OPEN ENDED LAB
Object:
To study the kinetics of the reaction between potassium permanganate and oxalic acid in an
acidic medium by determining the reaction rate, order using differential method.
Apparatus:
Test tubes (7)
Three Burettes (100ml)
Two pipettes (10ml)
Measuring cylinder (10 mL)
Stopwatch
Beaker (50 mL or 100 mL)
Chemicals:
Potassium permanganate (KMnO₄): 0.01 M solution
Oxalic acid (C₂H₂O₄): 0.025 M solution
Sulfuric acid (H₂SO₄): 1 M solution
Distilled water
Theory:
The reaction between potassium permanganate (KMnO₄) and oxalic acid (H₂C₂O₄) in an acidic medium is
a redox reaction.
2 KMnO4+5 C2H2O4+6 H2SO4→2MnSO4+10 CO2+8 H2O+ K2SO4
Permanganate ion (MnO₄⁻) acts as a strong oxidizing agent, while oxalic acid acts as a reducing agent. The
overall reaction is:
2MnO4− +5C2O42− +16H+ → 2Mn2+ +10CO2+8H2O
The color of KMnO₄ will change from purple to colorless as it is reduced to Mn²⁺.
Bubbling may be observed due to CO₂ formation.
The reaction rate law describes how the reaction rate depends on the concentrations of the
reactants. For the reaction between potassium permanganate (KMnO4) and oxalic acid (C2H2O4)
in the presence of sulfuric acid (H2SO4), the rate is expressed as: [1]
Rate=k[KMnO4]p[C2H2O4]q[H2SO4]r
Here, p, q and r are the reaction orders with respect to each reactant, and (k) is the rate constant.
The differential method is used to determine the reaction order by analyzing the initial rate of the
reaction as a function of the initial concentrations of the reactants. By performing a series of
experiments with different initial concentrations and measuring the initial rates, we can determine
the values of p, q, and r. To evaluate (dx/dt) an appropriate method is used, the change 'dx' in 'x' is
measured over an appreciable time interval '∆t' and (∆x/∆t) is equal to (dx/dt) corresponding to the
mean value of 'x' in that time interval.[2]
Rate= -∆[KMnO4]/∆T or Rate= -∆[C2H2O4]/∆T or Rate= -∆[H2SO4]/∆T
As the reaction proceeds, the permanganate ion (MnO4-) is reduced to Mn2+ ion. The added
reducing agent reacts with the MnO4- ion, preventing it from oxidizing the indicator. Once all the
reducing agent is consumed, the excess MnO4- ion reacts with the indicator, causing a color
change. The time taken for this color change to occur is recorded.[1]
By knowing the initial concentration of the reducing agent and the amount consumed during the
reaction, we can calculate the change in concentration of MnO4- ion (Δx). Since the concentration
of MnO4- ion was initially zero and becomes Δx at the time of the color change, the average
concentration of MnO4- ion during the time interval Δt can be approximated as Δx/2.
And the total order ‘n’ will be: n = p + q + r [1]
Safety Precautions:
Handling chemicals such as potassium permanganate (KMnO₄), oxalic acid (C₂H₂O₄), and sulfuric
acid (H₂SO₄) requires strict adherence to safety protocols due to their hazardous nature. Potassium
permanganate is a strong oxidizing agent and can cause skin irritation, while oxalic acid is toxic
and corrosive. Sulfuric acid is highly corrosive and can cause severe burns. Mandatory Personal
Protective Equipment (PPE) includes lab coats, safety goggles, and chemical-resistant gloves to
protect against splashes and contact with the skin or eyes.
Procedure:
Take 100ml each of sulfuric acid, oxalic acid and distilled water in separate burettes and keep the
potassium permanganate in a separate beaker. Prepare 7 test tubes, containing the volume of the
reactants mentioned against test tube number in the following table noting that the total volume of
the reactants in each test tube will amount to 12ml.
