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Unit 8: Acids and Bases flashcards

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8.1 Introduction to Acids and Bases

What is the autoionization of water, and what is K_w?
Water is amphiprotic, so it ionizes itself: 2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq). K_w = [H₃O⁺][OH⁻] (water is left out as the pure solvent). At 25 °C, K_w = 1.0 × 10⁻¹⁴, so pure water has [H₃O⁺] = [OH⁻] = 1.0 × 10⁻⁷ M.
How are pH and pOH defined, and how are they related?
pH = −log[H₃O⁺]; pOH = −log[OH⁻]. Reverse: [H₃O⁺] = 10⁻ᵖᴴ. pH + pOH = pK_w = 14.00 at 25 °C. Use log₁₀, not ln. pH 4.25 gives pOH 9.75.
What do acidic, basic and neutral mean in terms of ion concentrations?
Neutral: [H₃O⁺] = [OH⁻]. Acidic: [H₃O⁺] > [OH⁻]. Basic: [OH⁻] > [H₃O⁺]. Acidic solutions still contain OH⁻; find it with K_w.
Is a neutral solution always pH 7?
No, only at 25 °C. Autoionization is endothermic, so K_w increases with temperature. Heated pure water has pH below 7 but is still neutral: pH = pOH = pK_w ÷ 2.
How much does [H₃O⁺] change per pH unit?
A factor of 10 per unit; the scale is logarithmic. pH 3 has 100 times the [H₃O⁺] of pH 5. Lower pH means higher [H₃O⁺].
How do you find [OH⁻] from a pH at 25 °C?
pOH = 14.00 − pH, then [OH⁻] = 10⁻ᵖᴼᴴ, or [OH⁻] = K_w ÷ [H₃O⁺]. pH 9.00 gives pOH 5.00 and [OH⁻] = 1.0 × 10⁻⁵ M.
What is the significant-figure rule for pH?
The number of decimal places in a pH equals the number of significant figures in [H₃O⁺]. [H₃O⁺] = 1.0 × 10⁻³ M gives pH 3.00.

8.2 pH and pOH of Strong Acids and Bases

Which acids and bases are strong?
Strong acids: HCl, HBr, HI, HClO₄, HNO₃, H₂SO₄ (first proton only). Strong bases: group 1 hydroxides (LiOH, NaOH, KOH) and Ca(OH)₂, Sr(OH)₂, Ba(OH)₂. They ionize or dissociate completely, so no equilibrium table is needed.
How do you find the pH of a strong acid solution?
[H₃O⁺] = c, so pH = −log c. 0.0010 M HCl has pH 3.00. In the solution [H₃O⁺] = [A⁻] = c, [HA] ≈ 0 (HA is not a major species) and [OH⁻] = K_w ÷ c.
How do you find the pH of a strong base solution?
Group 1 hydroxide: [OH⁻] = c. Group 2 hydroxide: [OH⁻] = 2c. Then pOH = −log[OH⁻] and pH = 14.00 − pOH. 0.20 M NaOH: pOH 0.70, pH 13.30. 0.0050 M Ba(OH)₂: [OH⁻] = 0.010 M, pH 12.00.
How does diluting a strong acid change its pH?
c₂ = c₁V₁ ÷ V₂; each tenfold dilution moves pH 1 unit towards 7. 0.010 M HNO₃ (pH 2.00) diluted tenfold has pH 3.00.
What is the pH of 10⁻⁸ M HCl?
Close to 7, slightly below, not 8. Below about 10⁻⁶ M, water's own H₃O⁺ is not negligible, and an acid can never make water basic.
What is the difference between strong and concentrated?
Strong means completely ionized; concentrated means a lot of solute per litre. A dilute solution of a strong acid is still strong.

8.3 Weak Acid and Base Equilibria

How does a weak acid behave in water?
It reacts with water to make H₃O⁺, but only a small percentage of molecules ionize, so [H₃O⁺] ≪ c. Species: [HA] ≫ [H₃O⁺] ≈ [A⁻] ≫ [OH⁻]. pH is not −log c.
What are K_a and pK_a?
K_a = [H₃O⁺][A⁻] / [HA]; pK_a = −log K_a. Smaller K_a, larger pK_a = weaker acid. K_a changes only with temperature, not with dilution.
How do you calculate the pH of a weak acid solution?
x = [H₃O⁺] ≈ √(K_a × c), valid when x below 5% of c (check it; otherwise solve the quadratic). 0.20 M with K_a = 5.0 × 10⁻⁶: x = 1.0 × 10⁻³ M, pH 3.00, 0.50% ionized.
How do you calculate the pH of a weak base solution?
K_b = [OH⁻][HB⁺] / [B]; x = [OH⁻] ≈ √(K_b × c). Then convert: pOH = −log x, pH = 14.00 − pOH. 0.10 M with K_b = 1.0 × 10⁻⁵: [OH⁻] = 1.0 × 10⁻³ M, pH 11.00.
How do you find K_a from a measured pH?
x = 10^(−pH) = [H₃O⁺] = [A⁻], so K_a = x² / (c − x). Percent ionization = (x ÷ c) × 100%.
How are K_a and K_b of a conjugate pair related?
K_a × K_b = K_w and pK_a + pK_b = 14.00 at 25 °C. An acid with pK_a 4.00 has a conjugate base with pK_b 10.00, K_b = 1.0 × 10⁻¹⁰.
What happens to percent ionization when a weak acid is diluted?
It rises, even though [H₃O⁺] falls. K_a stays constant.

