Title:
Kinetic: Factor affecting rates of reaction Objectives:
To study the factors affecting rates of reaction Introduction:
Chemical kinetics is the area of chemistry concerned with the speeds or rates at which reaction occur. There are essentially three variables that affect the rate of a chemical reaction; concentration of reactants, presence of catalyst and temperature. A catalyst will increase the rate of reaction without being consumed in the reaction itsel f. As the temperature increases, the rate of reaction increases. For the effect of temperature and catalyst on the rate of reaction, we will study the reaction of oxalic acid H2C2O4 with acidified potassium permanganate, KMnO 4, Manganese Sulphate, MnSO4 will be used as catalyst. The rate of reaction is measured b y the time for the dark purple coour of KMnO 4 to disappear. The equation for the reaction is 2MnO4- + 5C2O42- + 16H+ → 2Mn2+ + 10CO2 + 8H2O
Apparatus
Chemicals
Stopwatch
0.1M hydrogen peroxide, H 2O2
Hot water bath
0.1M potassium iodide, KI
Test tubes
2M sulphuric acid, H2SO4
10cm3 measuring cylinder
Starch solution
Dropper
0.01M
potassium
permanganate,
KMnO4
50cm3 beaker
0.5M oxalic acid, H 2C2O4
Conical flask
0.1N manganese sulphate, MnSO 4
0.005M sodium thisulphate, Na 2S2O3
PPE (Personal Protective Equipment)
Appropriate PPE, goggles, lab coat and proper shoes
Procedure
A) Effect of Temperature 1. 3 cm3 of KMnO4 and H2SO4 is measured into a test tube 2. 3cm3 of oxalic acid is measured into another test tube and oxalic acid is added quickly into the first test tube. The mixture is shake and the stopwatch is started. The time is recorded until the dark purple colour KMnO 4 is disappeared. 3. Step 1 and 2 is repeated at 40°C, 50°C and 60°C.
B) Effect of presence of catalyst 1. Step 1 and 2 in B is repeated but this time, 3 drops of MnSO 4 is added to the first test tube
A. Effect of Temperature
Concentration of No.
KMnO4 [KMnO4], M
Temperature (oC)
Time (sec)
Rate of reaction (Ms-1)
1
0.01
27
346
- 0.0000289
2
0.01
60
38
- 0.0002632
3
0.01
50
44
- 0.0002273
4
0.01
40
58
- 0.0001724
Calculation to seek for rate of reaction as follows:
1
2
Rate = d[KMnO4] dt
Rate = d[KMnO4] dt
Rate = - [KMnO4]i tfinal
Rate = - [KMnO4]i tfinal
3
Rate = - (0.01 M) (346 sec)
Rate = - (0.01 M) (44 sec)
Rate = - 0.0000289 Ms-1
Rate = - 0.0002273 Ms-1
Rate = d[KMnO4] dt
Rate = d[KMnO4] dt
Rate = - [KMnO4]i tfinal
Rate = - [KMnO4]i tfinal
4
Rate = - (0.01 M) (38 sec)
Rate = - (0.01 M) (58 sec)
Rate = - 0.0002632 Ms-1
Rate = - 0.0001724 Ms-1
* rate = -ve values -ve means the rate is decreasing.
B. Effect of Catalyst
Concentration of No.
KMnO4 [KMnO4], M
1
0.01
Temperature o
( C) 25
Calculation to seek for rate of reaction as follows:
Rate = d[KMnO4] dt
1
Rate = - [KMnO4]i tfinal Rate = - (0.01 M) (315 sec) Rate = - 0.0000317 Ms-1
Catalyst
MgSO4
Time (sec) 315
Rate of reaction (Ms1
)
- 0.0000317
Discussion
Rate of reaction is referred to the speed of a chemical reaction occurs. Based on the collision theory, f or a reaction between two particles to occur, the particles must collide and the collision must be effective. The particles must collide with enough kinetic energy and the correct orientation. The more frequently effective collision occur, the faster the rate of reaction. The rate of a reaction can be increased by changing certain factors in a reaction like temperature or concentration of reactants and lowering th e activation energy. The factor of temperature and catalyst are being investigated in the experiment. Rate of reaction can be explained by using the Arrhenius equation:
Based on the result obtained in Procedure B , the increase in temperature lead to faster rate of reaction. The fastest time taken in this experiment is with the temperature at 60°C while the slowest time taken in this experiment is with the room temperature at 27°C. From the Arrhenius equation, it can be seen that the temperature i s directly proportional to the rate of reaction. This can be explain further by collision theory. In collision theory, increasing the temperature will lead to the increment of the average kinetic energy. It will make the particle more entropic and this will lead to the higher frequency of the collision. The particle s also will move in higher speed and thus, ease the process of breaking the bond between them. Thus, the chance of effective collisions increases and therefore, the rate of reaction increases. Activation energy is the minimum amount of energy required to initiate a reaction. The Arrhenius equation relates activation energy to the rate at which a chemical reaction proceeds. A substance that lowers the activation energy of a chemical reaction is called a catalyst. Basically, a catalyst acts by modifying the transition state of a reaction. Catalysts are not consumed by the chemical reaction and they don't change the equilibrium constant of the reaction. Based on the result obtained in Procedure C, the presence of the catalyst will increase the rate of reaction. The time taken for the reaction at the room temperature without using any catalyst was 263 second, but when the catalyst has been taken place, the time taken is changed to the 118 second. This can be concluded that the presence of the catalyst may affect the rate of reaction. Presence of catalyst may alter the rate of reaction by changing the reaction and made an alternative pathway with different activation ener gy of the reaction. Catalyst will speed up the rate of reaction by lowering the activation barrier. In conclusion, the presence of catalyst will make the activation energy become lower and end up with higher rate of reaction.
