📚Study Guide: Acids and Bases
Unit 8: Acids and Bases
Acids and bases are among the most important classes of chemicals in both industrial and biological contexts. This unit covers three acid-base definitions (Arrhenius, Bronsted-Lowry, Lewis), pH calculations for strong and weak acids/bases, buffer solutions, titration curves, and hydrolysis of salts. The Bronsted-Lowry definition is most useful for AP Chemistry: acids are proton donors and bases are proton acceptors. Conjugate acid-base pairs differ by one H+. Weak acid and base dissociation are governed by Ka and Kb equilibrium constants; pKa = -log Ka and pKb = -log Kb. Buffer solutions resist pH change upon addition of small amounts of acid or base and are described by the Henderson-Hasselbalch equation. Titration curves reveal the stoichiometry and strength of acids and bases through the shape of the pH curve and the choice of indicator. The AP exam heavily emphasizes buffer calculations, titration problems (especially at the half-equivalence point where pH = pKa), and the ability to predict whether a salt solution will be acidic, basic, or neutral based on the parent acid and base strengths. Polyprotic acids and the common ion effect are also frequently tested.
Key Concepts
- Acid-Base Definitions: Arrhenius: acids produce H+, bases produce OH-. Bronsted-Lowry: acids donate H+, bases accept H+. Lewis: acids accept electron pairs, bases donate electron pairs.
- Strong vs. Weak: Strong acids/bases dissociate completely in water. Weak acids/bases establish an equilibrium described by Ka or Kb. Ka x Kb = Kw = 1.0 x 10^-14 at 25 degrees C.
- pH and pOH: pH = -log[H+]; pOH = -log[OH-]; pH + pOH = 14.00 at 25 degrees C. For strong acids/bases, [H+] or [OH-] equals initial concentration times stoichiometry.
- Weak Acid/Base Dissociation: Use ICE tables with Ka or Kb. For a weak acid HA: Ka = [H+][A-] / [HA]. If Ka is small and [HA]0 is large, assume x << [HA]0. Percent ionization = ([H+]eq / [HA]0) x 100.
- Buffers: A mixture of a weak acid and its conjugate base (or weak base and its conjugate acid). Resists pH change. Henderson-Hasselbalch: pH = pKa + log([A-]/[HA]). At half-equivalence point, [A-] = [HA], so pH = pKa.
- Titration Curves: Strong acid-strong base: equivalence point at pH 7. Weak acid-strong base: equivalence point > 7 due to conjugate base hydrolysis. Weak base-strong acid: equivalence point < 7 due to conjugate acid hydrolysis. The buffer region is flat; the equivalence point is steep.
- Indicators: Weak acids whose color changes over a specific pH range. Choose an indicator whose pKa is close to the pH at the equivalence point.
Vocabulary
- Amphoteric: A substance that can act as either an acid or a base depending on what it reacts with (e.g., water, bicarbonate).
- Conjugate Acid-Base Pair: Two species that differ by one proton (H+); the acid has one more H+ than its conjugate base.
- Hydrolysis: The reaction of an ion with water to produce H+ or OH-, resulting in acidic or basic salt solutions.
- Equivalence Point: The point in a titration where moles of acid = moles of base (for monoprotic 1:1 titrations).
- Half-Equivalence Point: The point where exactly half the volume of titrant needed to reach the equivalence point has been added; pH = pKa for weak acid titrations.
- Polyprotic Acid: An acid that can donate more than one proton (e.g., H2SO4, H3PO4), with successive Ka values decreasing dramatically (Ka1 >> Ka2 >> Ka3).
Essential Formulas
- pH = -log[H+]; [H+] = 10^-pH
- pOH = -log[OH-]; [OH-] = 10^-pOH
- pH + pOH = 14.00 (at 25 degrees C)
- Ka x Kb = Kw = 1.0 x 10^-14
- pKa = -log Ka; pKb = -log Kb
- pH = pKa + log([A-]/[HA]) (Henderson-Hasselbalch)
- Percent Ionization = ([H+]_eq / [HA]_0) x 100
Common Mistakes
- Treating Weak Acids Like Strong Acids: For weak acids, [H+] does NOT equal the initial acid concentration. You must solve using Ka and an ICE table (or approximation).
- Using Henderson-Hasselbalch for Strong Acids/Bases: The Henderson-Hasselbalch equation applies only to buffer solutions containing a weak acid/base and its conjugate.
- Confusing Equivalence Point and Endpoint: The equivalence point is the stoichiometric point. The endpoint is where the indicator changes color. They should be close but are not identical.
- Adding Volumes Incorrectly in Titrations: When calculating concentration after mixing, use the TOTAL volume (acid volume + base volume) in the denominator.
AP Exam Strategies
- Use ICE Tables for Weak Acids/Bases: Even for simple problems, set up the ICE table to avoid algebraic errors and to clearly show your reasoning.
- Memorize the Six Strong Acids: HCl, HBr, HI, HNO3, H2SO4 (first proton only), HClO4. Everything else is weak.
- Calculate pH at Half-Equivalence: Recognize that pH = pKa at the half-equivalence point in a weak acid-strong base titration. This is a favorite exam question.
- Predict Salt pH: Strong acid + strong base -> neutral. Strong acid + weak base -> acidic. Weak acid + strong base -> basic. Weak acid + weak base -> depends on Ka vs Kb.
Real-World Applications
- Blood Buffer System: The carbonic acid/bicarbonate buffer maintains blood pH within the narrow range of 7.35-7.45; disruptions cause acidosis or alkalosis.
- Antacids: Bases like calcium carbonate and magnesium hydroxide neutralize stomach acid (HCl) to relieve heartburn.
- Ocean Acidification: Increased atmospheric CO2 dissolves in seawater, forming carbonic acid and lowering ocean pH, threatening calcifying organisms like corals and shellfish.