Bc Chemicals and Equipment Lab

Pre-lab should consist of the following ;

1- Purpose: A brief description of why you are performing the experiment. BRIEFLY state the expected outcome for the experiment that is being performed.

2- List of chemicals and equipment

3-
Procedure: This is the heart of your laboratory performance. There should be enough information outlined so that you can perform the experiment without having to refer to the lab manual or any other source.

You should write what you will do in your own words using proper grammar (DO NOTWRITE IN FIRST PERSON); copying verbatim from the book is unacceptable.

4- Safety and Disposal: Read the lab material provided. list all possible safety and disposal concerns for the experiment of the day.

Kinetics Experiment
This experiment will extend over two weeks. In the first week we will collect the data, in the
second week we will analyze the data.
Background
The rates of chemical reactions and the conditions, which control them, are important for many
aspects of chemistry. They help to determine the best conditions to perform certain reactions and
also give insight into the elementary processes of reactions. Consider a general reaction:
aA + bB→cC + dD
We can monitor how fast this reaction happens by measuring the rate of disappearance of A, dA/dt or B, -dB/dt or the rate of appearance of C, dC/dt or D dD/dt. The rate of the reaction is
defined in terms of the stoichiometric coefficients as:
Rate of reaction = –
1 dA
1 dB 1 dC 1 dD
==
=
a dt
b dt
c dt d dt
This accounts for the fact that reactants are disappearing (the negative sign) and products are
appearing and also that the rate of disappearance or appearance depends on how many of each
substance are involved in the reaction.
The rate of the reaction at a particular temperature depends on the concentration of the reactants
so that:
rate of reaction = k[ A]n [ B]m
(1)
where k is the rate coefficient, and n and m are typically the 0, 1, or 2. This expression is known
as the rate law. The exponents, n and m, are determined experimentally by measuring how the
rate of the reaction changes with the concentration of the reactants.
Picture the reaction of acetone with bromine.
O
H
H
O
H
C
H
C
C
H
H
+
Br
Br
H+
H
H
Br
C
H
C
C
H
H
H
Br
In this case, the concentration of one of the reactants, Br2, can be measured using a SPEC 20. If
we measure the concentration of Br2 versus time the rate of change or slope of the line will give
us:
− d [ Br2 ]
= k[acetone]n [ H + ]m [ Br2 ] p
dt
H+ is included because this reaction does not occur unless acid is present, and this was indicated
by the H+ over the reaction arrow. The goal is to find the exponents n, m, p, and the rate
coefficient, k. With that information, the reaction is understood, and a mechanism can be
proposed.
Here is the general theory.
If we start the reaction with a large excess of acetone and H+, compared to the concentration of
the bromine, their concentrations do not change measurably over the first few minutes so we can
assume they are constant. To determine the exponent for Br2, p, first assume the concentrations
of H+ and acetone are constant. Then the rate is:
− d [ Br2 ]
= k ‘ [ Br2 ] p
dt
where k ‘ = k[acetone]n [ H + ]m
Rearrange equation 2 to get
− d [ Br2 ] = k ‘ [ Br2 ] p dt
If p=0, integration gives
 − d [ Br ] =  k ‘[ Br ] dt
 − d [ Br ] =  k ‘ dt
p
2
2
2
[ Br2 ]t − [ Br2 ]0 = −k ‘ t
and the concentration of bromine will linearly decrease with time.
Figure 1
Zero order reaction
0.0045
0.004
y = -7E-06x + 0.004
R2 = 0.9996
Concentration Bromine
0.0035
0.003
0.0025
0.002
0.0015
0.001
0.0005
0
0
50
100
150
200
250
300
350
400
Time
If p= 1 the integrated rate equation is:
 − d [ Br ] =  k ‘[ Br ] dt
p
2
2
1
 − d [ Br ] [ Br ] =  k ‘ dt
2
2
ln[ Br2 ]t − ln[ Br2 ]0 = − k ‘ t
and the natural log of the concentration of Br2 decreases linearly with time.
Figure 2
450
500
First Order Reaction
0.001
y = -5E-06x + 0.0009
R2 = 0.9736
Concentration
0.0008
0.0006
0.0004
0.0002
0
0
20
40
60
80
100
120
140
160
Time
Figure 3
First Order Reaction
-6.8
-7
y = -0.008x – 6.905
R2 = 0.9995
Natural log Concentration
-7.2
-7.4
-7.6
-7.8
-8
-8.2
0
20
40
60
80
Time
100
120
140
160
It can be quite difficult to distinguish between the zero order reaction and the first order reaction.
However, if both concentration and the natural log of concentration versus time are plotted, one
graph will be straighter than the other. It may be easier to determine this by adding a trendline in
excel to both plots and comparing the R value. R is a measure of how well the data fits to a
straight line. If R is 1.00, the line is exactly straight. In comparing two data sets, the one with R
closer to 1.00 is more accurately represented by a straight line.
