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Title: CHEMISTRY EDEXCEL UNIT 4
Description: EDEXCEL Board A2 Level Chemistry Unit 4 SECTION 1: HOW FAST? - RATES
Description: EDEXCEL Board A2 Level Chemistry Unit 4 SECTION 1: HOW FAST? - RATES
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CHEMISTRY
Unit 4: General Principles of Chemistry I – Rates, Equilibria and Further Organic Chemistry
1) How fast? – Rates
a) Demonstrate an understanding of the terms ‘rate of reaction’, ‘rate equation’, ‘order of
reaction’, ‘rate constant’, ‘half-‐life’, ‘rate-‐determining step’, ‘activation energy’,
‘heterogeneous and homogenous catalyst’
i) Rate of (a chemical) reaction – the rate of change of concentration of a reactant or
product with time / moldm-‐3s-‐1
(1) Rate of reaction cannot be measured directly – can only be determined from
concentration and time data
∆ !"#!$#%&'%("# !"!!"#$%#&% !"#$"%&"& !"#! !"#$ !"!!"#$%&'
(2) Average rate of reaction =
ii)
iii)
iv)
v)
vi)
∆ !"#$
(3) This is only a reasonable assumption if the concentration of a reactant has fallen by
less than 10% during the time elapsed
(4) For a reaction to take place, reactant molecules must collide with kinetic energy
greater than or equal to the activation energy and with the correct orientation
Rate equation – comes from experiments, showing us how rate depends on the
concentration of each species
(1) For a reaction: aA + bB à cC + dD, the rate equation could = k [A]a[B]b[E]e
(2) The subscripts are the stoichiometries in the chemical equation
(3) The superscripts are the partial orders of reaction
(4) E could be a catalyst (to be involved in the rate equation of a reaction, the species
does not have to be a reactant or product necessarily)
Overall order of reaction – sum of the powers to which the concentrations of reactants
are raised in the experimentally determined rate equation (the total of the superscripts)
(1) Partial order of one reactant – the power to which the concentration of that
reactant/species is raised in the experimentally determined rate equation
(2) Cannot be predicted from the chemical equation, depending on both the
stoichiometry and the mechanism of the reaction – have to be found experimentally
Rate constant – ‘k’ is the rate constant, which varies with…
(1) Orientation factor – the complexity of the geometry of the molecules eg
...
1
(2) The activation energy of the reaction – a large activation energy results in a large,
negative exponent and therefore a small value for the rate constant
(3) The temperature – rise in temperature increases the ‘RT’ so the value for the
exponential term is less negative and k gets larger, so the rate of reaction increases
(4) The presence of any catalyst – lowers the activation energy so the exponent
becomes less negative and k gets larger, so the rate of reaction increases
Half-‐life – time taken for the concentration of a reactant to halve
Rate-‐determining step – the slow step that is so slow compared to the subsequent fast
step(s)
vii) Activation energy -‐ the minimum energy required before a reaction can occur (a certain
minimum energy that reactant particles collide to produce a reaction)
viii) Homogenous – catalysts in the same phase as the reactants, usually in gas phase or in
aqueous solution
(1) Work by reacting with one of the reactants to form an intermediate compound eg
...
Rate of reaction depends on the catalyst and one of the reactants
ix) Heterogeneous – catalysts in a different phase to the reactants
(1) Where gases are adsorbed onto active sites and products leave (desorbed)
(2) Eg
...
Iron catalyst stays in the reaction chamber as products pass
x) Maxwell-‐Boltzmann distribution of energy
(1) The higher the temperature of a solution, the faster the rate of reaction
(2) If we increase the temperature by 10˚C, the rate of reaction doubles (exponential)
(3) As the temperature of the solution increases the reactant particles (can be molecules
or ions) move more quickly as they now have an increased amount of kinetic energy
(4) They have a greater range of energies and the average energy is increased
(5) Greater proportion of particles have sufficient energy to overcome activation energy
(6) Reactant particles also come into contact more as the average speed of the
molecules is increased so there are more collisions/sec – an increased collision
frequency (negligible compared to the increased proportion of successful collisions)
(7) There is therefore an increased number of successful collisions per second with
energy in excess of the activation energy thus making the rate of reaction higher
(8) Total number of molecules with energy equal to or greater than a particular energy
value is given by the area under the graph to the right of that energy
(9) The shaded area to the right of the activation energy is the fraction of molecules that
have sufficient energy at the specified temperature to react on collision, providing
that the orientation of collision is correct
xi) Effect of pressure on the rate of a gaseous reaction
(1) The higher the pressure in a mixture of gasses, the faster the rate of reaction
(2) As the pressure is increased at a constant temperature, the molecules become
packed more closely together
(3) There is no change in their speed or energy so the proportion of successful collisions
that result in a reaction remains constant
(4) There are more reactant particles that can come into contact with one another per
given volume (increased concentration)
(5) There are more collisions per second – the collision frequency increases
(6) More likely to be an increased number of successful collisions per second with
energy in excess of the activation energy thus making the rate of reaction higher
b) Select and describe a suitable experimental technique to obtain rate data for a given
reaction, eg colorimetry, mass change and volume of gas evolved
i) Initial rate – the rate at the very beginning of a reaction at the instant that the chemicals
are mixed and before much of the reactant is used up
(1) Usually found by measuring the time taken for the concentration of a reactant or
product to change by a known amount, which must be less than 10% of the initial
concentration of the reactant
(2) It is the gradient of the tangent to the curve at t=0
(3) The initial rate is a useful when comparing reaction rates when a variable is changed
(4) The experiment is repeated, changing the concentration of only one of the reactants
and keeping the concentration of all the others constant to calculate partial orders…
(5) OR excess of all of the reagents is added except one (at least ten times the
concentration of the other reagent) so that addition of these reagents is negligible –
rate of reaction is now only dependent on one reagent and its order can be found
ii) Measuring the volume of gas
(1) Counting the number of bubbles at regular time intervals throughout the reaction
(2) The gas could be collected in a gas syringe or bubbled through a measuring cylinder,
displacing water – measuring the volume of gas produced at regular time intervals
(3) Volume of gas is proportional to the moles of gas and can therefore be used to
measure the concentration of the product
(4) Eg
...