Table. 1
Test KMnO4 C2H2O4.2H2O H2SO4 Distilled
tube (ml) (ml) (ml) water (ml)
1 4 4 4 0
2 3 4 4 1
3 2 4 4 2
4 4 3 4 1
5 4 2 4 2
6 4 4 3 1
7 4 4 2 2
Observations:
Table: 2 Experiment time
Test tubes 1 2 3 4 5 6 7
Time (sec) 663 433 413 338 208 160 128
Calculations:
For ∆x:
KMnO4 = C2H2O4. 2H2O
M1V1 = M2V2
n1 n2
∆x= M1= n1. M. V2 = (0.025) (4) (2) = 0.0033 M
V1. n2 (12) (5)
Final concentrations of KMnO4, H2SO4 and oxalic acid:
Table. 3
Conc “a” (M) Conc “b” (M) Conc “c” (M)
0.01 1 0.033
0.0075 1 0.033
0.005 1 0.033
0.01 1 0.025
0.01 1 0.0167
0.01 0.75 0.033
0.01 0.5 0.033
For “p” conc. of KMnO4
Table. 4
Test tubes Volume (ml) Conc “a” (M) Time (sec)
1 4 0.01 663
2 3 0.0075 433
3 2 0.005 413
p1 = ln 0.00333 - ln 0.00333
663 443
ln (0.0075- 0.00333) – ln (0.01- 0.00333)
p1 = 0.858
Similarly,
p2 = 0.0766
p3 = 0.342
Avg “p” = 0.858+ 0.0766+ 0.342 = 0.4255
3
For “q” Conc. of H2SO4
Table. 5
Test tubes Volume (ml) Conc “c” (M) Time (sec)
5 4 1 208
6 3 0.75 160
7 2 0.5 128
q1 = ln 0.00333 - ln 0.00333
208 160
ln (0.75- 0.00333) – ln (1- 0.00333)
q1 = 0.902
Similarly,
q2 = 0.6
q3 = 0.7
Avg “q” = 0.902+ 0.6+ 0.7 = 0.734
3
For “r” conc. of oxalic acid:
Table. 6
Test tubes Volume (ml) Conc “b” (M) Time (sec)
3 4 0.033 413
4 3 0.025 338
5 2 0.0167 208
r1 = ln 0.00333 - ln 0.00333
413 338
ln (0.025- 0.00333) – ln (0.033- 0.00333)
r1 = 0.637
Similarly,
r2 = 1.005
r3 = 0.8605
Avg “r” = 0.637+ 1.005+ 0.8605 = 0.8345
3
For order of reaction:
n = p+ q+ r
n = 0.4255+ 0.734+ 0.8345
n= 2
Discussion:
The aim of this experiment was to determine the reaction order for the oxidation of oxalic acid
(C₂H₂O₄) by potassium permanganate (KMnO₄) in an acidic medium (H₂SO₄) using the differential
method. The results obtained from the experiment indicate that the reaction follows a second-order
kinetic model (n=2), with partial reaction orders determined as p=0.42 for KMnO₄, q = 0.73 for
sulfuric acid, and r=0.85r = 0.85 for oxalic acid. These fractional reaction orders suggest that the
reaction mechanism is more complex than a simple integer-order dependence.
The fractional order for KMnO₄ (p=0.42) implies that the reaction rate does not proportionally
increase with the KMnO₄ concentration, suggesting a possible saturation effect or a step in the
mechanism involving an intermediate that regulates KMnO₄ utilization. The order with respect to
sulfuric acid (q=0.73) indicates its significant influence on the reaction rate, likely by stabilizing
the Mn+2 intermediate or facilitating electron transfer. The near-unit order with respect to oxalic
acid (r = 0.83) shows that the rate is closely tied to oxalic acid concentration, consistent with its
role as the reducing agent in the reaction.
Despite some deviations from whole numbers, the overall reaction order (n=p+q+r=2) aligns with
theoretical expectations for this redox reaction. The variation in time taken for the purple color of
KMnO₄ to disappear across the test tubes confirmed the dependence of the rate on the reactant
concentrations. However, the formation of a brown precipitate in some cases suggested that side
reactions, such as MnO2formation, may have slightly impacted the observed rates.
These results highlight the intricacies of reaction kinetics, where reaction orders often reflect a
combination of the molecular mechanism and intermediate stabilization. Minor experimental
errors, such as slight variations in timing or concentration measurements, could also have
contributed to the fractional values obtained.
Conclusion:
The experiment successfully determined the reaction order for the oxidation of oxalic acid by
KMnO₄ in the presence of H₂SO₄. The overall reaction was found to be second order (n=2) with
partial orders p=0.42, q= 0.73, and r=0.83 for KMnO₄, sulfuric acid, and oxalic acid, respectively.
These results indicate a complex reaction mechanism with dependencies on all three reactants.
The fractional orders highlight the importance of intermediates and catalytic effects in this
reaction. This experiment underscores the utility of the differential method in studying reaction
kinetics and provides valuable insights into the dynamics of redox processes in acidic media.
Further investigations could explore the mechanistic details to better understand the origins of
these fractional orders. [2]
Reference:
[1] Rate Law Determination ", Lab Manual University of Rhode Island Land-grant university in
South Kingstown, Rhode Island
[2] Chemical Reaction Engineering 3rd Edition by Octave Levenspiel