8.4 Acid-Base Reactions and Buffers

How do you calculate pH after mixing an acid and a base?
Convert every volume and molarity to moles. Let the strong reagent react completely (limiting reactant; one-way, no ICE table). Then identify what is left and choose the method. Use moles, not concentrations, and divide by the total volume.
What is the pH when equal moles of strong acid and strong base react?
7.00 at 25 °C: only water and spectator ions remain. 10.0 mL of 0.10 M HCl + 10.0 mL of 0.10 M NaOH gives pH 7.00.
How do you find pH when a strong reagent is in excess?
[H₃O⁺] or [OH⁻] = excess moles ÷ total volume. 5.0 × 10⁻³ mol weak acid + 6.0 × 10⁻³ mol NaOH in 100. mL leaves 1.0 × 10⁻³ mol OH⁻: 0.010 M, pH 12.00.
What forms when a weak acid is partly neutralized by a strong base?
A buffer: HA and A⁻ in comparable amounts. pH = pK_a + log(n base form ÷ n acid form). 4.0 × 10⁻³ mol HA (pK_a 4.20) + 1.0 × 10⁻³ mol NaOH: 4.20 + log(1.0/3.0) = 3.72.
What is the pH when equal moles of a weak acid and strong base react?
Not 7. Only A⁻ remains, a weak base: K_b = K_w ÷ K_a and [OH⁻] ≈ √(K_b × c), so pH is slightly above 7. For weak base + strong acid, only HB⁺ remains and pH is slightly below 7.
Do a weak acid and a weak base react completely?
No. They reach an equilibrium with K = K_a(HA) ÷ K_a(HB⁺). Products are favoured if HA is the stronger acid.

8.5 Acid-Base Titrations

Where is the equivalence point on a pH titration curve?
At the middle of the steepest part of the curve, not at the top. At equivalence, n(titrant) = n(analyte) for monoprotic acids and bases. 20.00 mL of weak acid needing 16.00 mL of 0.125 M NaOH is 0.100 M.
Does a weak acid need less base to reach equivalence than a strong acid?
No. Strength changes the shape of the curve, not the volume needed: equal moles of a strong or weak monoprotic acid need equal moles of base.
How does a weak acid–strong base titration curve differ from a strong acid–strong base curve?
It starts at a higher pH, has a flat buffer region, a shorter steep section, and an equivalence point above 7, because A⁻ is a weak base.
How do you find K_a from a titration curve?
At half the equivalence volume, [HA] = [A⁻], so pH = pK_a. Halve the volume, not the pH. If pH = 4.74 at 8.00 mL with V(eq) = 16.00 mL, K_a = 10⁻⁴·⁷⁴ = 1.8 × 10⁻⁵. This does not work for a strong acid.
What does a weak base–strong acid titration curve look like?
The mirror image: it starts high, falls through a buffer region and has an acidic equivalence point (NH₄⁺ is a weak acid). At half-equivalence, pH = pK_a of BH⁺.
What does the titration curve of a diprotic acid look like?
One jump per proton (if each is acidic enough), with equal volumes between jumps: first equivalence at 12.0 mL means the second at 24.0 mL. pH ≈ pK_a₁ at ½V(eq1) and ≈ pK_a₂ midway between V(eq1) and V(eq2).
How do you find the pH after the equivalence point?
Excess titrant moles ÷ total volume (analyte + titrant) gives [OH⁻] or [H₃O⁺]; then convert to pH.