every 15/30 seconds
(2) Seeing how quickly the mass decreases as the gas escapes from the solution
(3) Total mass of gas evolved at any given time is proportional to the concentration of
the reactant when timing started
(4) Record masses – difference is mass of gas
(5) The end amount of gas evolved = the concentration of reactant at the start until the
concentration of reactant = 0
(6) Eg
...
1g at room temperature and pressure
(b) 50cm3 of H2 weighs less than 0
...
by adding solution to ice-‐cold water
(b) Eg
...
CH3COCH3 (propanone) + I2 à CH2ICOCH3 (1-‐iodopropan-‐2-‐one) + HI
v) Colorimetric analysis (colorimetry)
(1) Can be used when one of the reactants or products of the reaction is coloured
(2) Reactants are mixed and a clock started
(3) The light absorbed is measured at set time intervals
(4) Detecting the relative amount of light of a particular frequency absorbed measured
throughout the reaction with a spectrophotometer/photoelectric colorimeter
(5) Colour changes ie
...
Br2(aq) + HCOOH(aq) (methanoic acid) à 2Br-‐(aq) + CO2(g) + 2H+(aq) – following
the absorption of light by bromine
vi) Infrared spectroscopy
(1) Spectrometer is set at a particular frequency
(2) The amount of infrared radiation absorbed at that frequency is measured at regular
time intervals
(3) Eg
...
A sample of one of the optical isomers of 2-‐iodobutane is mixed with NaOH(aq) –
hydrolysis to produce the racemic mixture of butan-‐2-‐ol
(7) Angle of rotation of the plane of polarisation of the plane-‐polarised light gradually
decreases as the single chiral isomer of 2-‐iodobutane is hydrolysed
(8) Eg2
...
0moldm-‐3, the pH
only changes by 1 unit (from 0 to 1) when 90% of the acid has reacted
(4) The pH rises to 2 when 99% of the acid has reacted
(5) This method requires a very accurate, and hence expensive, pH meter to monitor the
change in acid concentration – unsuitable for school laboratory use
c) Investigate reactions, which produce data that can be used to calculate the rate of the
reaction, its half-‐life from concentration or volume against time graphs, eg
...
Comparing A and B – concentration of HCl is halved, concentration of everything
else is constant, therefore the rate halves and the order with respect to HCl is 1
f) Deduce from experimental data for reactions with zero, first and second order kinetics:
i) Half-‐life (the relationship between half-‐life and rate constant will be given if required)
(1) Zero order kinetics – half life = [A]0/2k where k is the rate constant, [A]0 is the initial
concentration
(a) ½ [A]0 = [A]t1/2
(b) kt1/2 = ln(2)
(c) t1/2 = ln(2) / k
(2) First order kinetics – half-‐life is constant at a fixed temperature eg
...
If it
takes 25s for the concentration of any reactant to fall from 8 to 4 units, it will take
50s for it to fall from 4 to 2 units
ii) Order of reaction
(1) Zero order kinetics – change in concentration does not change rate
(2) First order kinetics – doubling concentration doubles rate
(3) Second order kinetics – doubling concentration quadruples rate
iii) Rate equation
(1) Zero order kinetics – overall order of 0
(2) First order kinetics – overall order of 1
(3) Second order kinetics – overall order of 2
iv) Activation energy (by graphical methods only; the Arrhenius equation will be given if
needed)
(1) Arrhenius equation – to calculate the activation energy ( 𝐸! ) for a particular reaction
(a) 𝑘 = 𝐴𝑒 !
!!
!"
(b) ln 𝑘 = −
rearranged as…
!!
!
!
!
+ ln (𝐶)
(i) Where 𝑘 is the rate constant (1/k)
(ii) 𝐴 is the pre-‐exponential factor (negligible)
(iii) 𝐶 is the Arrhenius constant
(iv) 𝑅 is the gas constant – usually has a value of 8
...
HUR and 50% ethanolic solution
(3) Step 1 – slow rate determining step
(4) Step 2 – fast attack of nucleophile on carbocation
(5) Rate = k [(CH3)3CBr]
(a) The rate no longer depends on [OH-‐]
(b) The first step in the two step mechanism is the slow, rate determining step
i) Demonstrate that the mechanisms proposed for the hydrolysis of halogenoalkanes are
consistent with the experimentally determined orders of reactions, and that a proposed
mechanism for the reaction between propanone and iodine is consistent with the data from
the experiment in 4
Title: CHEMISTRY EDEXCEL UNIT 4
Description: EDEXCEL Board A2 Level Chemistry Unit 4 SECTION 1: HOW FAST? - RATES
Description: EDEXCEL Board A2 Level Chemistry Unit 4 SECTION 1: HOW FAST? - RATES