WJEC Chemistry: 1.7- Equilibria and Acid-Base Reactions

Dynamic Equilibrium:

• DEFINITION: "Dynamic Equilibria is when the forward and reverse reactions occur at the same rate. "

• CONDITIONS: - The system is closed (Reactants and Products cannot escape).                                                                            - The reverse reaction is significant.

Le Chatelier's Principle:

• DEFINITION: "If a system at equilibrium is subjected to a change then the position of equilibrium will shift to minimise that change."

• The variables that can result in the position of the equilibrium changing are:                                                               - The addition or removal of one of the reactants or products.                                                                                   - A change in pressure.                                                                                                                                               - A change in temperature.                                                                                                                                         - The addition of a catalyst.

Change in Concentration:

• When the concentration of one of the species is increased, the system moves to remove the species. When the concentration is decreased, the system moves to replace the species.

Change in Pressure:

• When the pressure is increased, the system will move to the side with the fewest moles in order to decrease the pressure. When the pressure is decreased, the system will move to the side with more moles in order to increase the pressure.

Change in Temperature:

• When the temperature is increased, the system will move in the endothermic direction in order to decrease the temperature. When the temperature is decreased, the system moves in the exothermic direction to increase the temperature.

Addition of a Catalyst:

• When a catalyst is added, the rates of the forward and reverse reactions are increased by the same amount and the position of the equilibrium is unaffected.

Changes in Pressure, Concentration and Addition of a Catalyst do not affect the equilibrium constant. However, if the reaction is exothermic, then an increase in Temperature will cause the value of Kc to decrease, and if the reaction is endothermic, then an increase in Temperature will cause the value of Kc to increase.

1. Equilibrium Constant Expressions

• The position of equilibrium in a reversible reaction can be shown by combining the equilibrium concentrations to give a value for an equilibrium constant (Kc). 

• The products are put in the numerator and the reactants in the denominator. The concentrations are raised to powers corresponding to the mole ratio in the equation.

2. Units for Equilibrium Constants (Kc)

• The units of the equilibrium constant vary, depending on the number of species involved. The units can be deduced by multiplying out the units of concentration and cancelling as appropriate.

Acids:

• DEFINITION: "An acid is a substance that behaves as a proton (H+) donor."

• STRONG AND WEAK: The difference between strong and weak acids is that a strong acid fully dissociates in aqueous solution, where as a weak acid only partially dissociates in aqueous solution.                                               - Examples of Strong Acids: HCl, HNO3, H2SO4

• CONCENTRATED AND DILUTE: A concentrated acid consists of a large quantity of acid and a small quantity of water. A dilute acid contains a large quantity of water.

• The acidity of a solution depends on both the concentration of the solution and the strength of the acid.

Bases:

• DEFINITION: "A base is a substance that behaves as a proton (H+) acceptor."                                                 

• Examples of Strong and Weak Bases: NaOH, KOH (Strong), NH3 (Weak).

The acidity or alkilinity of a solution is often measured by its pH

pH Calculations:

• pH= -log[H+]  --->  where "[H+]" is the concentration of H+ in moldm-3

• Example: "What is the pH of a sample of rain water with a H+(aq) concentration of 3.9x10-6 mol dm-3?                     pH= -log(3.9x10-6)                                                                                                                                                           = 5.4

• To calculate the concentration of the H+ ions given the pH take the antilogarithm of both sides ---> [H+]= 10-pH
  

Neutralisation:

• Bases react with acids in neutralisation reactions to form a salt and water.                                                                   - Example: HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)                                                                                                                                                                                                                                                     
• A salt is a compound that forms when a metal ion replaces a hydrogen ion in a acid.

An acid-base titration is a type of volumetric analysis where the volume of one solution that reacts with a known volume of another solution is measured.

• Standard Solution: "A solution where the concentration is accurately known."

• Equipment: A burette containing one solution (i.e. an acid), A conical flask containing the other solution (i.e. a base). A pipette to accurately transfer the other solution to the conical flask. An indicator to show when the reaction is completed.

Double Titrations:

• If a solution contains a mixture of two bases which are of different strengths, one titration can be performed, but in two stages, using two different indicators, one added at each stage to calculate the concentrations of both bases, (as each indicator changes colour at a different pH).

Back Titrations:

• A known excess of one reagant A reacts with an unkown amount of reagant B. At the end of the reaction, the amount of reagant A that remains is found by titration. A simple calculation gives the amount of reagant A that has been used and the amount of reagant B that has reacted.