Acids and bases are everywhere in chemistry and in life. Unit 8 teaches you how to identify acids and bases, calculate pH and pOH, understand strong and weak acids, and solve problems involving buffers and titrations. You'll learn the relationship between acid and base strength and molecular structure. These concepts connect directly to equilibrium from Unit 7 and to electrochemistry in Unit 9.
🎯 What You Need to Know for the Exam
Unit 8 makes up about 11-15% of the AP Chemistry exam. This is one of the heavier units. Focus your energy on these priorities:
- pH, pOH, and water equilibrium: calculating pH and pOH from concentrations, understanding Kw and neutral solutions
- Strong acids and bases: complete ionization, calculating pH directly from concentration
- Weak acids and bases: partial ionization, using Ka or Kb to find pH, calculating percent ionization
- Buffers: understanding how they resist pH change, using the Henderson-Hasselbalch equation, calculating buffer capacity
- Acid-base titrations: interpreting titration curves, finding equivalence points, calculating concentrations
- Acid-base reactions: strong with strong, weak with strong, predicting products and calculating pH
What's in this review:
- Introduction to Acids and Bases
- pH and pOH of Strong Acids and Bases
- Weak Acid and Base Equilibria
- Acid-Base Reactions and Buffers
- Acid-Base Titrations
- Molecular Structure of Acids and Bases
- pH and pKa
- Properties of Buffers
- Henderson-Hasselbalch Equation
- Buffer Capacity
- pH and Solubility
- Study Tips for Unit 8
- Summary, Review Questions & Practice
Topic 8.1: Introduction to Acids and Bases
Water is the solvent in most acid-base chemistry. Water has a remarkable property: it undergoes autoionization, producing hydronium ions (H3O+) and hydroxide ions (OH-). In pure water, the concentration of H3O+ equals the concentration of OH-. The product of these concentrations is constant at a given temperature, called Kw.
At 25°C, Kw = [H3O+][OH-] = 1.0 × 10^-14. This is the fundamental equation of water equilibrium. It means that if you know the concentration of hydronium ions, you can calculate the concentration of hydroxide ions, and vice versa.
We measure acidity using pH, defined as pH = -log[H3O+]. Similarly, pOH = -log[OH-]. A key relationship is pH + pOH = 14.0 at 25°C. A neutral solution has pH = 7 and pOH = 7. Acidic solutions have pH < 7. Basic solutions have pH > 7.
Key concepts to know:
- Autoionization of water: H2O ⇌ H3O+ + OH-. Both hydronium and hydroxide ions are present in all aqueous solutions.
- Kw = 1.0 × 10^-14 at 25°C: The product [H3O+][OH-] is constant. Kw changes with temperature.
- pH scale: pH = -log[H3O+]. Ranges from 0 (very acidic) to 14 (very basic). pH = 7 is neutral at 25°C.
- pOH scale: pOH = -log[OH-]. Related to pH by pH + pOH = 14.0 at 25°C.
- Neutral solution: pH = pOH = 7.0 at 25°C.
⚠ Watch out for:
Don't forget that Kw changes with temperature. At temperatures other than 25°C, the pH of pure neutral water deviates from 7.0. The exam sometimes includes this detail. Also, remember that pH + pOH = 14.0 only at 25°C. At other temperatures, use the actual Kw value.
Topic 8.2: pH and pOH of Strong Acids and Bases
Strong acids completely ionize in water. When you dissolve a strong acid like HCl, it dissociates into H+ and Cl-. The concentration of H3O+ equals the initial concentration of the acid. If you dissolve 0.1 M HCl, you get [H3O+] = 0.1 M, and pH = 1.
Strong bases like NaOH completely dissociate. When you dissolve a strong base, it produces OH- ions. A 0.1 M NaOH solution gives [OH-] = 0.1 M. But be careful: for Group II hydroxides like Ca(OH)2, each formula unit produces two OH- ions. A 0.1 M Ca(OH)2 solution gives [OH-] = 0.2 M.
Common strong acids include HCl, HBr, HI, HClO4, H2SO4, and HNO3. Common strong bases include Group I and Group II hydroxides. If the acid or base is not on the strong list, treat it as weak.
Key concepts to know:
- Strong acid ionization: Complete. [H3O+] = initial concentration of acid.
- Strong base dissociation: Complete. [OH-] = initial concentration of base (or 2x for Group II hydroxides).
- Direct pH calculation: For strong acids, pH = -log(concentration). For strong bases, find pOH, then pH = 14 - pOH.
