Table of contents

1. Intro to General Chemistry3h 48m
2. Atoms & Elements4h 16m
3. Chemical Reactions4h 10m
4. BONUS: Lab Techniques and Procedures1h 38m
5. BONUS: Mathematical Operations and Functions48m
6. Chemical Quantities & Aqueous Reactions3h 53m
7. Gases3h 49m
8. Thermochemistry2h 37m
9. Quantum Mechanics2h 58m
10. Periodic Properties of the Elements3h 9m
11. Bonding & Molecular Structure3h 29m
12. Molecular Shapes & Valence Bond Theory1h 57m
13. Liquids, Solids & Intermolecular Forces2h 23m
14. Solutions3h 1m
15. Chemical Kinetics2h 53m
16. Chemical Equilibrium2h 29m
17. Acid and Base Equilibrium5h 1m
18. Aqueous Equilibrium4h 47m
19. Chemical Thermodynamics1h 50m
20. Electrochemistry2h 42m
21. Nuclear Chemistry2h 36m
22. Organic Chemistry5h 7m
23. Chemistry of the Nonmetals2h 39m
24. Transition Metals and Coordination Compounds3h 16m

15. Chemical Kinetics

Arrhenius Equation

General Chemistry

15. Chemical Kinetics

Arrhenius Equation

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Multiple Choice

The activation energy for the decomposition of HI(g) to H2(g) and I2(g) is 186 kJ/mol. The rate constant at 555 K is 3.52 × 10^-7 L/mol-s. What is the rate constant at 645 K?

4.56 × 10^-6 L/mol-s

1.23 × 10^-6 L/mol-s

2.89 × 10^-7 L/mol-s

5.67 × 10^-8 L/mol-s

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Verified step by step guidance

Start by understanding the Arrhenius equation, which relates the rate constant (k) to temperature (T) and activation energy (Ea): k = A * exp(-Ea / (R * T)), where A is the pre-exponential factor and R is the gas constant (8.314 J/mol-K).

Since we are given two temperatures and their corresponding rate constants, we can use the two-point form of the Arrhenius equation: ln(k2/k1) = (-Ea/R) * (1/T2 - 1/T1). This allows us to find the rate constant at the second temperature (T2) using the known rate constant at the first temperature (T1).

Substitute the given values into the equation: k1 = 3.52 × 10^-7 L/mol-s at T1 = 555 K, Ea = 186 kJ/mol (convert to J/mol by multiplying by 1000), and T2 = 645 K.

Calculate the difference in the reciprocal temperatures: (1/T2 - 1/T1). This involves finding 1/645 and 1/555, then subtracting the two values.

Use the calculated difference and the activation energy to find ln(k2/k1). Rearrange the equation to solve for k2: k2 = k1 * exp((-Ea/R) * (1/T2 - 1/T1)). This will give you the rate constant at 645 K.