Reaction rates and equilibrium
Reaction rates
Rate of reaction
• Reaction rate = change in concentration / time
• Rate of reaction is how fast a reactant is used up or how fast a product is formed
• Rate of reaction fastest at start of reaction as all reactants are at their highest concentration
• Rate slows down as reaction proceeds at reactants concentration decreases
• Once one of reactants completely used up the concentration stops changing and rate of reaction is 0
• Collision theory: two reacting particles must collide for a reaction to occur
• An effective collision is a collision that leads to a chemical reaction
• For an effective collision the particles must collide in the correct orientation, and have sufficient kinetic energy equal to or greater than the activation energy
How to alter rate of reaction
• Concentration (or pressure when reactants are gases) - increased amount of particles in the same volume – hence particles are closer together so collide more often. Therefore in a given period of time there will be more effective collisions so increased rate of reaction
• Temperature increases kinetic energy of the particles so they collide more often at higher speeds so higher chance of effective collision so higher rate of reaction
• Increased surface area increases the amount of surface available to collide with, so more collisions which means more reactions which increases rate of reaction (more solid in contact with other reactants)
• Use a catalyst - provides an alternative pathway of lower energy
How to measure rate of reaction
• Mass balance
• Measure amount of gas given off and can see when reaction stops
• Monitor rate of formation of a precipitate (cross is concealed)
Catalysts
• A catalyst is a substance which increases the rate of a chemical reaction without being changed itself
• The catalyst may react with a reactant to form an intermediate or may provide a surface on which the reaction can take place
• At end of reaction catalyst is regenerated
• Catalyst provides a surface for the reaction to take place on
• Homogeneous catalyst - in same state as the reactants and reacts with the reactants to form an intermediate which then breaks down to give the products while also regenerating the catalyst
• Intermediate creates extra ‘bump’ in energy profile diagram (see notes)
• Examples include sulfuric acid in formation of esters
• Heterogeneous catalyst - catalyst is in a different state to the reactants (usually solids working with gases or reactants in solution)
• Reactant molecules are absorbed (weakly bonded) onto the surface of the catalyst where the reaction takes place
• After reaction, the product molecules leave surface of catalyst by desorption
Catalysis
• Approx. 90% chemical materials produced using catalyst
• Catalysts reduce temps needed industrially which cuts costs and reduces impact on environment (less fuel needed etc.)
• Economic advantages > costs needed to develop catalysts
Boltzmann distribution
• In a gas/liquid, most molecules have close to average speed with some moving very slowly and some with lots of energy moving quickly
• This spread of gases is known as Boltzmann distribution (see textbook for graph)
• No molecules have 0 energy (starts at origin)
• Area under curve = total number of molecules
• No max energy (tangent to origin)
• At higher temperature more molecules have activation energy or greater than activation so greater proportion of collision will be effective (hence curve has lower peak but higher at activation energy line)
• A catalyst provides an alternative pathway so more molecules are able to react as the activation energy is lower – the activation energy on the curve will move left so more particles have the required energy
See notes for equilibrium constant Kc
Dynamic equilibrium and Le Chatelier's principle
• In equilibrium the rate of reaction forward = the rate of reverse reaction
• The concentration of products and reactants doesn’t change
• Equilibrium systems are dynamic - both the forward and the backwards reactions are taking place, but as they are happening at an equal pace the conc doesn’t change
• For a reaction to remain in equilibrium, the system must be closed
• The position of equilibrium indicates the extent of the reaction
• Le Chatelier’s Principle - when a system in equilibrium is subjected to an external change the system readjusts itself to minimise the effect of that change
• Used to determine what effect a change will have on a mixture at equilibrium, but doesn’t explain why the change occurs, or the extent of the change
• If more products formed position of equilibrium shifts so right, if more reactants shifts to left
• If more reactants added, position shifts to the right as the extra reactants form products to decrease conc of reactants
• An increase in temperature shifts position of equilibrium in the endothermic direction, decrease toward exothermic direction
• Decrease in pressure shifts equilibrium to the side with more gaseous moles, increase towards side with fewer gaseous moles
• A catalyst doesn’t change the position of equilibrium – only speeds up rate of forward and reversible reaction equally and increases rate at which equilibrium is established