Author: J R Reid Equilibrium - Introduction Equilibrium The Equilibrium Constant Factors Affecting Equilibrium

Equilibrium Reactions don’t always use up all of their reactants. This means that at the end of the reaction we don’t always have 100% products and no reactants – these are reactions in equilibrium. Some reactions take a long time to react completely – this is not the same as a reaction that is at equilibrium. Reactions can work in two directions – forward and backwards e.g. N 2 + H 2 → NH 3 NH 3 → N 2 + H 2 Sometimes a reaction will work in both directions spontaneously (without any help). When this occurs both the reactants and the products are being produced therefore there will always be levels of products and reactants present. Equilibrium equations are written with a double arrow: ⇋

The Equilibrium Constant - Intro Some reactions at equilibrium have forward and backward reactions that work at equal rates. This means that we will have an equal balance of reactants and products in reaction. Some reactions have the forward reaction that works at a much faster rate. This means that there is a greater concentration of products present. Some reactions have a backward reaction of a faster rate. This means that we have greater concentrations of reactants. To help us know what sort of reaction that we are dealing with we have an equilibrium constant – this gives us a number value based on the levels of reactants and products present at the end of a reaction

The Equilibrium Constant – Calculations (Part 1) The equilibrium constant for each reaction can be worked out using this generic chemical equation: aA + bB → cC + dD The lower case (small) letters stand for the numbers, the upper case (capital) letters stand for the chemical. Now we can use the formula above to calculate the equilibrium constant, but first a few clues: ‘K’ stands for ‘Constant’. ‘Kc’ stands for Constant based on concentration The square brackets mean ‘concentration of…’ i.e. [C] means concentration of ‘C’ The exponents match the numbers from a balanced equation

The Equilibrium Constant – Calculations (Part 2) Now we need a real life chemical equation e.g: N 2 + 3H 2 → 2NH 3 If we looked at the generic equation; aA + bB → cC + dD we see the following matches: Because there is no fourth chemical we just ignore the ‘D’ one. Note that the numbers before each chemical matches the lower case letter. When we substitute into our formula this: Becomes this

The Equilibrium Constant – Calculations (Part 3) Now that we know what letter match which part of the equation we can plug in all the concentrations into the formula e.g. [NH 3 ] = the concentration of ammonia gas at equilibrium (in moles per litre) [N 2 ] = the concentration of nitrogen gas at equilibrium [H 2 ] the concentration of hydrogen gas at equilibrium

Factors Affecting Equilibrium The concentrations used to calculate the equilibrium constant change at different temperatures, pressures, concentrations etc. So when ever we calculate or use a constant we must make sure that we remember the standard conditions: The temperature must be 25ºC The air pressure must by 1 atmosphere 1 molL -1 concentration of reactant must be used If we change any of these factors the equilibrium will shift to favour the reactants or products (more of one of them will be produced)

Le Chatelier’s Principle A chemist called Le Chatelier developed a law that predicted if there would be more reactants or products porduced when you changed the environmental conditions. It went like this: If a system in equilibrium is subjected to a stress then a change occurs to try to overcome that stress In other words, the chemicals don’t like change, and if change occurs they try to remove the change Examples: If we add heat, they will try to remove that heat (through an endothermic reaction) If we add more of a chemical they will try to remove that chemical If we add more (air) pressure they will remove gases to remove the pressure

Temperature and Equilibrium Exothermic reactions make heat as a product of a reaction. It could be written like this: H 2 + Cl 2 ⇋ 2HCl + heat Therefore if we add heat to the reaction the equilibrium will move in the direction that uses up the heat (in this case towards the left) Note: Exothermic reactions have a negative change in enthalpy (- Δ H) Endothermic reactions have a positive change in enthalpy (+ Δ H)

Pressure and Equilibrium In the following reaction we have gases being made on both sides of the reaction: N 2(g) + 3H 2(g) ⇋ 2NH 3(g) If we increased the pressure both sided get squeezed. The chemicals will ‘move’ to the side that has less gas to try to remove the pressure. In the reaction above the left hand side has a total of 4 moles of gas (1 nitrogen and 3 hydrogens). The right hand side has a total of 2 moles of gas (2 ammonia). So if we increased the pressure then the N 2 and H 2 would be used up, and we increase our NH 3 production If we decreased the pressure we will use up more of the ammonia and more N 2 and H 2 will be produced

Concentration and Equilibrium Adding more of a certain chemical will cause the reaction to try to use it up. In the following reaction if we increased the reactants the reaction would try to use them up by turning them into products. N 2(g) + 3H 2(g) ⇋ 2NH 3(g) Adding more N 2 will cause more NH 3 to be made. It will also cause more H 2 to be used up – why? Adding more H 2 will cause more NH 3 to be made. It will also cause more N 2 to be used up – why? Adding more NH 3 will cause more N 2 and H 2

Catalysts and Equilibrium Catalysts have no affect on equilibrium. They cause the reaction to run faster, but they do not change the concentrations of the products or reactants at the end of the reaction.

Exam Practice Can’t see the exam paper below? Go to the NCEA website and search for 90310NCEA Have a go at Questions: Two Four

Exam Practice Can’t see the exam paper below? Go to the NCEA website and search for 90310NCEA Have a go at Questions: Five Six

Exam Practice Can’t see the exam paper below? Go to the NCEA website and search for 90310NCEA Have a go at Questions: Five Six a)