Acid-Base Reactions to Know for AP Chemistry

Acid-base reactions are fundamental in chemistry, influencing everything from biological processes to industrial applications. Understanding different theories, pH calculations, and the behavior of acids and bases helps us predict reactions and their outcomes in various contexts.

1. Arrhenius theory of acids and bases

   • Acids produce hydrogen ions (H⁺) in aqueous solution.
   • Bases produce hydroxide ions (OH⁻) in aqueous solution.
   • This theory is limited to aqueous solutions and does not account for acid-base reactions in non-aqueous solvents.

2. Brønsted-Lowry theory of acids and bases

   • Acids are proton donors, while bases are proton acceptors.
   • This theory expands the definition of acids and bases beyond aqueous solutions.
   • It introduces the concept of conjugate acid-base pairs, where the acid becomes a base after donating a proton.

3. Lewis theory of acids and bases

   • Acids are electron pair acceptors, and bases are electron pair donors.
   • This theory encompasses a wider range of chemical reactions, including those that do not involve protons.
   • It highlights the role of electron pairs in acid-base interactions.

4. pH scale and calculations

   • The pH scale measures the acidity or basicity of a solution, ranging from 0 (strongly acidic) to 14 (strongly basic).
   • pH is calculated using the formula: pH = -log[H⁺].
   • A change of one pH unit represents a tenfold change in hydrogen ion concentration.

5. Strong acids and strong bases

   • Strong acids completely dissociate in water, releasing all their H⁺ ions (e.g., HCl, H₂SO₄).
   • Strong bases completely dissociate in water, releasing all their OH⁻ ions (e.g., NaOH, KOH).
   • They have a pH close to 0 for strong acids and close to 14 for strong bases.

6. Weak acids and weak bases

   • Weak acids partially dissociate in water, establishing an equilibrium (e.g., acetic acid).
   • Weak bases partially accept protons, also establishing an equilibrium (e.g., ammonia).
   • They have a pH that is not as extreme as strong acids or bases, typically between 3 and 11.

7. Conjugate acid-base pairs

   • A conjugate acid is formed when a base accepts a proton, while a conjugate base is formed when an acid donates a proton.
   • The strength of an acid and its conjugate base are inversely related; strong acids have weak conjugate bases.
   • Understanding these pairs is crucial for predicting the direction of acid-base reactions.

8. Acid-base titrations

   • A titration is a quantitative method to determine the concentration of an acid or base by neutralizing it with a base or acid of known concentration.
   • The endpoint is often indicated by a color change, which can be monitored using indicators.
   • The titration curve illustrates the pH change as the titrant is added, showing the equivalence point.

9. Buffer solutions

   • Buffers are solutions that resist changes in pH when small amounts of acid or base are added.
   • They typically consist of a weak acid and its conjugate base or a weak base and its conjugate acid.
   • Buffers are essential in biological systems and many chemical processes to maintain stable pH levels.

10. Acid-base indicators

    • Indicators are substances that change color at a specific pH range, signaling the acidity or basicity of a solution.
    • Common indicators include phenolphthalein (colorless in acid, pink in base) and litmus (red in acid, blue in base).
    • The choice of indicator depends on the expected pH range of the titration.

11. Neutralization reactions

    • Neutralization occurs when an acid reacts with a base to produce water and a salt.
    • The general reaction can be represented as: Acid + Base → Salt + Water.
    • These reactions typically result in a solution with a pH close to 7, depending on the strength of the acid and base.

12. Hydrolysis of salts

    • Hydrolysis occurs when a salt dissolves in water and reacts with water to produce an acidic or basic solution.
    • The nature of the salt (derived from a strong acid and weak base or vice versa) determines whether the solution will be acidic or basic.
    • Understanding hydrolysis is important for predicting the pH of salt solutions.

13. Ka and Kb (acid and base dissociation constants)

    • Ka is the equilibrium constant for the dissociation of an acid, while Kb is for the dissociation of a base.
    • These constants provide a measure of the strength of acids and bases; larger values indicate stronger acids or bases.
    • The relationship between Ka and Kb for a conjugate acid-base pair is given by: Ka × Kb = Kw (where Kw is the ion product of water).

14. Common ion effect

    • The common ion effect describes the shift in equilibrium that occurs when a common ion is added to a solution.
    • This effect can suppress the dissociation of weak acids or bases, leading to a decrease in their ionization.
    • It is important in buffer solutions and in predicting the behavior of salts in solution.

15. Le Châtelier's principle in acid-base equilibria

    • Le Châtelier's principle states that if a system at equilibrium is disturbed, the system will shift to counteract the disturbance.
    • In acid-base equilibria, adding or removing reactants or products will shift the equilibrium position.
    • This principle helps predict the direction of shifts in acid-base reactions when conditions change.