- Common strong acids and bases: Know the list. Any acid or base not on it is weak.
⚠ Watch out for:
Remember that H2SO4 is a strong acid in its first dissociation but the second dissociation is weak. For typical problems, treat it as strong. Also, for Group II bases, don't forget to multiply by 2. A 0.1 M Ba(OH)2 solution has [OH-] = 0.2 M, not 0.1 M.
Topic 8.3: Weak Acid and Base Equilibria
Weak acids only partially ionize. A weak acid HA in water establishes an equilibrium: HA ⇌ H+ + A-. The equilibrium constant for this reaction is Ka (the acid ionization constant). Only a small fraction of the acid molecules ionize. The vast majority remain as HA molecules.
The relationship is: Ka = [H3O+][A-] / [HA]. You can also express this as pKa = -log Ka. A larger Ka means a stronger acid (ionizes more). A smaller Ka means a weaker acid (ionizes less).
Similarly, weak bases establish an equilibrium: B + H2O ⇌ BH+ + OH-. The equilibrium constant is Kb. The percent ionization describes what fraction of the base (or acid) ionizes. It's calculated from the ratio of ionized to initial concentration, multiplied by 100%.
For a conjugate acid-base pair, Ka × Kb = Kw. This is a critical relationship. If you know Ka for an acid, you can find Kb for its conjugate base.
Key concepts to know:
- Weak acid equilibrium: Only a small percentage ionizes. Most remains un-ionized.
- Ka expression: Ka = [H3O+][A-] / [HA]. Larger Ka means stronger acid.
- pKa: pKa = -log Ka. A lower pKa means a stronger acid.
- Weak base equilibrium: Kb = [OH-][HB+] / [B]. Larger Kb means stronger base.
- Percent ionization: (ionized concentration / initial concentration) × 100%. Increases with temperature and decreases as concentration increases.
- Conjugate relationships: Ka × Kb = Kw. A strong acid has a weak conjugate base; a weak acid has a stronger conjugate base.
⚠ Watch out for:
Weak acid problems require equilibrium setup. Don't treat them like strong acids. You need Ka or Kb to find the pH. Also, remember that percent ionization is not the same as the degree of ionization. Know both definitions. And don't confuse Ka with Kb. Ka is for acids, Kb is for bases.
Topic 8.4: Acid-Base Reactions and Buffers
When you mix acids and bases, they react. A strong acid and strong base produce water: H+ + OH- → H2O. The pH of the resulting solution depends on which reactant was in excess. If the acid is in excess, the solution is acidic. If the base is in excess, the solution is basic.
When a weak acid and strong base mix, they form a buffer solution if the weak acid is in excess. A buffer contains both the weak acid (HA) and its conjugate base (A-). The buffer resists pH changes when small amounts of acid or base are added.
When a weak base and strong acid mix, a buffer forms if the weak base is in excess. The buffer contains the weak base (B) and its conjugate acid (BH+).
When two weak species mix, the reaction reaches equilibrium. Calculate which species is present in greater concentration, then use its Ka or Kb to find pH.
Key concepts to know:
- Strong acid + strong base: Produces water. pH depends on excess reactant.
- Weak acid + strong base: Forms a buffer if weak acid is in excess. Products include conjugate base and water.
- Weak base + strong acid: Forms a buffer if weak base is in excess. Products include conjugate acid and water.
- Weak acid + weak base: Reaction reaches equilibrium. Identify the major species and use Ka or Kb.
- Buffer formation: Requires a weak acid and its conjugate base (or weak base and its conjugate acid) in similar concentrations.
⚠ Watch out for:
When a weak acid and strong base mix, the quantitative reaction (stoichiometry) happens first. The strong base consumes the weak acid. Then, if excess weak acid remains, the weak acid-conjugate base equilibrium takes over. Track which species dominates after the reaction is complete.
Topic 8.5: Acid-Base Titrations
A titration is a controlled experiment where you add acid (or base) from a burette to a flask containing the other reactant. You measure the volume of titrant added when the reaction reaches the equivalence point, where moles of acid equal moles of base.
A titration curve plots pH against the volume of titrant added. The curve has a distinctive shape. At the beginning, pH changes slowly. Near the equivalence point, pH changes rapidly (this is the steep part). After the equivalence point, pH changes slowly again.
For a weak acid titrated with strong base, the equivalence point occurs at a pH above 7 (basic). This is because the conjugate base of the weak acid hydrolyzes and makes the solution basic. The half-equivalence point (halfway to the equivalence point) is useful: at this point, pH = pKa of the weak acid.
