Iodine Clock Reaction Report CHM 215 BO1 TA: James Pander Mohamed Shalan Due: May 3rd, 2014
2-
Purpose: The goal of this experiment is to measure the rate of reaction of persulfate (S2O8 ) with iodide (I-), using a delayed indicator that turns b lue after a certain concentration of I2 is produced. By measuring the time it takes for the indicator to activate, we can measure the rate and its dependence on varying concentrations and temperatures. Procedure: The procedures used in this lab experiment were derived from A General Chemistry Lab Manual: The Rediscovery Book . No deviations from the instructions occurred. See official citation below: Pickering, Miles. The Rediscovery Book: A General Chemistry Lab Manual . Glenview, IL: Scott, Foresman/Little, Brown Higher Education, 1990. Print. Results: Run 1: 34.44 s S2O8
-
Volume Added
Concentration of Solution after Mixing
0.1 M
20.00 mL
0.4 M
-
.239 M
20.00 mL
0.0956 M
0.00575 M
10.00 mL
0.00115 M
Varying the S2O3 Run 2: 64.69 s S2O8
Rate
()()
KI S2O3
Concentration before Mixing
-
-
concentration: K 2SO4 as the ionic filer
Concentration before Mixing
Volume Added
Concentration of Solution after Mixing
0.1 M
10.00 mL
0.2 M
Rate
()()
.239 M
KI
20.00 mL
0.0956 M
-6
8.89 x 10 S2O3
-
0.00575 M
10.00 mL
0.00115 M
M/s
Run 3: 175.47 s
Concentration before Mixing
Volume Added
Concentration of Solution after Mixing
Rate
S2O8
0.1 M
6.00 mL
0.012 M
3.28 x 10 M/s
KI
.239 M
20.00 mL
0.0956 M
0.00575 M
10.00 mL
0.00115 M
2-
S2O3
-
-
-
Varying the I concentration: KCl as the ionic filler
Run 4: 67.45 s S2O8
-
KI 2-
S2O3
Run 5: 114.25 s S2O8
-
KI
S2O3
-
Concentration before Mixing
Volume Added
Concentration of Solution after Mixing
Rate
0.1 M
20.00 mL
0.04 M
8.52 x 10 M/s
.239 M
10.00 mL
0.0478 M
0.00575 M
10.00 mL
0.00115 M
Concentration before Mixing
Volume Added
0.1 M
20.00 mL
Concentration of Solution after Mixing 0.04 M
.239 M
6.00 mL
0.0287 M
0.00575 M
10.00 mL
0.00115 M
-
Rate
-
5.03 x 10 M/s
Run 5a (with water as ionic filler): 121.16 s S2O8
Concentration before Mixing
Volume Added
0.1 M
20.00 mL
Concentration of Solution after Mixing 0.04 M
KI
.239 M
6.00 mL
0.0287 M
0.00575 M
10.00 mL
0.00115 M
-
S2O3
Rate
-
4.75 x 10 M/s
Rate Dependence on Temperature o
Run 6: 48 C 8.06 s S2O8 KI S2O3
-
-
o
Run 7: 39 C 9.31 s S2O8 KI S2O3
-
-
o
Run 8: 25 C 27.62 s S2O8 KI S2O3
-
-
o
Run 9: 6 C 102.12 s S2O8 KI S2O3
-
-
Concentration before Mixing
Volume Added
0.1 M .239 M 0.00575 M
10.00 mL 20.00 mL 20.00 mL
Concentration before Mixing
Volume Added
0.1 M .239 M 0.00575 M
10.00 mL 20.00 mL 20.00 mL
Concentration before Mixing
Volume Added
0.1 M .239 M 0.00575 M
10.00 mL 20.00 mL 20.00 mL
Concentration before Mixing
Volume Added
0.1 M .239 M 0.000575 M
10.00 mL 20.00 mL 20.00 mL
Concentration of Solution after Mixing 0.04 M 0.0956 M 0.00115 M
Concentration of Solution after Mixing 0.04 M 0.0956 M 0.00115 M
Concentration of Solution after Mixing 0.04 M 0.0956 M 0.00115 M
Concentration of Solution after Mixing 0.04 M 0.0956 M 0.00115 M
Rate
()() -5
7.14 x 10 M/s
Rate
-
6.18 x 10 M/s
Rate
-
2.08 x 10 M/s
Rate
-
1.36 x 10 M/s
-
Log rate Log (1.67 x 10 ) = -4.777
Run 1 Run 4
Log[I ] Log (0.0956)=-1.020
-
Log (8.52 x 10 ) = -5.070 Log (5.03 x 10 ) = -5.298
Run 5
Log (0.0478)=-1.321 Log (0.0287)= -1.542
log rate vs log[iodide] 0 -5.4
-5.3
-5.2
-5.1
-5
-4.9
-4.8
-0.2
-4.7
-0.4
y = 1.0031x + 3.7697 R² = 0.9997
-0.6
e t a r g o l
-0.8 -1
Series1 Linear (Series1)
-1.2 -1.4 -1.6 -1.8
log[iodide]
-
Log rate Log (1.67 x 10 ) = -4.777 Log (8.89 x 10 )= -5.051 Log (3.28 x 10 )= -5.484
