1
Q
Collision theory
A
Based on idea that for a chemical reaction to occur between two or more reactant particles, they must collide
2
Q
Requirements for a chemical reaction
A
- reactant particles must collide with correct orientation
- reactant particles must collide w/ sufficient energy to overcome energy barrier for reaction (activation energy)
3
Q
Unsuccessful collision
A
When the requirements for a successful collision aren’t met, reactant particles may simply bounce apart without reacting
- also happens when colliding particles don’t have energy equal to or greater than activation energy of reaction
4
Q
For a chemical reaction to take place, the following conditions must be met
A
- Reactant particles must collide
- Reactant particles must collide with the correct orientation
- Reactant particles must collide with energy equal to or greater than activation energy
5
Q
Activation energy of a reaction
A
Minimum amount of kinetic energy that colliding particles must have for a chemical reaction to occur
- amount of kinetic energy needed to overcome the energy barrier between reactants and products
- difference in energy between the reactants and the transition state of the reaction
6
Q
Transition state
A
Highest energy state on a reaction coordinate
- indicates a point at which new bonds are being formed at the same time as old bonds are being broken
7
Q
Chemical reactions and activation energy
A
- faster reactions have lower activation energies
- slower reactions have higher activation energies
8
Q
Maxwell-Boltzmann distribution
A
Theory that, in an ideal gas, the kinetic energy of the molecules is spread over a range of values
9
Q
Maxwell-Boltzmann curve
A
The total area beneath the curve is equal to total no. of particles in the sample
- area under any region of the curve is directly proportional to no. of molecules having a value of kinetic energy in that range
- shaded region = no. of particles that have energy equal to, or greater than, the activation energy (Ea)
- shape of distribution changes as temperature is increased
- when temperature is increased, a greater proportion of particles have energy equal to, or greater than, Ea
10
Q
Maxwell-Boltzmann curve and changes in temperature
A
Temperature of a gas sample is increased:
- peak of distribution curve shifts to right, increase in most likely value for kinetic energy of particles. At higher temp., average KE of particles increases
- curve flattens, becoming broader so total area under it remains constant (there are still same no. of particles in sample)
- increase in area under curve to right of Ea value
NB/ at higher temp., a greater proportion of particles will have energy equal to or greater than Ea
11
Q
Temperature
A
Measure of average kinetic energy of particles in a substance
12
Q
Why does increasing temp. increase rate of reaction?
A
- when temp. increases, average kinetic energy of particles in a substance increases
- particles start moving faster, collide more frequently and more energetically
- for many reactions, increase of 10 degrees C, doubles rate of reaction- because more particles have energy equal to, or greater than, Ea
13
Q
Effect of a catalyst
A
Provides a reaction pathway that requires a lower Ea
- hence, a greater proportion of reactant particles will have energy equal to, or greater than, Ea
14
Q
Kelvin scale
A
An absolute temperature scale
- an absolute unit of measurement
- doubling temp. on this scale = average KE of particles in sample is also doubled
15
Q
Average KE and Kelvin scale
A
The absolute temperature in K is directly proportional to average kinetic energy of particles in a sample
16
Q
Absolute 0
A
Lowest possible temperature on Kelvin scale- it is the origin of x-axis of distribution
At this temp. :
- motion of particles is minimal
- a substance has no transferrable energy
- an ideal gas at constant pressure would reach 0 volume
17
Q
Conversion between degrees Celsius and Kelvin
A
Kelvin = Degrees Celsius -273.15
18
Q
Catalysts
A
Used to increase the rate of a chemical reaction
- achieve this by providing an alternate reaction pathway that has a lower Ea than the uncatalysed pathway
- remains chemically unchanged at the end of the reaction. In some, catalyst can be re-used, provides a further economic benefit.
19
Q
Catalysed reaction pathway
A
Has a lower Ea than uncatalysed reaction pathway
- original Ea for reaction remains unchanged
20
Q
Catalysts and activation energy
A
Catalysts don’t lower Ea of a reaction
- they proved an alternate reaction pathway that has a lower Ea
21
Q
How do catalysts increase rate of a chemical reaction?