For monoprotic acids or bases, moles of titrant at equivalence point equals moles of analyte. You can use this to calculate the concentration of the unknown acid or base.
For polyprotic acids (like H2SO4 or H3PO4), the titration curve shows multiple equivalence points, one for each proton. Each equivalence point requires additional titrant. Polyprotic acid titration curves have multiple steep regions, and each half-equivalence point corresponds to a different pKa.
Key concepts to know:
- Equivalence point: Moles of acid = moles of base. Concentrations determine the actual volume at this point.
- Titration curve: Shows pH vs. volume of titrant. Steep section near equivalence point.
- Half-equivalence point: For weak acid titrations, pH = pKa at this point.
- Equivalence point pH: Strong acid + strong base = 7. Weak acid + strong base = pH > 7. Weak base + strong acid = pH < 7.
- Monoprotic calculation: moles = Molarity × Volume. Use this relationship for both acid and base.
⚠ Watch out for:
The equivalence point is not always pH = 7. It depends on the type of acid and base being titrated. For weak acid titrations, the equivalence point pH is determined by the hydrolysis of the conjugate base. Don't assume equivalence = neutral.
Topic 8.6: Molecular Structure of Acids and Bases
The strength of an acid or base relates to molecular structure. A strong acid has a very weak conjugate base, meaning the conjugate base is stabilized. Strong acids have protons on atoms that are electronegative, or they have structures where the negative charge on the conjugate base is spread out (resonance stabilization).
Carboxylic acids (R-COOH) are a common weak acid class. The OH in the carbonyl group is more acidic than an ordinary hydroxyl because the negative charge on the conjugate carboxylate anion is stabilized by resonance. The O=C-O- structure distributes the negative charge over two oxygens.
Strong bases have very weak conjugate acids. Group I and II hydroxides are strong bases. The conjugate acids (water and water, respectively) are very weak. Weak bases include ammonia (NH3) and amines, where nitrogen has a lone pair available for accepting protons.
Key concepts to know:
- Strong acids: H-X where X is very electronegative (Cl, Br, I, ClO4, NO3, SO4). Conjugate bases are stabilized.
- Carboxylic acids: Weak acids with -COOH group. Conjugate base is stabilized by resonance.
- Electronegative elements: When bonded to H, make the H more acidic. Increase acid strength.
- Resonance stabilization: Distributes negative charge over multiple atoms, stabilizing the conjugate base.
- Strong bases: Group I and II hydroxides. Weak conjugate acids.
- Weak bases: Nitrogenous bases like NH3 and carboxylate ions.
⚠ Watch out for:
Don't assume that a molecule with many OH groups is a strong base. Electronegativity, resonance, and inductive effects determine strength. NH3 is a weak base even though it's a Brønsted base. And remember that the carboxylate ion (COO-) is a weak base, while carboxylic acid is a weak acid.
Topic 8.7: pH and pKa
The relationship between pH and pKa tells you the protonation state of an acid or base. When pH < pKa, the acid form (HA) is more abundant. When pH > pKa, the base form (A-) is more abundant. When pH = pKa, the two forms are equal in concentration.
This is a powerful tool. If you know the pKa of an acid and the pH of the solution, you can predict which form dominates without calculating. For example, if pKa of acetic acid is 4.74, then at pH 3, acetic acid (HA) dominates. At pH 6, acetate ion (A-) dominates.
Acid-base indicators are molecules that have different colors in their protonated and deprotonated states. The color changes when pH crosses the pKa of the indicator. To choose an indicator for a titration, select one whose pKa is close to the equivalence point pH.
Key concepts to know:
- pH vs. pKa relationship: pH < pKa means acid form dominates. pH > pKa means base form dominates.
- At pH = pKa: Concentrations of HA and A- are equal.
- Buffer effectiveness: Buffers work best when pH is near the pKa of the weak acid.
- Indicators: Choose an indicator with pKa near the equivalence point pH for accurate titration.
⚠ Watch out for:
The pH vs. pKa relationship works for weak acids and their conjugate bases. For strong acids, pKa is very negative, so pH > pKa always holds. Don't try to use this relationship for strong acids.
Topic 8.8: Properties of Buffers
A buffer solution contains a weak acid and its conjugate base (or a weak base and its conjugate acid) in significant concentrations. The buffer resists pH changes when small amounts of acid or base are added. When you add a small amount of strong acid to a buffer, the conjugate base reacts with the added H+, removing it from solution and preventing a large pH drop. When you add a small amount of strong base, the weak acid reacts with the added OH-, removing it from solution and preventing a large pH rise.
