Experiment 1Factors Affecting Reaction
Rates
Chiu, Ina Cathrina R.Salindo, Elysse S.
Introduction
Chemical Kinetics Deals with rates of reactions
(how fast a reaction progresses)
Rate is a measure of quantity (formed/consumed) per unit of time
Introduction Collision Theory
For a reaction to proceed, successful collision must happen!
Collision frequency is directly proportional to reaction rate
Activation energy is inversely proportional to reaction rate
Transition State Theory
– For a reaction to occur, it must reach a transition state (with enough energy) before proceeding on to the products.
Introduction
Factors Affecting Reaction Rates Nature of Reactants Concentration of Reactants Temperature Surface Area Catalyst
RESULTS AND DISCUSSION
Part A. Nature of Reactants
Reactant Observation
MgBubbles (slower
reaction)
NaExplosion (faster
reaction)
Activation Energy (Ea)— the minimum amount of
energy required in order for a reaction to proceed
—Inversely proportional to the rate of reaction
Part A. Nature of Reactants
↓Ea = ↑reactivity = ↑reaction rate
Na has a relatively
lower activation
energy than Mg.
Part A. Nature of Reactants
Part B. Concentration of Reactants
Table 2: Effect of [Na2S
2O
3] on rate with constant [HCl]
[Na2S
2O
3][HCl
]ln
[Na2S
2O
3]
Time(s)
Rate (1/time
)
Ln Rate
0.125 0.5 -2.079 15 0.067 -2.703
0.100 0.5 -2.303 19 0.053 -2.937
0.075 0.5 -2.590 27 0.037 -3.297
0.050 0.5 -2.996 64 0.016 -4.135
0.025 0.5 -3.689 142 7.04 E-3
-4.956
Plot ln rate vs ln [Na2S
2O
3]
Part B. Concentration of Reactants
-5.5 -5 -4.5 -4 -3.5 -3 -2.5
-4
-3.5
-3
-2.5
-2
-1.5
-1
-0.5
0
ln rate
ln [Na
2S
2O
3]
Table 3: Effect of [HCl] on rate with constant [Na2S
2O
3]
[Na2S
2O
3][HCl] ln
[HCl]Time Rate
(1/time)
ln Rate
0.1 1.0 0.000 15 0.067 -2.703
0.1 0.8 -0.233 22 0.043 -3.147
0.1 0.6 -0.511 26 0.038 -3.270
0.1 0.4 -0.916 34 0.029 -3.540
0.1 0.2 -1.609 41 0.024 -3.730
Part B. Concentration of Reactants
Plot ln rate vs ln [HCl]
Part B. Concentration of Reactants
-1.8 -1.6 -1.4 -1.2 -1 -0.8 -0.6 -0.4 -0.2 0
-4
-3.5
-3
-2.5
-2
-1.5
-1
-0.5
0
ln rate
ln [HCl]
Solving for the order with respect to Na
2S
2O
3:
ln (rate = k[Na2S
2O
3]x[HCl]y)
ln Rate = x ln [Na2S
2O
3] + y ln [HCl] + ln K
y = m x + b
r2 = 0.99
m = order with respect to [Na2S
2O
3] =
1.458
b = y ln [HCl] + ln K = 0.373
Part B. Concentration of Reactants
Solving for the order with respect to HCl:
ln (rate = k[Na2S
2O
3]x [HCl]y)
ln Rate = y ln[HCl] + x ln[Na2S
2O
3] + ln k
y = m x + b
r2 = 0.86
m = order with respect to [HCl] = 0.579
b = x ln[Na2S
2O
3] + ln k = -2.901
Part B. Concentration of Reactants
Actual vs. Theoretical Order
Actual Theoretical
Order wrt
Na2S
2O
3
1.458 2
Order wrt HCl
0.579 0
Overall order
2.037 2
Part B. Concentration of Reactants
Rate = k [Na2S
2O
3]1.458 [HCl]0.579
Solving for specific rate constant, k (differential method):
b = y ln [HCl] + ln K
0.373 = (0.579) ln (0.5) + ln k
k = 2.169 M-1 s-1
b = x ln[Na2S2O3] + ln k
-2.901 = (1.458) ln (0.1) + ln k
k = 1.578 M-1 s-1
Part B. Concentration of Reactants
As [Na2S
2O
3 ] increases, the rate of
reaction also increases As [HCl] increases, rate increases. BUT!