8.6 Molecular Structure of Acids and Bases

What determines how strong an acid is, structurally?
The stability of the conjugate base A⁻ with its negative charge. The better A⁻ holds or spreads the charge, the stronger the acid. Explain using the conjugate base, not just the acid.
Which hydrogens in a molecule are acidic?
Those bonded to electronegative atoms, usually O or a halogen. C–H hydrogens are not acidic in water. HCOOH has one acidic H, the one on O.
How does electronegativity affect acid strength across a period and in oxoacids?
Same period: a more electronegative atom bearing the charge holds it better (weakest to strongest acid): CH₄, NH₃, H₂O, HF. Oxoacids: a more electronegative central atom pulls density from O–H (weakest to strongest): HOI, HOBr, HOCl.
How do nearby electronegative atoms affect carboxylic acid strength?
The inductive effect spreads the negative charge through bonds, stabilizing A⁻ (weakest to strongest): CH₃COOH, ClCH₂COOH, Cl₂CHCOOH, Cl₃CCOOH.
Why does acid strength increase from HClO to HClO₄?
More O atoms on the central atom add inductive pull and allow resonance that shares the negative charge of the conjugate base (weakest to strongest): HClO, HClO₂, HClO₃, HClO₄.
Why is acetic acid much stronger than ethanol?
In acetate the negative charge is shared over two O atoms by resonance; in ethoxide it sits on one O. So ethoxide is the stronger base and ethanol the much weaker acid.
Why does acid strength increase from HF to HI?
Down the group the anion is larger, so the charge is spread over a bigger volume and the H–X bond is weaker. HF is weak; HI is the strongest. Electronegativity does not decide this trend.
What are typical weak acids and weak bases by structure?
Weak acids: carboxylic acids, R–COOH. Weak bases: nitrogen bases (NH₃, amines) and carboxylate ions, R–COO⁻. The weaker the acid, the stronger its conjugate base (K_a × K_b = K_w).

8.7 pH and pKa

How does comparing pH with pK_a tell you the protonation state?
pH below pK_a: HA (protonated) predominates. pH = pK_a: [HA] = [A⁻]. pH > pK_a: A⁻ (deprotonated) predominates. An acid with pK_a 5.0 is 50% deprotonated at pH 5.0.
How large is the ratio [A⁻]/[HA] at a given pH?
[A⁻]/[HA] = 10^(pH − pK_a). 1 unit above pK_a → 10; 2 units → 100. HA with pK_a 8.0 at pH 6.0 has [A⁻]/[HA] = 0.01, so HA predominates.
How do you find the predominant form of a weak base at a given pH?
Use pK_a of BH⁺ = 14.00 − pK_b (25 °C). pH below pK_a(BH⁺): BH⁺ predominates. A base with pK_b 3.4 has pK_a(BH⁺) = 10.6. Do not compare pH with pK_b.
How does an acid–base indicator work?
It is a weak acid, HIn, whose two forms have different colours. Its colour changes over roughly pK_a ± 1. It is used in tiny amounts so it does not change the pH.
How do you choose an indicator for a titration?
Choose one whose pK_a is close to the equivalence-point pH, so it changes on the steep part of the curve. Weak acid + strong base (pH above 7): phenolphthalein (8.3–10.0). Weak base + strong acid (below 7): methyl red (4.4–6.2). Strong + strong: several work.
Does pH > pK_a mean the acid is fully dissociated?
No, only that A⁻ is the larger form. And pH = pK_a does not mean neutral; it means the two forms are equal.

8.8 Properties of Buffers

What is a buffer, and what does it contain?
A solution that resists pH change when small amounts of acid or base are added. It contains large concentrations of both members of a conjugate pair: HA with A⁻, or B with BH⁺. Its pH sits near the pK_a of the weak acid, not necessarily near 7.
How does a buffer neutralize added acid and added base?
Added H₃O⁺ is removed by the base member: A⁻ + H₃O⁺ → HA + H₂O. Added OH⁻ is removed by the acid member: HA + OH⁻ → A⁻ + H₂O. Show the buffer member reacting, not H⁺ + OH⁻ → H₂O.
Why does a buffer's pH change only slightly?
The reactions go essentially to completion, swapping the strong acid or base for a weak one, and only slightly change the ratio of the pair. The pH does change, just much less than in water.
How do amounts change when acid is added to a buffer?
Base member − x, acid member + x. 0.30 mol HA and 0.20 mol A⁻ plus 0.050 mol HCl gives HA 0.35 mol and A⁻ 0.15 mol.
Which mixtures make a buffer?
Yes: weak acid + salt of its conjugate base; weak base + salt of its conjugate acid; excess weak acid + some strong base. No: weak acid + equal moles of strong base (only A⁻ remains, e.g. 0.10 mol HF + 0.10 mol NaOH); strong acid + its salt (HCl + NaCl); weak acid alone.
Why can HCl + NaCl not act as a buffer?
Cl⁻ is the negligible conjugate base of a strong acid, so it cannot remove added H₃O⁺.