Buffers don't prevent pH change entirely. They reduce the change. A buffer is effective when pH is within about one unit of the pKa of the weak acid. Outside this range, the buffer capacity becomes poor.
Key concepts to know:
- Buffer composition: Weak acid + conjugate base, or weak base + conjugate acid.
- Buffer mechanism: Weak acid reacts with added base; conjugate base reacts with added acid.
- pH resistance: Buffers minimize pH change when small amounts of acid or base are added.
- Effective range: Buffers work best when pH is within ±1 unit of the pKa.
⚠ Watch out for:
A buffer requires both the acid and base forms in significant concentrations. Just having a weak acid or just having the conjugate base doesn't make a buffer. Both must be present. Also, buffers don't prevent pH change, they reduce it.
Topic 8.9: Henderson-Hasselbalch Equation
The Henderson-Hasselbalch equation is a shortcut for calculating the pH of a buffer:
pH = pKa + log([A-]/[HA])
This equation comes directly from the Ka expression and is used to quickly find the pH when you know the pKa and the ratio of conjugate base to acid concentrations. If the concentrations of A- and HA are equal, the log term is zero and pH = pKa. If A- is 10 times more concentrated than HA, the pH is about 1 unit higher than pKa.
The equation is particularly useful for buffer problems. When you add acid or base to a buffer, the ratio [A-]/[HA] changes slightly, but the Henderson-Hasselbalch equation lets you calculate the new pH quickly. Adding a small amount of strong acid consumes some A-, so the ratio decreases and pH drops slightly. Adding a small amount of strong base consumes some HA, so the ratio increases and pH rises slightly.
Key concepts to know:
- Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA])
- When [A-] = [HA]: pH = pKa
- Ratio interpretation: The log term tells you how many units pH differs from pKa based on the concentration ratio.
- Buffer addition effects: Adding acid decreases [A-]/[HA] ratio; adding base increases it.
⚠ Watch out for:
The Henderson-Hasselbalch equation assumes that the pH is within the buffer's effective range and that adding acid or base doesn't dramatically change the concentrations. Also, use pKa, not Ka, in the equation. And remember the log is base 10, not natural log. Note: The CED excludes calculations of the exact pH change when strong acid or base is added to a buffer. Focus on qualitative predictions of how the ratio and pH shift.
Topic 8.10: Buffer Capacity
Buffer capacity is the amount of acid or base a buffer can neutralize while maintaining a relatively constant pH. A buffer with high capacity can neutralize a large amount of added acid or base. A buffer with low capacity can only handle small additions.
Buffer capacity depends on the concentrations of the weak acid and conjugate base. A buffer made with high concentrations of HA and A- has greater capacity than one made with low concentrations (assuming the same ratio). Increasing the concentration of the buffer components increases capacity without changing pH.
The ratio of acid to base also affects capacity asymmetrically. If you have more conjugate base than acid, the buffer has greater capacity for neutralizing added acid. If you have more acid than base, it has greater capacity for neutralizing added base.
Key concepts to know:
- Buffer capacity definition: The amount of acid or base the buffer can neutralize while maintaining relatively constant pH.
- Concentration effect: Higher concentrations of buffer components increase capacity.
- Ratio asymmetry: Buffers with excess base handle acid better; buffers with excess acid handle base better.
- pH unchanged, capacity changed: You can increase capacity by adding more buffer components without changing pH (keep the ratio constant).
⚠ Watch out for:
Buffer capacity is not just about how much acid or base you add; it's about how much the pH changes. A buffer can fail if you add too much and overwhelm its capacity. Also, the ratio of acid to base determines directional capacity, not just total capacity.
Topic 8.11: pH and Solubility
The solubility of a salt is pH sensitive when one of the ions is a weak acid, weak base, or hydroxide ion. For example, the solubility of CaCO3 increases at low pH. The carbonate ion is the conjugate base of a weak acid (HCO3-). At low pH, H+ ions protonate the CO3^2- ions, forming HCO3- and H2CO3. This removes CO3^2- from the equilibrium, allowing more CaCO3 to dissolve.
Similarly, the solubility of a metal hydroxide like Mg(OH)2 decreases at high pH. Adding OH- shifts the dissolution equilibrium left, precipitating more Mg(OH)2.
These effects can be understood qualitatively using Le Châtelier's principle. The exam does not require quantitative calculations of solubility as a function of pH. Just understand the direction of the effect and the mechanism.
Key concepts to know:
- pH-sensitive salts: Contain weak acid conjugate bases, weak bases, or hydroxide ions.
