EXPERIMENT 1 - FACTORS GOVERNING THE SPEED OF CHEMICAL
REACTIONS
OBJECTIVE To determine how concentration and temperature affect the speed of a chemical reaction.
THEORY
Part A - Effect of Concentration on Reaction Rate
In order for a chemical reaction to take place, the reacting molecules must collide or come very close together. Thus, for the general reaction
a A + b B ® products rxn (1-1)
in which a moles of reactant A combine with b moles of reactant B, one would expect
the rate of collision, and thus, the rate of product formation to be related to concentrations of reacting species.
At a given temperature, the rate is proportional to the concentration as
rate µ [A]m [B]n eqn (1-1)
or
rate = k [A]m [B]n eqn (1-2)
where the proportionality constant, k, is the rate constant. The rate constant will have a unique value for any particular process at a given temperature. [A] and [B] are concentrations of species A and B in moles per liter. m and n are the exponents to which the concentrations of A and B are raised respectively. The value of m determines the order of the reaction with respect to reactant A. Similarly, the value of n is the order of the reaction with respect to reactant B. The sum of the exponents m and n gives the overall order of the reaction. The powers to which the concentrations are raised, m and n, MAY or MAY NOT be the same as the stoichiometric coefficients a and b for the reaction. Table (1-1) gives of some reactions and their rate equations.
Table (1-1) - Some chemical reactions and their rate equations.
Example / REACTION / RATE LAW / OVERALLORDER
a / 2 HI ® H2 + I2 / Rate = k[HI]2 / 2
b / 2 NO + H2 ® N2O + H2O / Rate = k[NO]2 [H2]1 / 3
c / 2 N2O5 ® 4 NO2 + O2 / Rate = k [N2O5] / 1
In general, the order of a reaction CANNOT be determined by inspecting the chemical equation but must be determined experimentally [see example c in Table (1-1)].
Everyday experience tells us that the rates of chemical reactions increase with increasing temperature. Many foods spoil at room temperature but resist decay when stored in a freezer. Familiar fuels such as gas, oil and coal are relatively inert at room temperature but burn rapidly at elevated temperatures.
The order of a reaction is very useful because it allows us to predict the influence of concentration on the speed of the reaction. For a 1st order reaction (example c) doubling the concentration, doubles the reaction rate. But if a reaction is 2nd order (example a) doubling the concentration increases the reaction rate by 4 times.
Part B - Effect of Temperature on Reaction Rate
Temperature affects the speed of a reaction. At higher temperatures, the velocity of the molecules increases and thus the molecules collide more frequently and with more energy. If a reaction is investigated at a number of different temperatures, the rate constant k values found usually show a quite dramatic increase, often several orders of magnitude, over a relatively small temperature range. For many reactions there is approximately a two- to three-fold increase in rate for every 10oC rise in temperature.
The relationship between the temperature, in K, and the rate constant k is given by the Arrhenius equation,
eqn (1-3)
or
eqn (1-4)
or
eqn (1-5)
( y = mx + b )
where A is the pre-exponential factor, R is the gas constant, T is the absolute temperature, and Eact is the activation energy. Equation (1-5) shows that log k is a linear function of the reciprocal absolute temperature.
In order for the reactants to form products, collisions between the reacting molecules are necessary. But collision alone is an insufficient criterion for the production of the products. The kinetic energy of the colliding molecules must be greater than the required minimum energy level before the reactants can be converted to products. This energy level is called the activation energy, Eact. Activation energy is unique for a particular chemical reaction. For a given temperature, reactions that have large activation energies would be slower than the ones that have smaller activation energies.
In this experiment we will note how concentration and temperature affect the speed of the following reaction.
S2O82- (aq) + 2 I- (aq) ® 2 SO42- (aq) + I2 (aq) rxn (1-2)
persulfate iodide sulfate iodine
ion ion ion
For this reaction, the rate law takes the form:
rate = k [I-]m [S2O82-]n eqn (1-6)
The rate of this reaction depends on the concentrations of the reactants, I- and S2O82- ions and temperature. Our method for measuring the rate of the reaction involves what
is called a "clock" reaction. Rxn (1-2) is conducted in the presence of a starch indicator and thiosulfate, S2O32-, ions which serve to remove I2 as soon as it is formed. Therefore, the following reaction will also occur simultaneously in the reaction flask.
