-base reaction. Every BL acid is a Lewis acid, but not vice versa. Lewis theory explains metal complex formation and many organic reactions.

Q: What is salt hydrolysis and how do you predict whether a salt solution is acidic, basic, or neutral?

A: Salt hydrolysis occurs when ions from a dissolved salt react with water to change pH. Rules: Strong acid + Strong base → neutral salt (NaCl, pH 7 — neither ion hydrolyzes). Weak acid + Strong base → basic salt (CH3COONa — acetate ion accepts H⁺ from water, raising pH). Strong acid + Weak base → acidic salt (NH4Cl — ammonium ion donates H⁺ to water, lowering pH). Weak acid + Weak base → depends on relative Ka vs Kb values. Key insight: the conjugate of a strong acid/base does NOT hydrolyze; the conjugate of a weak acid/base DOES.

Memory tricks for pH, buffers & acid-base theory

Memory Tricks