A
- provide an alternative reaction pathway w/ a lower Ea
- a greater proportion of particles now have energy equal to, or greater than, original Ea for reaction
- results in an increased frequency of successful collisions between reactant particles
- increases rate of reaction
22
Q
Enzymes
A
Biological catalysts
- are large protein molecules that enable biochemical reactions in living things to take place at relatively low temp.
- enzyme molecules are folded into special shapes that can accommodate reactant molecules in their ‘active sites’
- incredibly efficient, most only catalyse on reaction involving one particular molecule or pair of molecules (specific)
23
Q
Substrate
A
Molecule that fits into active site of an enzyme and reacts
24
Q
Enzymes as catalysts
A
Provide an alternative reaction pathway with a lower Ea than the uncatalysed reaction
25
Rate of reaction
The change in concentration of a reactant or product per unit of time
- measure of the 'speed' of reaction
Unit: mol/dm^3/s
26
Equation for rate of reaction
Rate of reaction = (increase/decrease in reactant concentration)/ change in time
27
Graphs and rates of reaction
Concentration of reactant vs. time:
- gradient of line represents rate of reaction (gradient decreases with time)
Concentration of product vs. time:
- gradient of line decreases with time
28
Instantaneous rate of reaction
Rate of reaction at a particular time
29
Average rate of reaction
Rate of reaction averaged over the whole time period
30
Determining instantaneous rate of reaction graphically
- determined graphically from a change in reactant or product concentration against time
- equivalent to gradient of line at that particular point in time
- calculated by drawing a tangent to curve at a particular point, and dividing rise/run
- gradient will be positive or negative, depending on whether y-axis plots product or reactant conc.
- steeper the gradient, faster the rate of reaction at that time
31
How does the rate of reaction change over course of a reaction?
Product concentration vs. time graph:
1. At start of reaction, rate is fastest (gradient is the steepest)
- at this point, conc. of reactants is at its highest, high frequency of collisions between reactant particles
2. As reaction proceeds, gradient becomes less steep as conc. of reactants continues to decrease- fewer collisions between reactant particles
3. Once all reactants have been consumed in the reaction, there is no more product formed, line becomes horizontal
32
Experimental determination of rates of reaction
Experiments to measure rate of reaction measure rate of increase or decrease in concentration of a reactant or product, either directly or indirectly
33
Experimental techniques used to measure rates of reaction
1. Measuring production of a gas
- volume of gas produced, measured using a gas syringe, from a reaction is recorded at set time intervals, and results are plotted
- measured using a mass balance, change in mass per unit time is measured as gas escapes from flask
2. Measuring change in ion concentration
- use conductivity meter to measure change in ion concentration of an aqueous solution
- increase/decrease in hydrogen ion conc. can be measured using a pH probe
3. Measuring time taken for formation of a precipitate
- time taken for reaction to reach a certain point is measured
4. Measuring change in concentration by titration
- involves removing a sample of reaction mixture and stopping the reaction by placing it in cold water
- concentration of reactant or product is determined by titration w/ standard solution
- samples of reaction mixture are removed at regular time intervals, and reaction is stopped (remaining amount of reactant is determined through titration)
34
Temperature and effect on rate of reaction
Particles in solids, liquids and gases are heated
- gain KE, moving with higher velocity
- frequency of collisions between reactant particles increases, as particles have more KE at higher temp.
