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The decomposition of N_{2}O_{5} in CCl_{4} at 318K has been studied by monitoring the concentration of N_{2}O_{5} in the solution. Initially, the concentration of N_{2}O_{5} is 2.33 mol L^{–1} and after 184 minutes, it is reduced to 2.08 mol L^{–1}. The reaction takes place according to the equation

2 N_{2}O_{5} (g) → 4 NO_{2} (g) + O_{2} (g)

The rate of production of NO_{2} during this period is-

1. 5.72 × 10^{–3} mol L^{–1} min^{–1}

2. 2.72 × 10^{–3} mol L^{–1} min^{–1}

3. 1.72 × 10^{–5} mol L^{–1} min^{–1}

4. 6.72 × 10^{–4} mol L^{–1} min^{–1}

Subtopic: Definition, Rate Constant, Rate Law |

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In a first order reaction, time required for completion of 99.9% is X times of half-life (t_{1/2}) of the reaction. When reaction is completed 99.9%, [R]_{n} = [R]_{0} – 0.999[R]_{0} .The value of X is-

1. 5

2. 10

3. 15

4. 20

Subtopic: First Order Reaction Kinetics |

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If 75 % of a first-order reaction was completed in 90 minutes, 60 % of the same reaction would be completed in approximately (in minutes):

(Take : log 2 = 0.30 ; log 2.5 = 0.40)

1. 50 min

2. 60 min

3. 70 min

4. 65 min

Subtopic: First Order Reaction Kinetics |

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The rate of a reaction is decreased by 3.555 times when the temperature was changed from 40°C to 30°C. The activation energy (in kJ ${\mathrm{mol}}^{-1}$) of the reaction is:

(Take R=8.314 J ${\mathrm{mol}}^{-1}{\mathrm{K}}^{-1}$ In 3.555=1.268)

1. 100 kJ/mol

2. 120 kJ/mol

3. 95 kJ/mol

4. 108 kJ/mol

Subtopic: Arrhenius Equation |

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For the non – stoichiometry reaction 2A + B → C + D, the following kinetic data were obtained in three separate experiments (all at 298 K).

Initial Concentration (A) | Initial Concentration (B) | Initial rate of formation of C (mol L^{–} S^{–}) |

0.1 M 0.1 M 0.2 M |
0.1 M 0.2 M 0.1 M |
1.2 × 10^{–3}1.2 × 10 ^{–3}2.4 × 10 ^{–3} |

The rate law for the formation of C is:

1. $\frac{\mathrm{dc}}{\mathrm{dt}}=\mathrm{k}\left[\mathrm{A}{]}^{2}\right[\mathrm{B}]$

2. $\frac{\mathrm{dc}}{\mathrm{dt}}=\mathrm{k}\left[\mathrm{A}\right][\mathrm{B}{]}^{2}$

3. $\frac{\mathrm{dc}}{\mathrm{dt}}=\mathrm{k}\left[\mathrm{A}\right]$

4. $\frac{\mathrm{dc}}{\mathrm{dt}}=\mathrm{k}\left[\mathrm{A}\right]\left[\mathrm{B}\right]$

Subtopic: Order, Molecularity and Mechanism |

76%

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The time for the half-life period of a certain reaction A → Products is 1 hour. When the initial concentration of the reactant 'A' is 2.0 mol L^{-1}, the time taken for its concentration to come from 0.50 to 0.25 mol L^{-1},if it is a zero-order reaction, is:

1. 1h

2. 4 h

3. 0.5 h

4. 0.25 h

Subtopic: Order, Molecularity and Mechanism |

66%

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Consider the reaction, 2A + B → Products.

When concentration of B alone was doubled, the half-life did not change. When the concentration of A alone was doubled, the rate increased by two times. The unit of rate constant for this reaction is:

1. L mol^{–1} s^{–1}

2. no unit

3. mol L^{–1}s^{–1}

4. s^{–1}

Subtopic: Order, Molecularity and Mechanism |

From NCERT

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The rate equation for the reaction 2A+B→C is found to be:

rate = k [A] [B]

The correct statement in relation to this reaction is that the:

1. | Unit of k must be s^{-1} |

2. | t_{1/2} is a constant |

3. | Rate of formation of C is twice the rate of disappearance of A |

4. | Value of k is independent of the initial concentrations of A and B |

Subtopic: Definition, Rate Constant, Rate Law |

54%

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For a first-order reaction A → B the reaction rate at a reactant concentration of 0.01M is found to be $2.0\times {10}^{-5}\text{mole}{\mathrm{L}}^{-1}{\text{}\mathrm{s}}^{-1}$. The half-life period of the reaction is:

1. | 300s | 2. | 30s |

3. | 220s | 4. | 347s |

Subtopic: First Order Reaction Kinetics |

84%

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Half-life of substance A following first order kinetics is 5 days. Starting with 100g of A, the amount left after 15 days will be:

1. 25 g

2. 50 g

3. 12.5 g

4. 6.25 g

Subtopic: First Order Reaction Kinetics |

86%

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