CHEM UNIT 4: RATES OF REACTION

Reaction rate - change in the amount of reactants and products per unit time

Ways to follow the rate of reaction

• gas volume - collect it in a gas syringe and record how much you've got at regular time intervals (eg reaction between an acid and a carbonate in which CO2 is produced)
• loss of mass - if gas is given off, system will lose mass so you can measure this at regular time intervals
• colour change- track change using a colorimeter (eg reaction between propanone and iodine, brown colour of iodine fades)
• clock reaction - sudden colour change when product reaches a certain conc. measure time it takes for the colour change to happen - shorter the time, the faster the rate
• electrical conductivity - if no. of ions changes, so will the electrical conductivity

work out rate from a conc-time graph - gradient of the line will tell you the rate at that point on the graph. If the graph is a curve draw a tangent. - the steeper the line, the faster the rate. negative gradient if you're measuring the reactant conc. and positive if its the product conc.  

Other graphs

• absorbance vs time - bromine(aq) [orange/yellow] + methanoic acid followed using a colorimeter. percentage absorbance falls as bromine is used up - find the gradient of the graph to find the rate at that point
• gas volume vs time - decomposition of hydrogen peroxide followed by recording the volume of oxygen produced at regular time intervals - volume should increase - find the gradient

Order of reaction with respect to a reactant tells you how it's conc affects the rate

• if you increase the conc by x and the rate stays the same, the order = 0
• if you increase the conc by x and the rate increases by x, the order = 1
• if you increase the conc by x and the rate increases by x^2, the order = 2

A + B --> C + D

m + n = overall order of the reaction

k = rate constant, the bigger it is the faster the reaction. rate constant is always the same for a certain reaction at a particular temp. - if you increase the temp, the rate constant rises too

Calculate rate constant from orders and rate of reaction

reaction rates can be used to calculate the half-life of a reaction

•  zero order t(1/2) = [X] / 2k - the decomposition of X to Y and Z is zero order
• first order t(1/2) = 0.69 / k - the decomposition of N2O5 to NO2 and O2 is first order
• second order t(1/2) = 1 /k[X] - the decomposition of hydrogen iodide is second order

half lives can be used to figure out the rate equation - if the half life is contant then it is first order with respect to that reactant.

Work out orders of reactions and rate equations by experimentation

• do a series of experiments monitoring the rate of the reaction
• in each separate experiment, vary the conc of only one reactant, keep everything else the same
• plot each experiment on a conc-time graph and calculate the initial rate of reaction (gradient at time=0)
• analyse the results to see how changing the conc affects the rate, work out rate equation

Rate determining step is the slowest step in a multistep reaction

Reactants in the rate equation affect the rate

• if a reactant appears in the rate equation, it must be affecting the rate so this reactant must be in the rate determining step
• if a reactant doesn't appear in the rate equation it must be involved in a faster step
• rate determining step doesn't have to be the first step in a mechanism
• reaction mechanism can't usually be predicted from just the chemical equation

You can predict the rate equation from the rate determing step

• order of a reaction with respect to a reactant shows the no. of molecules of that reactant which are involved in the rate determining step

-

-

Use the Arrhenius equation to calculate the activation energy

• k = rate constant
• Ea = activation energy (J)
• T = temp
• R = gas constant (8.31 JK^-1mol^-1)
• A = another constant
• as the Ea gets bigger, k gets smaller . So a large Ea will mean a slow rate
• as temp. rises, k increases

you can use this equation to create an Arrhenius plot by plotting ln k against 1/T whch produces a graph with gradient of -Ea/R and rearrange to find Ea

Catalysts lower the activation energy

catalyst - increases the rate of a reaction by providing an alternative reaction pathway with a lower activation energy. The catalyst is chemically unchanged at the end of reaction.

• don't get used up in the reaction so you only need a little bit of catalyst to catalyse a huge amount of stuff. they all take part in the reaction but are remade
• have high specifity - catalysts pick and choose which reactions they catalyse

homogeneous catalysts are in the same state as the reactants

heterogenous catalysts are in a different physical state from the reactants

• solid - provide a surface for the reaction to take place on so usually a mesh or fine powder to increase surface area
• can be easily separated from the products and leftover reactants
• can be poisoned though - poison - subtance that clings to the catalyst's surface more strongly than the reactant, stopping the catalyst from getting involved in the recation it is meant to be speeding up  

enthalpy profiles and maxwell-boltzmann distribution curves show why catalysts work

• alternative pathway with lower Ea - more particles with enough energy to react when they collide so in a certain amount of time, more particles will react
• saves a lot of money in industry because they mean reactions can happen at lower temps.