Theoretically, it shouldn’t. GENERALLY, ↑ concentration = ↑
collision frequency = ↑ reaction rate
Part B. Concentration of Reactants
Part C. Temperature
Temp (o C)
1/T (K)
Time (s)
Rate (1/tim
e)k ln k
173.45E-
0385 0.012 1.49 0.399
273.34E-
0354 0.019 2.36 0.859
373.23E-
0327 0.037 4.59 1.524
Arrhenius Plot
0.0032 0.00325 0.0033 0.00335 0.0034 0.00345 0.00350
0.2
0.4
0.6
0.8
1
1.2
1.4
1.6
1.8
In k
1/T (K)
Part C. Temperature
ln k = -EaR
1T+ ln A
y = mx + b
Ea = -R(slope)*find slope using linear regressionSlope = -5054.287569Ea = -8.314 J (-5054.287569)Ea = 42,021.35 J
Part C. Temperature
↑temperature = ↑rate of reaction
Increasing temperature increases the fraction of molecules that posses enough kinetic energy to overcome Ea, thus increasing reaction rate.
Fra
ctio
n of
mol
ecul
es
T2 > T1
Part C. Temperature
Part D. Surface Area
Reactants Visible Results
Strip of MgBubbles (slower
reaction)
Pieces of MgBubbles (faster
reaction)
↑surface area = ↑rate of reaction
—A greater surface area exposed increases the probability of effective collisions between reactant molecules and results to an increase in reaction rate.
Part D. Surface Area
Part E. Catalyst
Reactants Visible Results
H2O2 + Rochelle SaltSlow bubble formation
H2O2 + Rochelle Salt
+ CoCl2
Faster bubble formation; color change (pink → green → pink)
- a substance that speeds up reaction by providing an alternative pathway with lower activation energy for the reactant molecules, but is not consumed in the reaction
CoCl2 in aqueous solution is pink
because of [CoCl2 (H2O)4]●2H2O or
cobaltous chloride hexahydrate
formed an intermediate
activated complex of Co3+ (green) and
tartrate ions
converted back to its original form,
CoCl2
Color Change:
pink
pink
green
Conclusion
—5 main factors that can affect the rate of reaction – nature of reactants, concentration of reactants, temperature, surface area and presence of a catalyst.
—For a reaction to proceed, molecules must acquire enough energy to overcome the activation energy.
—Generally, increasing the frequency of collisions and decreasing the activation energy would hasten a reaction.
—Concentration of reactants, temperature and surface area are directly proportional to reaction rate. An increase in these factors would increase reaction rate.
—Adding a catalyst and using reactants with lower activation energy would also hasten the reaction rate.
Conclusion
Recommendations
—Use a stopwatch instead of a clock or watch because of its greater accuracy.
—It is also recommended that this experiment should not be performed spontaneously and without proper preparation (glassware, reagents, personal protective equipment, etc).
References:• Clark, J. (March 2011). Rates of Reaction Menu. Retrieved
April 30, 2011, from
• http://www.chemguide.co.uk/physical/basicratesmenu.html#top
• The Royal Society of Chemistry. Classic Chemistry Demonstration:A Visible Activated Complex. Retrieved May 1, 2011, from
• www.rsc.org/images/oscillating_tcm18-188828.pdf
• Purchon, N.D. (Novermber 10, 2006). Rates of Reaction. Retrieved April 30, 2011, from
• http://www.purchon.com/chemistry/rates.htm
• Engle, Harry and Luciana Ilao, Learning Modules in General Chemistry 2 (2007 Edition),