AP Chemistry College Board

This subject is broken down into 91 topics in 9 modules:

  1. Acids and Bases 10 topics
  2. Applications of Thermodynamics 10 topics
  3. Atomic Structure and Properties 8 topics
  4. Chemical Reactions 9 topics
  5. Equilibrium 14 topics
  6. Intermolecular Forces and Properties 13 topics
  7. Kinetics 11 topics
  8. Molecular and Ionic Compound Structures and Properties 7 topics
  9. Thermodynamics 9 topics
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This page was last modified on 28 September 2024.

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Chemistry

Acids and Bases

Acid-Base Reactions and Buffers

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Acid-Base Reactions and Buffers

Acid-Base Reactions

  • An acid-base reaction is also known as a neutralisation reaction. It involves the direct transfer of a proton (H+) from an acid to a base. The products usually include water and a salt.
  • Acids are proton donors. They typically taste sour, react with metals to produce hydrogen gas, and turn blue litmus paper red. Common examples include hydrochloric acid (HCl) and sulfuric acid (H2SO4).
  • Bases are proton acceptors. They taste bitter, feel slippery to the touch and turn red litmus paper blue. Common examples include sodium hydroxide (NaOH) and ammonia (NH3).
  • The reaction of a strong acid with a strong base results in a completely neutral solution with a pH of 7. For example: HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l).
  • However, the reaction of a weak acid with a strong base, or a strong acid with a weak base, does not necessarily give a neutral solution. It depends on the relative strengths of the acid and base.

pH Scale and the Role of Water

  • The pH scale is used to measure how acidic or basic a solution is. It ranges from 0 to 14, with 7 being neutral. Values below 7 indicate acidity, and values above 7 indicate basicity.
  • Water plays a critical role in pH. It can act as both an acid and a base. When pure water self-ionises, it produces equal amounts of H+ (hydrogen ion) and OH- (hydroxide ion). This state of balance means that pure water has a neutral pH of 7.

Buffer Solutions

  • A buffer solution is one that resists changing its pH, even when small amounts of acid or base are added. This is accomplished by the presence of significant amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid.
  • In a buffer, the weak acid reacts with any added base and the conjugate base reacts with any added acid, helping to keep the pH more or less constant.
  • An example of an acid buffer system consists of ethanoic acid (CH3COOH) and its conjugate base, the ethanoate ion (CH3COO-).
  • Buffer solutions are important in many biological systems. For example, the human blood stream is buffered to maintain a pH of approximately 7.4.

Remember: Accurately determining pH levels and understanding how to create and manipulate buffer solutions is essential in many scientific and industrial applications.

Course material for Chemistry, module Acids and Bases, topic Acid-Base Reactions and Buffers

Chemistry

Equilibrium

Properties of the Equilibrium Constant

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Properties of the Equilibrium Constant

Chemical Equilibrium

  • Chemical equilibrium is achieved in a reaction when the rate of the forward reaction is equal to the rate of the reverse reaction.
  • At equilibrium, the concentrations of reactants and products remain constant over time, but they aren't necessarily equal.

Equilibrium Constant (K)

  • The equilibrium constant (K) is a measure of the ratio of product concentrations to reactant concentrations when a chemical reaction reaches equilibrium.
  • This constant is determined using the expression: K = [C]^c [D]^d / [A]^a [B]^b. Here, [A], [B], [C] and [D] represent molar concentrations of reactants (A and B) and products (C and D); while a, b, c, and d are their respective coefficients in the balanced chemical equation.
  • K is calculated at a specific temperature and remains constant unless the temperature changes.

Interpreting Equilibrium Constant

  • If K > 1, the reaction favours formation of products at equilibrium. This means the equilibrium lies to the right.
  • If K < 1, the reaction favours formation of reactants at equilibrium, i.e., the equilibrium lies to the left.
  • If K = 1, the amounts of products and reactants at equilibrium are approximately equal.

Relating K to ΔG

  • The equilibrium constant is related to the standard free-energy change (ΔG°) according to the following equation: ΔG° = -RT lnK. Here, R refers to the gas constant and T stands for temperature in Kelvin.
  • A large positive ΔG° (ΔG° >> 0) means equilibrium constant K is very small (K << 1), showing reaction favours the reactants. Conversely, a large negative ΔG° (ΔG° << 0) corresponds to large K (K >> 1), indicating reaction favours the products.
  • At equilibrium, ΔG° = 0.

Effect of Changing Conditions on K

  • Changes in concentration or pressure do not change the value of K. They may shift the position of equilibrium but K will remain the same.
  • Only changes in temperature can alter the value of the equilibrium constant K. If the forward reaction is exothermic (releases heat), increasing temperature will cause K to decrease. If the forward reaction is endothermic (absorbs heat), increasing temperature will cause K to increase.

Significance of Equilibrium Constant

  • Knowledge of the equilibrium constant helps scientists predict the direction of the reaction in given conditions, allowing us to understand, predict and control chemical reactions in laboratories and industrial processes.

Course material for Chemistry, module Equilibrium, topic Properties of the Equilibrium Constant

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