- at a higher temp., a greater proportion of reactant particles will collide w/ energy equal to or greater than Ea for reaction
Low temp:
- lower frequency of collisions
- collisions are less energetic
Higher temp:
- greater frequency of collisions
- reactant particles also collide w/ greater energy
NB/ for most reactions, increase in temp. of 10 degrees C doubles rate of reaction
35
Concentration and rate of reaction
Increasing concentration of a solution increases rate of reaction
- rate of reaction is greater in conc. solution
- increased frequency of collisions between reactants
- thus, greater probability of successful collisions between reactant particles
- increase in rate of reaction
36
Pressure and rate of reaction
increase in pressure for reaction w/ gases = more gas particles in a given volume
- increase in concentration
- increase in frequency of collisions between reactant particles
- increase in rate of reaction
Lower pressure: lower frequency of successful collisions between reactant particles
Higher pressure: higher frequency of successful collisions between reactant particles
37
Surface area and rate of reaction
- braking up a large substance into smaller pieces= increases SA per unit volume
- smaller pieces have a larger SA
- rate of reaction increases as SA:V increases
38
When rate of reaction increases...
Frequency of collisions is increased in all cases, as particles collide more often
- there is a greater probability of successful collisions resulting in a chemical reaction
39
How does concentration increase the rate of a chemical reaction?
- an increase in concentration of reactant leads to an increase in rate of reaction
- due to more frequent collisions between reactant particles in aqueous solutions
- changes in conc. of reactants can have varying effects on rate of reaction, and in some, have no effect on reaction rate
40
Order of reaction
Factor by which concentration of a reactant affects rate of a reaction
- order of a reaction w/ respect to a particular reactant can only be determined from experimental data
- order of reaction w/ respect to a particular reactant will be 0, 1st or 2nd order
- power to which conc. of a reactant is raised in the rate expression
41
Rate expression
Relationship between reactant concentrations and overall rate of reaction
- once the order of reaction w/ respect to each reactant is known, the rate expression for the reaction can be deduced
- can be used to predict rate of a reaction between reactants of known conc. and help establish a reaction mechanism
42
Rate expression for a general reaction
General reaction:
W + X --> Y + Z
Rate expression:
rate = k[W]^a [X]^b
a and b = orders of reaction w/ respect to reactants W and X
k= rate constant
[ ]= specify concentration of reactants
Overall order of reaction = a + b
43
Rate constant, k
Constant for a reaction at a specific temperature, is temperature-dependent
- the only factor that affects its value is a change in temperature
44
Overall order of reaction
Sum of individual orders of reaction for each reactant in the rate expression
45
Zero order reaction
Rate expression:
rate = k [A]^0
Rate is independent of [A], any change in conc. of A doesn't affect rate of reaction
46
First order reaction
Rate expression:
rate = k [A]^1 or rate = k [A]
Rate is directly proportional to [A], if conc. of A is doubled, rate also doubles
47
Second order reaction
Rate expression:
rate = k [A]^2
Rate changes as square of [A]
- if conc. of A is 3x, rate increases by 9 (3 squared)
48
Different orders of reaction and units of rate constant, k
Zero-order:
rate = k[A]^0
k = units of rate = mol/dm^3/s
First-order:
rate= k[A]
k= units of rate/units of conc. = 1/s
Second-order:
rate= k[A]^2
k= units of rate/ (units of concentration)^2 = dm^3/mol/s
Third-order:
rate= k[A]^3
k= units of rate/ (units of concentration)^3 = dm^6/mol^2/s
49
Concentration vs. time and rates of reaction
Used to find order of a reaction w/ respect to a reactant- shape of graph indicates order w/ respect to reactant
- zero-order reactants, graph is a straight line
- first-order reactants, graph is an exponential curve
- second-order reactants, line is a quadratic curve
50
Rate vs. concentration graph
Zero-order:
- produce a horizontal straight line graph
- rate of reaction is independent of conc. of reactant
First-order:
- straight line
- rate is directly proportional to concentration of reactants
Second-order:
- parabolic curve
- rate of reaction increases to a power of 2 as concentration of reactants increases
51
Half-life
Time it takes for the concentration of a reactant to decrease by half
- important in radioactive decay, involves emission of different types of radiation (alpha, beta and gamma)
