M3 B6 Rates of Reactions

The rate of a reaction is the change in concentration of a reactant or a product in a given time. (how fast a reaction is).

• Measure the volume of gas produced (product) at regular time intervals using gas collection (gas syringe or measuring cylinder). • Measuring loss in mass using a balance. e.g- reaction of calcium carbonate with dilute hydrochloric acid

• Temperature • Surface area • Pressure • Concentration • Adding a catalyst

Particles must collide with enough energy to react (aka a successful collision)- this is the activation energy. The molecules must also collide in the correct orientation.

• The rate of the reaction increases • because the number of particles in a given volume increases. • So the number of collision increases- collisions happen more frequently • Therefore the number of successful collisions increases.

Concentration

• Increasing the pressure increases the rate of a reaction • because particles are pushed closer together • This means there are more frequent collisions • Number of successful collisions also increases.

• Increasing temperature increases the rate • because molecules are moving faster as they have more kinetic energy • The number of collisions increases. • The number of successful collisions increases as more molecules have an energy greater than or equal to the activation energy.

• The number of particles (distribution) in a sample of gas with a particular energy.

The area under the curve represents the total number of particles in the reaction.

The higher temperature (2nd) curve has a maximum (peak) that is lower and to the right of the lower temperature (1st) curve.

A catalyst is a substance that increases the rate of a reaction without being used up by the overall equation. It is regenerated.

A catalyst provides an alternate route for the reaction with a lower activation energy.

A catalyst increases the rate of a reaction because 1. A lower activation energy means that more molecules have an energy equal to or greater than the activation energy. 2. So the number of successful collisions increases as more frequent collisions with energy > or = to activation energy.

Enthalpy profile diagram: y-axis = enthalpy (kJ mol⁻¹) x-axis = reaction progress (not time) Shows energy change (ΔH) + activation energy Reactants and products Boltzmann distribution: x-axis = energy y-axis = number of molecules skewed curve (starts high, tails off) Shows distribution activation energy as a line.

A catalyst that is in the same state as the reactants (solid, liquid or gas).

An intermediate (the intermediate then breaks down to give the product and regenerates the catalyst)

A species that is formed in one step of a reaction and used up in a later step, so it does not appear in the overall equation.

• Making esters- carboxylic acid and alcohol are liquids- so is the catalyst H2SO4 • Ozone depletion- O3 + O = 2O2 the chlorine radicals are a gas

A catalyst which is in a different state to the reactants.

• ADsorption- Gas molecules adsorb onto the surface of the metal catalyst (form a weak bond) • Weakening of bonds- The bonds within the reactant molecule are weakened by the interaction with the surface so the reaction happens faster. • Desorption- The product molecules are desorbed from the surface- into the atmosphere.

• The Haber process- the reactants and products are gases- iron (solid) is the catalyst • Hydrogenation of unsaturated oils (alkenes)- Nickel (solid) is used as a catalyst for the gas reactants and products • Catalytic convertors- CO and NO are turned into CO2 and N by passing over a solid platinum, palladium, or rhodium catalyst.

• Lowering the activation energy means that the reaction can run at a lower temp- REDUCING COSTS. • This also means that LESS FOSSIL FUELS are being burned- LESS CO2 EMITTED. • Since an alternative pathway is used, they may have a BETTER ATOM ECONOMY, or LESS HAZARDOUS WASTE produced. • They can also BE REUSED.

Rate= Change in volume of gas produced (y axis) divided by change in time (x axis)