If p = 2 the integrated rate equation is:
 − d[ Br ] =  k ‘[ Br ] dt
p
2
2
1
 − d[ Br ] [ Br ] =  k ‘ dt
2
2
2
1
1
−
= −k ‘ t
[ Br2 ]t [ Br2 ]0
and the inverse of the concentration of Br2 increases linearly with time.
Figure 4
Second Order Reaction
0.0025
y = -4E-06x + 0.0016
R2 = 0.8558
Concentration
0.002
0.0015
0.001
0.0005
0
0
50
100
150
200
Time
Figure 5
250
300
350
Second Order Reaction
-6
-6.2
y = -0.0045x – 6.3875
R2 = 0.9635
-6.4
Natural log Concentration
-6.6
-6.8
-7
-7.2
-7.4
-7.6
-7.8
0
50
100
150
200
250
300
350
-8
Time
Figure 6
Second Order Reaction
2500
y = 5.1391x + 498.79
R2 = 0.9999
1/Concentration
2000
1500
1000
500
0
0
50
100
150
200
Time
250
300
350
To determine the reaction order for bromine, p in equation 2, we examine how the concentration
varies with time. If [Br2] versus time is a straight line, the reaction is zero order in Br2. If ln
[Br2] versus time is a straight line, it is first order in Br2, and if 1/[Br2] versus time is a straight
line, the reaction is second order in Br2. For each case the slope of the straight line is k’, and in
all cases C, the constant of integration, depends on the initial concentration of bromine, and will
not need to be determined. As can be seen from the above plots, it is best to plot concentration
versus time, natural log of concentration versus time and the inverse of concentration versus
time, add a trendline to the data and compare the values of R for these to make the final
assessment. This will give the order with respect to Br2.
To determine the reaction order for the acetone and the H+ vary the initial concentrations of
these reactants and see how they affect the reaction rate. If we double the amount of acetone,
and the rate does not change, the reaction if zero order for acetone. If we double the acetone and
the rate doubles, it is first order in acetone, and if you double the acetone and the rate increases
by a factor of 4 the reaction is second order. The same will hold true for H+.
Procedure:
This experiment requires careful planning. Be sure you have assembled everything you need and
plan ahead. If you do not remember how to use the SPEC 20 see the appendix.
First you will need to measure the absorbance of some known concentrations of bromine and
make a standard curve.
1. Turn on the spectrometer and let it warm-up for 15 minutes. Set the wavelength to 395 nm
and make sure it is zeroed. Take about 10 ml of each of the three bromine standards. Measure
the absorbance of the standards, and record their concentrations. Make sure your standard curve
is a straight line. Plot absorbance versus concentration for Br2. This will allow you to convert
absorbance measured into concentration of Br2. You will use this later to determine the
concentration of bromine for your kinetics runs.
2. Label a small beaker and put about 15 mL of the 8M acetone solution in it. Also take about
15mL of the 2M HCl solution, and about 50 mL of the 1x 10-3 M bromine solution. Get 4, 10
mL volumetric flasks and label them one through four. You will measure the rate for 4 different
mixtures. To do this you must add the following amounts of the reactants to the volumetric
flasks and fill to the line with distilled water. Each of these will be measured one at a time. Do
not prepare a sample until you are ready to take your readings.
Table 1.
Sample
mL Bromine
mL HCl
mL Acetone
distilled water
1
2
3
2.0
2.0
2.0
1.0
2.0
1.0
1.0
1.0
2.0
to the 10 mL line
to the 10 mL line
to the 10 mL line
3. Quickly mix together sample #1 in a 10 mL volumetric flask, by placing 2.0 mL of the Br2
solution, 1.0 mL of the H+ solution, 1.0 mL of Acetone solution and then adding distilled water
to the 10.0 mL line. Cover it with a small amount of parafilm and mix it thoroughly. Transfer
the solution to a cuvette, place in the spectrometer and measure the absorbance every 15 seconds
for about 5 minutes or until the absorbance goes to zero.
4. Quickly mix together sample #2 in a 10 mL volumetric flask, by placing 2.0 mL of the Br2
solution, 2.0 mL of the H+ solution, 1.0 mL of Acetone solution and then adding distilled water
to the 10.0 mL line. Cover it with a small amount of parafilm and mix it thoroughly. Transfer
the solution to a cuvette, place in the spectrometer and measure the absorbance every 15 seconds
for about 5 minutes or until the absorbance goes to zero.
5. Quickly mix together sample #1 in a 10 mL volumetric flask, by placing 2.0 mL of the Br2
solution, 1.0 mL of the H+ solution, 2.0 mL of Acetone solution and then adding distilled water
to the 10.0 mL line. Cover it with a small amount of parafilm and mix it thoroughly. Transfer
the solution to a cuvette, place in the spectrometer and measure the absorbance every 15 seconds
for about 5 minutes or until the absorbance goes to zero.