# Rate of Reaction University of Lincoln presentation

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Rate of Reaction University of Lincoln presentation

Why the difference? Is it the enthalpy change (Heat of combustion) ?
Paraffin wax 42 MJ kg-1 Petrol 45 MJ kg-1 Is it the temperature? Yellow/white – 1300oC Pale orange/yellow – 1100oC What is it then? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Approximately how long will a 2 litre pool of petrol burn for?
Important values: Petrol density = 0.8 kg litre-1 Heat of combustion is 45 MJ kg-1 2 litres of petrol has a mass of 1.6 kg (from the density) Total energy available from 1.6 kg petrol = 1.6 kg x 45 MJ kg-1 = 72 MJ 2 litre petrol pool is a 1 MW fire (this is a measured value) 1 MW = 1 MJ s-1 so at this rate it would take 72 s This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

How do ignitable liquids burn?

2 litre petrol bomb takes about 10s to burn
2 litre petrol bomb takes about 10s to burn. What is the rate of heat release? 72 MJ in 10 s = 7.2 MW 2 litre petrol fully evaporated takes about 1 s to burn. What is the rate of heat release? 72 MJ in 1s = 72 MW Conclusion: Same total energy available but released at a faster rate This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

How long to burn a 1.6 kg candle?
1.6 kg paraffin wax at 42 MJ kg-1 can release 67.2 MJ Candle flame has a heat release rate of 80 W (80 Js-1) This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

A candle bomb? NASA are researching the paraffin rocket!!
How can this work? Increase rate of combustion Increase concentration of the oxidant; use 100% oxygen Paraffin as small liquid droplets Study of the rates of reaction - kinetics This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Factors affecting the rate of a chemical reaction

Measuring Reaction Rate
Use a characteristic of the products or the reactants that can be used as a measure of amount. Volume of gas Change in mass Absorption of light rate of decrease of reactant or rate of increase of a product This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Calculating Rate of Reaction
The gradient of tangent to the curve is the rate of reaction What happens to the reaction rate with time? Concentration = 0.3 mol dm-3 s-1 Rate = gradient = mol dm-3 s-1 Concentration = 0.1 mol dm-3 s-1 Rate = gradient = mol dm-3 s-1 This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

A Mathematical Relationship
Select two other points on the curve and calculate the rate of reaction at that concentration of H2O2 Plot a graph of Rate of Reaction as a function of H2O2 concentration This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

What does the graph show?
Graph is a straight line through the origin The two variables have a linear mathematical relationship We can say: Rate of Reaction is directly proportional to Hydrogen Peroxide concentration Easy to predict what happens to reaction when [H2O2] is changed [H2O2] x2 Rate x 2 First Order with respect to H2O2 k is the rate constant; first order reaction has units of s-1 when the rate of reaction is measured in mol dm-3 s-1. Show this by rearranging the rate equation and why are the units of rate important. This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Rate of reaction can be measured from the rate that oxygen gas is produced.

Vary the starting concentration and measure the initial rate
[H2O2]= 0.40 mol dm-3 [H2O2]= 0.32 mol dm-3 [H2O2]= 0.24 mol dm-3 [H2O2]= 0.16 mol dm-3 [H2O2]= 0.08 mol dm-3 This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Initial Rate can be measured

Plot Initial Rate as a function of starting concentration

Summary Decomposition of H2O2 can be followed by measuring the decrease in H2O2 concentration or the volume of O2 evolved. Rate of reaction can be calculated from the progress curve at different times or initial rate measurements. Plots of rate as a function of reagent concentration can be used to determine the mathematical relationship Order of reaction can be determined Rate equation can be written This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

For you to do: Initial rate data
[H2O2]/mol dm-3 Rate/cm3 O2 s-1 0.08 0.1 0.16 0.215 0.24 0.32 0.41 0.4 0.51 Determine the order of reaction with respect to hydrogen peroxide and calculate the value of the rate constant. This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

General Rate equations
Zero order Units of k ? First order Units of k ? Second order Units of k ? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Further Analysis of data
Logarithms can be very useful Plot of log rate as a function of log concentration (p439 Housecroft) Gradient is n; Intercept is log k Use this method on the initial rate data in slide 21 to determine order and the value of k This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Half-life Time taken for the concentration of reactant A at time t, [A]t to fall to half its value. A constant half-life for a first order reaction Progress curve and measure t½ at several different points. This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Constant half-life 27s 27s 26s
Going from 200 x 10-5 mol to 100 x 10-5 mol takes 27s Going from 100 x 10-5 mol to 50 x 10-5 mol takes 27s Going from 50 x 10-5 mol to 25 x 10-5 mol takes 26s 27s 27s 26s This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Using Luminol to detect blood stains Exponential decay curve