Run 1 Run 2 Run 3
Log[S2O8 ] Log (0.04)= -1.398 Log (0.02)= -1.699 Log (0.012)= -1.921
log rate vs log[persulfate] 0 -5.6
-5.4
-5.2
-5
-4.8
-4.6 -0.5
e t a r g o l
y = 0.7192x + 1.9984 R² = 0.954
-1 Series1 -1.5 -2
log persulfate
-2.5
Linear (Series1)
Since both graphs indicate linear relationships between the rate and the ion concentration, the -
2-
m
order of each ion in the rate law is 1 (m=n=1). Rate = k[I ][S2O8 ] Run
Rate
1
1.67 x 10 M/s
-
-
- n
[KI] after mixing (M)
[S2O8 ] after mixing (M)
[KI] after mixing (M)
[S2O8 ] after mixing (M)
0.0956
0.04
0.0956
0.04
Rate Constant
()()
-
0.0956
0.02
0.0956
0.02
00437 0.00465
-
0.0956
0.0112
0.0956
0.0112
0.00406
-
0.0478
0.04
0.0478
0.04
0.00446
-
0.0287
0.04
0.0287
0.04
0.00438
2
8.89 x 10 M/s
3
3.28 x 10 M/s
4
8.52 x 10 M/s
5
5.03 x 10 M/s
°
Run
Rate (M/s)
Rate Constant
ln rate constant
T ( C)
1/T (1/K)
6
7.14 x 10-5 M/s
( )( )
-3.981
48°C
7
6.18 x 10-5 M/s
0.01616
-4.125
39 C
.00320
8
2.08 x 10-6 M/s
0.000544
-7.156
25°C
.00335
9
1.36 x 10-6 M/s
0.0003556
-7.942
6°C
.00358
0.00311
Log rate vs (1/T) 0 -10.003 0.0031 0.0032 0.0033 0.0034 0.0035 0.0036 0.0037 -2 -3 e t a r g o l
-4
y = -6032.7x + 25.067 R² = 0.87791 Series1
-5
Linear (Series1)
-6 -7 -8 -9
1/T (in Kelvin)
-Ea/R is the slope of the line where R= 8.314 J/mol*K. Ea = (6032.7)(8.314) = 50155 J or approximately 50.2 kJ. DISCUSSION:
The objective of this experiment was to, through experimentation, interpret a rate law for this reaction as well as measure the effects of varying either th e persulfate ion or the iodide ion or the temperature. After determining the rate law, which was a second order reaction, we were able to find out the k proportionality constant and subsequently, the activation energ y for the reaction by plotting the k with the temperature. The reason for the addition of the thiosulfate was to aid in the indication of the “blue” reaction time. Test 1 was the standa rd reaction to be compared to. Tests two and three took place with varying persulfate ions while four and five took place with varying iodide concentrations. This is done to determine the order of each of the species. In test 5a, however, we substituted water for a salt solution in order to measure the kinetic salt effect. Experimentally, it was observed that the reaction rate was slightly slower, indicating that the species reacting at the rate-determining step have the same charge. As stated in the manual, “the masking effect of added charges also lowers the repulsion between likecharged reacting ions.” (35). Tests 6- 9 was done to determine the effect of temperature of the reaction rate. As predicted, the rate was higher at higher temperature. This is because a decreased amount of time for a reaction is an indicator for a higher rate of reaction. This also indicates that the activation energy for the reaction decreases as temperature increases. As with any lab, there will be systematic errors. An indication of these is the gap between the data points and the best fit line. Assuming that all the volumes were exact, this is an indication of wrong concentrations of ions. The prepared solution could’ve had a different actual concentration. Also, the thiosulfate solution has been in a bottle in a dark drawer for a couple of weeks, which means that its ion concentration could have changed with the experimentally determined one when that lab was preformed. However, the results are agreeable with the chemical principles within the uncertainties.