- Low pH increases solubility: H+ ions protonate bases, removing them from equilibrium.
- High pH decreases solubility: OH- ions shift dissolution equilibria left, precipitating hydroxides.
- Le Châtelier mechanism: Changes in [H+] or [OH-] shift solubility equilibria.
⚠ Watch out for:
The CED specifically excludes quantitative solubility-pH calculations. Know the qualitative effects. Understand why they happen. But don't expect to calculate exact solubility values at different pH values on the AP exam.
Study Tips for Unit 8
Unit 8 is the most mathematical unit in AP Chemistry. Here are strategies that work:
Memorize the strong acid and base lists. This is non-negotiable. If it's not on the list, it's weak. Spend 5 minutes memorizing: HCl, HBr, HI, HClO4, H2SO4, HNO3 (acids) and Group I and II hydroxides (bases).
Master the weak acid setup. Weak acid problems follow a pattern: write the equilibrium, set up Ka expression, use ICE table, solve for [H+], then pH. Practice this 10 times until it's automatic.
Understand buffers conceptually, then mathematically. First, understand that a buffer is a weak acid and its conjugate base. Then, understand that the Henderson-Hasselbalch equation is a shortcut. Finally, practice applying both the concept and the equation.
Practice titration curve interpretation. Draw the shape from memory: flat at start, then steep near equivalence point, then flat again. Label where pH = pKa, equivalence point pH, and half-equivalence point.
Connect to equilibrium. Unit 8 is just equilibrium applied to acid-base chemistry. Use Le Châtelier's principle to explain buffer behavior, indicator choice, and pH effects on solubility.
StarSpark Practice Prompts:
- "Give me a weak acid problem with Ka and initial concentration, and ask me to find pH and percent ionization"
- "Show me a buffer problem where I need to use Henderson-Hasselbalch to find pH, then predict pH after adding acid"
- "Describe a titration of weak acid with strong base, ask me to identify the equivalence point pH and the half-equivalence point"
Summary, Review Questions & Practice
You've covered all the topics in Unit 8. Before you move on, test yourself with these scenario-based questions. If you can answer them confidently, you're in great shape for this section of the exam.
Review Questions: Test Yourself
- Calculate the pH of a 0.05 M HNO3 solution and a 0.05 M NH3 solution. Explain why the pH values are different even though the concentrations are the same. (Ka for NH4+ = 5.6 × 10^-10)
- A buffer contains 0.15 M acetic acid and 0.10 M acetate ion. Calculate the pH using the Henderson-Hasselbalch equation. (pKa of acetic acid = 4.74). Then, predict what happens to the pH if you add a small amount of HCl.
- You are titrating 25.0 mL of 0.1 M weak acid (Ka = 1.8 × 10^-5) with 0.1 M NaOH. Calculate the pH at the equivalence point and explain why it's not 7.
- For the CaCO3(s) ⇌ Ca^2+(aq) + CO3^2-(aq) equilibrium, explain what happens to the solubility of CaCO3 when you add HCl. Use Le Châtelier's principle and protonation of carbonate.
- A solution contains 0.01 M weak base B with Kb = 1.5 × 10^-5. Calculate the pH, pOH, and percent ionization. Explain the relationship between Kb and the percent ionization.
Want more practice? Paste these questions into StarSpark to generate a full quiz with explanations.
Explore the Full AP Chemistry Study Guide
Unit 8 connects equilibrium (Unit 7) to electrochemistry (Unit 9). Understanding acid-base chemistry is essential for both. Buffers are equilibrium systems that resist pH change. Titration curves show you the practical application of equilibrium. And electrochemical cells depend on the oxidation and reduction of species in acidic or basic environments.
Check out the full AP Chemistry study plan to see how this unit connects to the rest of the course.
Other Unit Reviews:
- AP Chemistry Unit 1: Atomic Structure and Properties
- AP Chemistry Unit 2: Compound Structure and Properties
- AP Chemistry Unit 3: Properties of Substances and Mixtures
- AP Chemistry Unit 4: Chemical Reactions
- AP Chemistry Unit 5: Kinetics
- AP Chemistry Unit 6: Thermochemistry
- AP Chemistry Unit 7: Equilibrium
- AP Chemistry Unit 9: Thermodynamics and Electrochemistry
For official AP Chemistry resources, visit apcentral.collegeboard.org.
This review is aligned with the AP Chemistry Course and Exam Description. AP is a registered trademark of the College Board, which was not involved in the production of this guide.