2 S2O32- (aq) + I2 (aq) ® S4O62- (aq) + 2 I- (aq) rxn (1-3)
thiosulfate ion iodine tetrathionate ion iodide ion
Compared to rxn (1-2), rxn (1-3) is essentially instantaneous. The I2 produced in
rxn (1-2) reacts completely with the S2O32- ions present in the solution, until all the S2O32- ions has reacted. At this point, any further formation of I2 from rxn (1-2) turns the starch indicator blue.
By carrying out rxn (1-2) in the presence of S2O32- ion and a starch indicator, a built-in "clock" is introduced into the system. The clock tells us "when" there will be sufficient I2 produced by rxn (1-2) to use up all the S2O32- ions originally added. By introducing a fixed amount of S2O32- ions in each run, the blue colour will always occur after the same amount of I2 is produced. The time it takes to produce this amount of I2 can be measured and this time is inversely proportional to the average rate of reaction. If the amount of S2O32- ions is small compared to the amount of I- and S2O82- used, the colour change will occur before any appreciable amounts of reactants are used up, and the concentrations of reactants and the rate in eqn (1-6) will remain essentially constant during the time interval over which the rate is measured.
In experiment 5 we will been examining the catalytic effect of Cu2+ on rxn (1-2).
PROCEDURE
[NB - All glassware has to be CLEAN and DRY. The drying process involves
rinsing with acetone and blasting the inside surface with compressed air.]
Part A - Effect of Concentration on Reaction Rate
[NB - Part A is carried out at 20oC]
1. Set up three burettes containing the following solutions:
(Remember burettes need to be acclimatized, and need not be dried).
(i) 0.200 M KI,
(ii) 0.100 M (NH4)2S2O8,
(iii) distilled H2O.
The Na2S2O3 will be dispensed using a bottle-top dispenser.
2. Obtain fourteen 125 mL Erlenmeyer flasks. Clean, dry and label them as follows:
A1, A2, A3, A4, A5, A6 and A7
B1, B2, B3, B4, B5, B6 and B7.
3. Prepare the 'A' solutions according to Table (1-2).
[NB - The total volume of each Erlenmeyer flask is 30.00 mL and 3 drops
of starch indicator]
Table (1-2) - Contents of the seven 'A' solutions.
SOLUTION / 0.005 M Na2S2O3(mL) / 0.200 M
KI
(mL) / distilled
H2O
(mL) / 3 % starch
indicator
A1 / 10.00 / 20.00 / 0.00 / 3 drops
A2 / 10.00 / 20.00 / 0.00 / 3 drops
A3 / 10.00 / 20.00 / 0.00 / 3 drops
A4 / 10.00 / 20.00 / 0.00 / 3 drops
A5 / 10.00 / 15.00 / 5.00 / 3 drops
A6 / 10.00 / 10.00 / 10.00 / 3 drops
A7 / 10.00 / 5.00 / 15.00 / 3 drops
4. Prepare the 'B' solutions according to Table (1-3).
[NB - The total volume of each Erlenmeyer flask is 20.00 mL]
Table (1-3) - Contents of the seven 'B' solutions.
SOLUTION / 0.100 M (NH4)2S2O8(mL) / distilled
H2O
(mL)
B1 / 5.00 / 15.00
B2 / 10.00 / 10.00
B3 / 15.00 / 5.00
B4 / 20.00 / 0.00
B5 / 20.00 / 0.00
B6 / 20.00 / 0.00
B7 / 20.00 / 0.00
5. Use lead donuts to stabilize the fourteen Erlenmeyer flasks in the 20oC water bath.
Allow the flasks to come to thermal equilibrium by leaving them in the bath for at
least 5 minutes.
[In the following steps you will be mixing the 'A' solutions with the 'B' solutions.]
ie - Mix: (i) A1 with B1,
(ii) A2 with B2,
.
.
.
(vii) A7 with B7.
6. Pour the content of solution A1 rapidly into B1 while swirling. Start the timer
immediately and record the time (in sec) for the appearance of the blue colour.
For the duration of the reaction, keep swirling the flask containing the combined solutions and keep the flask immersed in the 20oC water bath.
