Equilibrium Rapid Review for AP Chemistry
For a more thorough review, refer to these concepts:
- Le Chatelier's Principle for AP Chemistry
- Acid–Base Equilibrium for AP Chemistry
- Ka, Kw, Kb - The Acid, Water, and Base Dissociation Constant for AP Chemistry
- Buffers for AP Chemistry
- Titration Equilibria for AP Chemistry
- Solubility Equilibria for AP Chemistry
- A chemical equilibrium is established when two exactly opposite reactions occur in the same container at the same time and with the same rates of reaction.
- At equilibrium the concentrations of the chemical species become constant, but not necessarily equal.
- For the reaction aA + bB cC + dD, the equilibrium constant expression would be: . Know how to apply this equation. c
- Le Châtelier's principle says that if an equilibrium system is stressed, it will reestablish equilibrium by shifting the reactions involved. A change in concentration of a species will cause the equilibrium to shift to reverse that change. A change in pressure or temperature will cause the equilibrium to shift to reverse that change.
- Strong acids completely dissociate in water, whereas weak acids only partially dissociate.
- Weak acids and bases establish an equilibrium system.
- Under the Bronsted–Lowry acid–base theory, acids are proton (H+) donors and bases are proton acceptors.
- Conjugate acid–base pairs differ only in a single H+; the one that has the extra H+ is the acid.
- The equilibrium for a weak acid is described by Ka, the acid dissociation constant. It has the form: . Know how to apply this equation.
- Most times the equilibrium concentration of the weak acid, [HA], can be approximated by the initial molarity of the weak acid.
- Knowing Ka and the initial concentration of the weak acid allows the calculation of the [H+].
- Water is an amphoteric substance, acting either as an acid or a base.
- The product of the [H+] and [OH–] in a solution or in pure water is a constant, Kw, called the water dissociation constant, 1.0 × 10–14. Kw = [H+] [OH–] = 1.0 × 10–14 at 25°C. Know how to apply this equation.
- The pH is a measure of the acidity of a solution. pH = –log[H+]. Know how to apply this equation and estimate the pH from the [H+].
- On the pH scale 7 is neutral; pH > 7 is basic; and pH < 7 is acidic.
- pH + pOH = pKw = 14.00. Know how to apply this equation.
- Kb is the ionization constant for a weak base. . Know how to apply this equation.
- Ka × Kb = Kw for conjugate acid–base pairs. Know how to apply this equation.
- Buffers are solutions that resist a change in pH by neutralizing either an added acid or an added base.
- The Henderson–Hasselbalch equation allows the calculation of the pH of a buffer solution: . Know how to apply this equation.
- The buffer capacity is a quantitative measure of the ability of a buffer to resist a change in pH. The more concentrated the acid–base components of the buffer, the higher its buffer capacity.
- A titration is a laboratory technique to determine the concentration of an acid or base solution.
- An acid–base indicator is used in a titration and changes color in the presence of an acid or base.
- The equivalence point or endpoint of a titration is the point at which an equivalent amount of acid or base has been added to the base or acid being neutralized.
- Know how to determine the pH at any point of an acid–base titration.
- The solubility product constant, Ksp, is the equilibrium constant expression for sparingly soluble salts. It is the product of the ionic concentration of the ions, each raised to the power of the coefficient of the balanced chemical equation.
- Know how to apply ion-products and Ksp values to predict precipitation.
- Formation constants describe complex ion equilibria.
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