Errors
During the experiment, there are some errors that occur which may affect the results of the experiment i) The volume of chemicals in each test tube may varies and not accurate due to error that
occured when calibrating, reading and transferring the chemicals. These errors might affect the concentration of each test tube, thus affect their rate of reaction. ii) The instrument used along the experiment such as test tube is not clean enough, may have
redundant dust or residual from the experiment before it. So, it can affect the observations of the reaction. iii) The temperature specified to each test tube may varies a little from its recommended
temperature and thus may affect the accuracy of the results. iv) The shaking process may not consistent throughout the experiment, thus may affect the
results obtained.
Safety Precautions 1. Must wear proper attire, goggles and gloves when handling hazardous chemicals.
Chemical such as H2SO4 is corrosive and it may infect our skin. Thus, it is important to protect our body when conducting this experiment
2.
Dispose chemical waste in the disposable container
Never pour chemical waste such as H2SO4 in the sink because it may affect the pipe and also environment. Instead, keep it in the disposable container which usuall y is in fume Chamber
Suggestion
When conducting the experiment, there is some factor that affected the result. The results are not accurate and precise
Make sure the apparatus used clean. If needed, rinse the apparatus with distilled water and let it to dry or dry it with clean tissue before using.
The students should read the measuring apparatus perpendicular with eyes. Students should avoid parallax errors during taking the measurements.
Each experiment must be repeated three times and take the average value. This practice is essential to improve the results’ accuracy.
The process of shaking the mixture manually can be improved by using agitation machine. This machine will enhance the consistency and accurac y of the results.
Questions
1
Define activation energy. What role does activation energy play in chemical kinetics?
Answer: The minimum amount of energy required to initiate a reaction. Activation energy is the quantity of energy necessary in order for a reaction to carry on from the reactants to the products.
2
Write the Arrhenius equation and define all terms.
Answer: K = Ae-Ea/RT
K = Rate constant Ea = Activation energy
A = Frequency factor T = Kelvin temperature
R = The gas constant
3
Use the Arrhenius equation to show why the rate constant of reaction: a) Decreases with increasing activation energy. b) Increases with increasing temperature.
Answer:
(a) Let's presume activation energy of 25 kJ mol -1. In the equation, we have to note down that as 50000 J mol-1. The value of the gas constant, R, is 8.31 J K -1 mol-1. At 20°C (293 K) the value of the fraction is:
By raising the activation energy just a little bit (to 50 kJ mol -1), this increases:
Thus, according to comparison both two different activation energy above, when the activation energy being increased then the rate constant of reaction, k decreases. Therefore, the rate constant is inversely proportionally to activation energy.
(b) Let's presume activation energy of 50 kJ mol -1. In the equation, we have to note down that as 50000 J mol-1. The value of the gas constant, R, is 8.31 J K -1 mol-1. At 20°C (293 K) the value of the fraction is:
By raising the temperature just a little bit (to 303 K), this increases:
Based on comparison two different temperatures above shows that when the temperature increase, the rate constant of reaction, k also increases. Therefore, the rate constant is directly proportional to temperature.
4
Variation of the rate constant with temperature for the first-order reaction is given in the following table. Determine graphically the activation energy for the reaction.
T (K)
K
(s-1)
298
1.74 X 10-5
308
6.61 X 10-5
318
2.51 X 10-4
328
7.59 X 10-4
338
2.40 X 10-3
Answer: T (K)
(s-1)
ln K
(1/T)
298
1.74 X 10-5
-10.96
0.003356
308
6.61 X 10-5
-9.62
0.003247
318
2.51 X 10-4
-8.29
0.003145
328
7.59 X 10-4
-7.18
0.003049
338
2.40 X 10-3
-6.03
0.002959
ln K = -Ea (1/T) + ln A R y = mx +c c = -5.2
K
slope, m = = -Ea R ln K = -Ea (1/T) + ln A R ((-6.03) – (-10.2)) = -Ea (0.003050 – 0.003290) + (-5.2) R -Ea = R ((-6.03) – (-10.2)) - (-5.2) (0.003050 – 0.003290) -Ea = - 3.25 x 105 J/mol Ea = 3.25 x 105 J/mol = 3.25 x 102 kJ/mol
Conclusion As conclusion, the rate of reaction is affected by the temperature and presence of the catalyst. Higher temperature of the reactant will lead to higher rate of reaction. The presence of catalyst makes the rate of reaction increases since the formation of products faster than uncatalyzed reaction.
References Boardworks Ltd. (2007) Rates of Reaction. Retrieved from http://www.montgomery.k12.ky.us/userfiles/1893/Classes/20673/ratesofreaction.ppt. on 27/11/2017. Clark, J. (2013).The Collision Theory of Reaction Rates. Retrieved from https://www.chemguide.co.uk/physical/basicrates/introduction.html on 27/11/2017. Clark, J. (2013). Rate Constants and The Arrhenius Equation. Retrieved from https://www.chemguide.co.uk/physical/basicrates/arrhenius.html on 27/11/2017.
Deoudes, A. (2010). A Kinetics Experiment The Rate of a Chemical Reaction: A Clock Reaction. Retrieved from http://cms.montgomerycollege.edu/WorkArea/DAsset.aspx?id=26338 on 27/11/2017.
The Collision Theory and Activation Energy. (n.d.). Retrieved from http://www.chalkbored.com/lessons/chemistry-12/collision-theory.ppt on 27/11/2017.