7. Repeat step 6 for the remaining 6 pairs of solutions.
Part B - Effect of Temperature on Reaction Rate
[NB - Part B is carried out at 0oC, *20oC, 30oC, and 40oC]
* - In Part B you will be studying the effect of temperature
when solutions A4 and B4 are mixed at the temperatures
indicated above. You will not need to repeat mixing the
solutions at 20oC. Simply obtain the data for this temperature
from Part A.
1. Obtain six 125 mL Erlenmeyer flasks and clean and dry the flasks as per
instructions given previously in Part A.
2. Prepare three A4 solutions and three B4 solutions by following instructions
given in Table (1-2) and Table (1-3).
3. Immerse one pair of A4/B4 solutions in 0oC ice bath, the second pair of solution
in the 30oC water bath, and the third pair of solution in the 40oC water bath.
Stabilize the flasks using lead donuts and allow the flasks to immerse in the bath for 5
minutes to come to thermal equilibrium
[Remember the 20oC temperature does not need to be repeated here.
Data for A4/B4 at 20oC is obtained from Part A .]
4. Measure the temperatures of the ice/water bath using a thermometer. Record
the actual temperatures on the data sheet.
5. Pour the content of solution A4 rapidly into B4 while swirling. Start the timer
immediately and record the time (in sec) for the appearance of the blue colour.
For the duration of the reaction, keep swirling the flask containing the combined solutions and keep the flask immersed in the water bath.
6. Repeat step 5 for the remaining two pairs of solutions.
DATA SHEET
Part A - Effect of Concentration on Reaction Rate
Temperature of reaction = ______
Table (1-4) - Data for determining 'n', the reaction order with respect to S2O82-.
Solution / [S2O82-](moles/L) / log [S2O82-] / t (sec) / 1/t (sec-1) / log (1/t)
A1/B1
A2/B2
A3/B3
A4/B4
Show sample calculation of [S2O82-] using solution A4/B4.
Table (1-5) - Data for determining 'm', the reaction order with respect to I-.
(moles/L) / log [I-] / t (sec) / 1/t (sec-1) / log (1/t)
A4/B4
A5/B5
A6/B6
A7/B7
Show sample calculation of [I-] using solution A4/B4.
Summary of result from Part A
rate = k [I-]m [S2O82-]n
Order of reaction with respect to[ I- ], 'm'
Order of reaction with respect to [S2O82-], 'n'
Overall reaction order
Part B - Effect of Temperature on Reaction Rate
From the sample calculations in Part A, transfer the [I-] and [S2O82-].
[ I- ] = ______
[S2O82-] = ______
Table (1-6) Effect of Temperature on the Reaction Rate
Temperature, T(oC) / Temperature, T
(K) / 1/T
(K-1) / time, t
(sec) / log (1/t)
Note: The 20oC data is obtained from Part A
INTERPRETATION OF DATA
Part A - Effect of Concentration on Reaction Rate
By studying the contents of the mixtures A1/B1, A2/B2, A3/A3, and A4/B4 carefully, you will notice that these four sets of solutions have a constant [I-] while [S2O82-] varies. Data from this set of solutions can be applied to rate equation (1-5). We take advantage of the constant [I-] and simplifiy eqn (1-6) to become
rate = k'[S2O82-]n eqn (1-7)
where k' = k [I-]m. k' is a constant since k and [I-]m are constants. Taking logs of both sides of eqn (1-7) we have
log (rate) = n log [S2O82-] + log k' eqn (1-8)
( y = mx + b )
A plot of "log versus log [S2O82-]" should be a straight line with slope equal to n, which is the order of the reaction in terms of S2O82-. The value of n can be rounded
off to the nearest integer.
Similarly, the contents of the mixtures A4/B4, A5/B5, A6/B6, and A7/B7
have constant [S2O82-] and varying [I-]. Data from this set of solutions can be applied to rate equation (1-6). This time, we take advantage of the constant [S2O82-] and simplifiy eqn (1-6) to become
rate = k'' [I-]m eqn (1-9)
where k'' = k [S2O82-]n. k'' is a constant since k and [S2O82-]n are constants. Taking logs of both sides of eqn (1-9) we have
log (rate) = m log [I-] + log k'' eqn (1-10)