8.9 Henderson-Hasselbalch Equation

What is the Henderson-Hasselbalch equation?
pH = pK_a + log([A⁻]/[HA]), with the base form on top. It is the K_a expression in log form. An acetate buffer (pK_a 4.76) with [A⁻]/[HA] = 10 has pH 5.76.
Why can you use the amounts mixed directly in a buffer calculation?
Both components are present in large amounts, so their equilibrium amounts are almost exactly the amounts mixed; no ICE table is needed. Volume cancels, so n(A⁻)/n(HA) can replace the concentration ratio after mixing.
What pH do ratios of 0.10, 1 and 10 give?
[A⁻]/[HA] = 0.10: pH = pK_a − 1. Ratio 1: pH = pK_a. Ratio 10: pH = pK_a + 1. A useful buffer stays within this range.
How do you find the ratio needed for a target buffer pH?
[A⁻]/[HA] = 10^(pH − pK_a). For pH 6.90 with pK_a 7.20: 10^(−0.30) = 0.50.
How do you use the Henderson-Hasselbalch equation for an NH₃/NH₄⁺ buffer?
Use pK_a of NH₄⁺ = 14.00 − pK_b = 9.25, not pK_b 4.75. Equal [NH₃] and [NH₄⁺] gives pH 9.25.
Does diluting a buffer change its pH?
No. Dilution leaves the ratio [A⁻]/[HA] unchanged, so the pH is unchanged.
How do you choose a conjugate pair for a buffer of a given pH?
Choose the weak acid whose pK_a is closest to the target pH, then adjust the ratio.

8.10 Buffer Capacity

What is buffer capacity?
How much strong acid or strong base a buffer can neutralize while its pH stays nearly constant. Once the limiting component is used up, excess H₃O⁺ or OH⁻ changes the pH sharply.
What sets the limit on how much acid or base a buffer can absorb?
Each reaction is 1:1 and complete. Moles of A⁻ set the limit for added acid; moles of HA set the limit for added base. 0.20 mol HA and 0.05 mol A⁻ can neutralize at most 0.05 mol HCl and 0.20 mol NaOH.
How does concentration affect buffer pH and capacity?
Multiplying both concentrations by 10 at the same ratio leaves pH unchanged but increases capacity tenfold in both directions. Two buffers can share a pH and differ tenfold in capacity.
How does the ratio of HA to A⁻ affect capacity?
More HA than A⁻: pH below pK_a, greater capacity for added base. More A⁻ than HA: pH above pK_a, greater capacity for added acid. Equal amounts: pH = pK_a, equal capacity both ways.
How should you compare capacity of buffers with different volumes?
Compare moles, not concentrations. 100 mL of 0.30 M HA/0.30 M A⁻ and 1.00 L of 0.030 M HA/0.030 M A⁻ have equal capacity: 0.030 mol of each.
How does dilution affect a buffer's pH and capacity?
pH is unchanged (ratio unchanged), but capacity per litre falls. Diluting to twice the volume halves capacity per millilitre.

8.11 pH and Solubility

Why does pH affect the solubility of some salts?
When one of the salt's ions is a weak base, a weak acid or OH⁻, added acid or base can react with it and remove it. Removing an ion makes Q below K_sp, so equilibrium shifts towards the dissolved ions and more solid dissolves. K_sp does not change; only solubility does.
Why are salts with basic anions more soluble in acid?
H₃O⁺ converts the basic anion to its conjugate acid: F⁻ + H₃O⁺ → HF + H₂O; CO₃²⁻ + 2H₃O⁺ → CO₂ + 3H₂O. CaF₂, CaCO₃, BaCO₃ and Ca₃(PO₄)₂ dissolve more in acid than in water.
How does pH affect the solubility of metal hydroxides?
In acid they are more soluble (OH⁻ + H₃O⁺ → 2H₂O removes OH⁻). In base they are less soluble because OH⁻ is a common ion: NaOH lowers the solubility of Mg(OH)₂ or Fe(OH)₃.
Why does adding acid not change the solubility of AgCl, AgBr or AgI?
Cl⁻, Br⁻ and I⁻ are conjugate bases of strong acids, so they do not react with H₃O⁺. Note HCl would actually lower a chloride's solubility by the common-ion effect.
How does pH affect a salt with a weakly acidic cation BH⁺?
OH⁻ removes BH⁺ (BH⁺ + OH⁻ → B + H₂O), so solubility increases as pH rises; acid has about no effect. BH⁺ is mostly removed once pH is above its pK_a.
What should a full explanation of pH-dependent solubility include?
Which ion is removed and by what reaction, that Q falls below K_sp, and which way the equilibrium shifts. Do not say K_sp increases. The proton goes to the basic anion or OH⁻, not the metal ion.

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