Reactions with more than one reactant
A + B → products e.g. C12H22O11 + H2O → C6H12O6 + C6H12O6 sucrose glucose fructose First order with respect to each reactant Second order reaction (sum of orders in rate equation) This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Determining order and rate equations
Difficulties with more than one reactant? Experimental Design Principle Vary one concentration and keep other(s) constant while measuring rate. Initial rate method Isolation method This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

An Example Reaction peroxodisulfate (VI) and iodide ions
Task: Determine the rate equation and a value for k Design the experiment 1. initial rate method (vary each concentration) 2. Plot a graph of log rate as a function of log initial concentration for each reactant. Gradient of each line is order of reaction for each reactant. 3. k is determined by rearranging the rate equation. This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Iodine clock data from experiment

Collision theory Molecules have to collide if they are to react – increasing frequency of collisions? Increasing concentration increases the frequency of collisions Increasing pressure increases frequency of collisions Increasing temperature increases frequency of collision But not just about rate of collisions – how do we explain slow reactions? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Activation Energy (Enthalpy) Ea
Energy of the collision must be above a certain value for reactants to react Why? Energy is needed to break bonds (remember bond enthalpies) This then creates reactive species to make new bonds The minimum energy required for a collision to result in chemical reaction is Ea This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Only those molecules with sufficient energy can react

Increasing Temperature increases Rate of Reaction
Number of molecules with kinetic energy E Kinetic energy (E) Activation enthalpy Ea =50kJ mol-1 Number of molecules with energy greater than 50kJ mol-1 at 300 K Number of molecules with energy greater than 50kJ mol-1 at 310 K 300 K 310 K This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Back to petrol Petrol vapour reacts with oxygen (air)
But not spontaneous at room temperature Needs ignition. What does ignition do? Provides energy to break bonds (endothermic) Creates reactive species (free radicals) Self-sustaining (can remove ignition source and it carries on). Why???? Energy released from the reaction breaks more bonds and the reaction continues This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Combining Activation Energy and enthalpy
Both can be shown on an enthalpy level diagram Exothermic reaction Positive enthalpy Add Ea to the diagram A+B ΔH Negative enthalpy C+D Draw a diagram for an endothermic reaction Reaction coordinate This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Rate equations and Temperature
The Arrhenius equation k is the rate constant; A is the pre-exponential factor; Ea is the activation energy; R is the molar gas constant (8.314 J mol-1 K-1); T is the absolute temperature (Kelvin). How does it work? It might be easier to do this increase temperature increase k increase rate decrease Ea increase k increase rate This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

The well known ‘rule of thumb’
Reaction rate doubles if temperature is increased by 10 oC Temp/K Ea/kJ mol-1 A/L mol-1 s-1 k/L mol-1 s-1 313 54 8.7 x 106 8.5 x 10-3 323 1.6 x 10-2 Check the values of k by calculating them from the Arrhenius equation using the other values in the table Calculate k at 333 K. What is happening to the value of k? How will this affect the rate of this reaction? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

An experiment to determine Ea
Determine order and rate equation for the reaction Measure the rate of reaction at different temperatures keeping the initial concentrations the same Calculate k at the different temperatures Plot lnk against 1/T: gradient = -EA/R; intercept = lnA This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Calculating Ea Temperature/K k/dm3 mol-1 s-1 296 2.9 x 10-3 302
Use the data below to calculate a value for the activation energy for this reaction Temperature/K k/dm3 mol-1 s-1 296 2.9 x 10-3 302 4.2 x 10-3 313 8.3 x 10-3 323 1.9 x 10-2 This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

How do we explain catalysis?

What are catalysts? Definition and some examples; reactions and catalysts Hydrogen peroxide , metals and natural substances Enzymes Gases on metal surfaces What is a different reaction route? This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

A catalyst provides an alternative path for the reaction with a lower activation enthalpy
Uncatalysed reaction Activation enthalpy of uncatalysed reaction Catalysed reaction Enthalpy Activation enthalpy of catalysed reaction Reactants Products Progress of reaction This work is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 2.0 UK: England & Wales License

Acknowledgements JISC HEA