First-order reaction:
- half-life is constant and independent of initial concentration of reactants
52
Elementary steps
Many reactions don't take place in a single step, but instead, occur as a sequence of steps, known as elementary steps
53
Rate-determining step
Slowest elementary step, determines rate at which reaction can proceed
- overall rate of reaction is dependent on the rate determining step, it is the step with the highest activation energy
- slowest step in a sequence of steps
54
RDS (rate-determining step) and rate expression
- only the reactants in the slowest, rate-determining step appear in the rate expression
55
Requirements for a reaction mechanism to be plausible
- elementary steps must add together to give overall balanced equation for reaction
- mechanism must be consistent w/ experimentally determined rate expression
56
Reaction intermediate
A substance that appears in both steps but not in the overall equation
57
Molecularity
Number of reactant particles in an elementary step
58
Unimolecular
If only one particle is involved in RDS
Elementary step:
A --> products
Rate expression:
rate = k[A]
Involve decomposition or dissociation of a molecule into two or more smaller species
59
Bimolecular
If two particles are involved in RDS
Elementary step:
A + A --> products
A + B --> products
Rate expression:
rate = k[A]^2
rate = k[A][B]
Involve two species colliding and reacting with each other
60
Termolecular
If three particles are involved in RDS
- probability of this occurring is low
- requires simultaneous collision of three reactant particles w/ correct orientation
Elementary step:
A + A + A --> products
A + A + B --> products
A + B + C --> products
Rate expression:
rate = k[A]^3
rate = k[A]^2 [B]
rate = k[A][B][C]
Involve collision of 3 reactant particles but have a low probability of occurring
61
Reaction intermediate
a species formed from the reactants in a chemical reaction which then goes on to react further to form the products
- don't appear in the overall equation as they're produced and used up in consecutive steps in reaction mechanism
- formation of intermediates is a common feature of multi-step reactions
62
Transition state
- also called activated complex
- the highest energy state on a reaction coordinate
- indicates a point at which new bonds are being formed at the same time as old bonds are being broken
63
Catalyst
increases the rate of a chemical reaction by providing an alternative reaction pathway w/ a lower activation energy
- they are chemically unchanged in a chemical reaction- aren't considered to be reactants or products
- appear in elementary step of a reaction and in RDS but not in overall balanced equation for reaction
- works by increasing value of rate constant, k
64
Factors that affect the rate constant
1. Catalyst
| 2. Temperature
65
Effect of temperature on a rate of reaction
- when temp. is increased, rate of reaction also increases
- for many reactions, an increase in temp. of just 10 degrees Celsius can double the rate of reaction
- due to more frequent and energetic collisions between reactant particles
66
Rate constant, k and temperature
- k is temp. dependent
- its value depends on the temp. at which the reaction occurs
- relationship between k and temp. is described by Arrhenius equation
67
Arrhenius equation
k=Ae ^ (−Ea/ RT)
A= Arrhenius constant; pre-exponential factor or frequency factor
Ea = activation energy (J/mol)
T= absolute temp. in kelvin (K)
R = universal gas constant (8.31 J/K/mol mol)
e = a constant; Euler’s number
68
Arrhenius constant, A
- takes into account frequency of collisions and probability they have the correct orientation (or geometry)
- expression e ^ −Ea /RT (exponential factor)
- fraction of molecules that have sufficient kinetic energy to react at a certain temp.
- it shows that an increase in temp. will increase value of k, and thus the rate of reaction
69
Using Arrhenius' equation to calculate activation energy of a reaction
ln(k) = ln(A)− (Ea/RT)
or
ln(k) = ln(A)−(Ea/R) 1/T
70
Graph of ln(k) against 1/T
- is a straight line
- gradient of line = - Ea/R
- y-intercept is equal to ln(A)
- after determining Ea, calculate Arrhenius' equation by substituting the Ea value into it
R = universal gas constant; 8.31 J/K/mol
71
Calculating activating energy (graphically)
The activation energy can be calculated as follows:
Gradient= − Ea/ R
- units: Kelvin (K)
−Ea = gradient × R
- the reaction w/ the steeper gradient has the highest activation energy
72
Activation energy and rate constants
ln (k1 / k2) = Ea/R (1/ T2 − 1/ T1)
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