Kinetics What do you understand about rate of reaction? L.O: To know how the rate of a reaction can be affected. To be able to explain changes in rate of reaction in terms of particle collisions What do you understand about rate of reaction? What does rate mean?

What affects the rate of a reaction? Concentration of reactants. Pressure in a reaction (gaseous). Temperature of reaction. Particle size (surface area) of a solid. Intensity of radiation (specific reactions). Using a catalyst.

but, overall, is NOT consumed RATE CAN BE INCREASED BY : 1. Increasing reactant concentrations. 2. Increasing reactant pressures. 3. Increasing temperature. 4. Adding a catalyst – species which causes rate to increase but, overall, is NOT consumed 5. Using powdered solids - increases surface area Explained using : A. the COLLISION THEORY and B. the BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGIES

Collision Theory A chemical reaction will ONLY happen if the reacting particles: 1. COLLIDE and 2. Collide with a kinetic energy EQUAL TO or GREATER THAN the reaction’s ACTIVATION ENERGY, (Ea). Activation Energy (Ea) = the MINIMUM energy of collision needed for reaction to occur. and 3. Collide with PROPER ORIENTATION (bonds are correctly aligned to react)

Energy barrier separates R from P An Enthalpy Profile for a Chemical Reaction Energy barrier separates R from P Enthalpy Time Larger Ea  SLOWER reaction Ea Smaller Ea  FASTER reaction Reactants Products Remember: Ea = activation energy

An Enthalpy Profile for the reaction between chlorine atoms and ozone Point X Cl--------O-------O O The arrangement at point X only exists for a very short time. This is a simple one step reaction. Others (combustion of methane) go through a series of steps and there will be a point X for each step. Enthalpy Time Ea Cl + O3 ClO + O2

 can react when they collide Boltzmann distribution of molecular energies Explains why some molecules have energy  Ea  can react when they collide while others have energy < Ea  cannot react even if they collide. ENERGY, E  NUMBER OF MOLECULES WITH ENERGY E TOTAL AREA under curve = TOTAL number of molecules in system A Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision Ea Most molecules (this area) have insufficient energy to react

How TEMPERATURE influences rate Ea ENERGY, E  NUMBER OF MOLECULES WITH ENERGY E A N.B: Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision How TEMPERATURE influences rate Graph B represents the distribution of molecular energies at B HIGHER TEMPERATURE but SAME CONCENTRATION than for graph A  more mols with ≥ Ea  more collisions per second with ≥ Ea  faster reaction for B

Effect of temperature on distribution of energies As the temperature increases, more molecules have a higher kinetic energy. The increase in energy helps the collision to happen more effectively: More molecules have more energy, so it is MORE LIKELY that a collision will have enough energy to overcome the Ea. Higher speed of molecules means it takes less time for molecules to get close enough to react. There are MORE collisions. Raising the temperature will always increase the rate of a reaction because it makes molecular collisions MORE EFFECTIVE

Variation of boltzmann distribution with temperature ENERGY, E  NUMBER OF MOLECULES WITH ENERGY E TOTAL AREA under curve = TOTAL number of molecules in system T1 As temperature increases T1  T2  T3 T2  total number molecules unchanged but more molecules with higher energy T3  distribution shrinks vertically and moves to right

How CONCENTRATION influences rate ENERGY, E  NUMBER OF MOLECULES WITH ENERGY E Ea A N.B: Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision How CONCENTRATION influences rate Graph B represents the distribution of molecular energies at LOWER CONCENTRATION but SAME TEMPERATURE than for graph A B  fewer mols with ≥ Ea  fewer collisions per second with ≥ Ea  slower reaction for B

How PRESSURE influences rate ENERGY, E  NUMBER OF MOLECULES WITH ENERGY E Ea A N.B: Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision How PRESSURE influences rate Graph B represents the distribution of molecular energies at LOWER PRESSURE but SAME TEMPERATURE than for graph A B  fewer mols with ≥ Ea  fewer collisions per second with ≥ Ea  slower reaction for B

Variation of boltzmann distribution with concentration / pressure TOTAL AREA under curve = TOTAL number of molecules in system ENERGY, E  NUMBER OF MOLECULES WITH ENERGY E As conc / pressure decreases  total number molecules decreases  distribution shrinks vertically C / P decreasing

Catalysis Catalyst = a substance which INCREASES rate of a reaction without being consumed Causes reaction to follow a different reaction route with lower activation energy, Ea Energy Time Ea without catalyst Ea with catalyst Reactants Products

Rate Examination Questions 1(a) The collision theory states that reactions can only occur when particles collide. Give a reason why collisions between particles do not always lead to a reaction. Energy of collision may be LESS than activation energy of the reaction 1(b) Explain why, at a given temperature, hydrochloric acid reacts faster with powdered calcium carbonate than it does with lumps of calcium carbonate. Carbonate has greater surface area  more collisions per second with HCl  more collisions per second having E ≥ Ea

1(d). State and explain the effect on rate of 1(d) State and explain the effect on rate of increasing the temperature. 1(c) State and explain the effect on rate of increasing the concentration of a reactant at a fixed temperature. Increases because TOTAL number of collisions per second increases.  greater number of collisions per second having E ≥ Ea Increases % of molecules with energy greater than activation energy increases  greater number of collisions per second having E ≥ Ea

Q2. The curves labelled 1- 5 could be obtained by Q2 The curves labelled 1- 5 could be obtained by reacting excess zinc with different solutions of HCl at different temperatures and measuring the volume of hydrogen gas produced at regular time intervals 100cm3 0.1M HCl at 25C + excess Zn  curve 3 (a) 50cm3 0.2M HCl at 25C + excess Zn  curve ____ (b) 100cm3 0.1M HCl at 15C + excess Zn  curve ____ (c) 100cm3 0.2M HCl at 25C + excess Zn  curve ____ 100cm3 0.1M HCl at 25C + excess Zn + catalyst  curve ____ 2 4 1 2

Vol /cm3 t /s 1 2 3 4 5 Graphs for Q2

How A CATALYST influences rate ENERGY, E  NUMBER OF MOLECULES WITH ENERGY E Ea A N.B: Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision How A CATALYST influences rate At SAME TEMPERATURE and SAME CONCENTRATION but with a CATALYST  lower Ea Ea with catalyst  more mols with ≥ Ea  more collisions per second with ≥ Ea  faster reaction with catalyst

Using powdered solids (reactant or catalyst)  total surface area of solid increased  more collisions per second with surface  more collisions per second with E  Ea  faster reaction.

Boltzmann distribution of molecular energies Ea with catalyst ENERGY, E  NUMBER OF MOLECULES WITH ENERGY E Ea Rate  number of collisions per second with E  Ea B A without catalyst, rate  area A but with catalyst, rate  area (A+B) Catalysts influence Ea and hence rate, not overall yield!!!