Laboratory Experiment: Investigating Chemical Kinetics

Categories: Chemistry

Chemical kinetics is the study of rates of chemical reactions, influenced by various factors like reactant nature, concentration, temperature, and catalyst presence. This experiment aims to elucidate the application of chemical kinetics by determining the rate law, employing the Arrhenius equation for activation energy, investigating catalyzed reactions, and examining the impact of temperature and concentration on reaction rates. Hydrochloric acid (HCl) reacting with sodium thiosulfate (Na2S2O3) in water serves as the model system for these investigations.

Experimental Setup:

Six setups were devised to test the reaction rate under varying conditions.

The factors explored included concentration, temperature, and catalyst presence. Initial concentrations were calculated, and rates were determined by measuring the inverse of time.

Results and Discussion:

Part 1: Effect of Concentration on Reaction Rate:

  • The reaction order with respect to S2O3^2- and H+ were found to be 1.48 and 0.29, respectively, resulting in an overall kinetic order of 1.77.
  • The visible change was the formation of sulfur (S), indicating reaction occurrence.

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  • Possible errors included inconsistent reaction timing and variations in beaker size.

Part 2: Effect of Temperature on Reaction Rate:

  • Results indicated a direct proportionality between temperature and reaction rate.
  • Increasing temperature increased collision frequency, leading to faster reactions.
  • Activation energy (Ea) sign (positive) suggested a temperature rise always increased rate constant.

Part 3A: Oxidation of Tartrate by Hydrogen Peroxide:

  • The equation involved the reaction of H2O2 with KNaC4H4O6.
  • Effervescence indicated CO2 formation, catalyzed by CoCl2, with regeneration evidenced by the return of a pink color.

Part 3B: Reaction of Oxalate with Permanganate:

  • The equation involved the reaction of KMnO4 with Na2C2O4 in the presence of H2SO4.

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  • The addition of KMnO4 drops influenced the reaction rate, indicating the role of the catalyst.

Calculations and Formulas:

  • Initial Concentration: Calculated by multiplying substance concentration by volume.
  • Rate Determination: Derived as the inverse of time.
  • Kinetic Order: Determined based on reaction orders with respect to reactants.
  • Arrhenius Equation: Relates rate constant (k), activation energy (Ea), and temperature (T).


Table 1: Experimental Results (Concentration) InsertTable1herewithRunNo.,Time(s),1/Time(s−1),[S2O32−]initial,[H+]initial

Table 2: Experimental Results (Temperature) InsertTable2herewithTemp.(K),1/Temp.(K−1),Time(s),1/time(s−1)

Table 3: Observation (Oxidation of Tartrate) InsertTable3herewithTestTube,Observation,RelativeReactionTime

The experiment revealed that concentration, temperature, and catalyst presence significantly affect reaction rates. The kinetic order, Arrhenius equation, and observations from catalyzed reactions provided valuable insights. Understanding these factors aids in predicting chemical reaction behavior and enhancing efficiency.

This comprehensive laboratory investigation contributes to a deeper comprehension of chemical kinetics, providing a foundation for further explorations in reaction mechanisms and optimization.

Updated: Feb 29, 2024
Cite this page

Laboratory Experiment: Investigating Chemical Kinetics. (2024, Feb 29). Retrieved from

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