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Q. No.In a reaction, A + B → Product, the rate is doubled when the concentration of B is doubled, and the rate increases by a factor of 8, when the concentrations of both the reactants (A and B) are doubled. The rate law for the reaction can be written as:
1. Rate = k[A][B]2
2. Rate = k[A]2[B]2
3. Rate = k[A][B]
4. Rate = k[A]2[B]
| 1. | \(\dfrac{-\Delta[\mathrm{H}I]}{\Delta t}=\dfrac{2 \Delta\left[\mathrm{H}_2\right]}{\Delta t}\) | 2. | \(\dfrac{-\Delta[\mathrm{HI}]}{\Delta t}=\dfrac{4\Delta\left[\mathrm{I}_2\right]}{\Delta t}\) |
| 3. | \(\dfrac{-\Delta[\mathrm{HI}]}{\Delta t}=\dfrac{4 \Delta\left[\mathrm{H}_2\right]}{\Delta t}\) | 4. | \( \dfrac{-\Delta[\mathrm{HI}]}{\Delta t}=\dfrac{\Delta\left[\mathrm{H}_2\right]}{\Delta t}\) |
The rate equation of a reaction is expressed as, Rate = \(k(P_{CH_{3}OCH_{3}})^{\frac{3}{2}}\)
(Unit of rate = bar min–1)
The units of the rate constant will be:
1. bar1/2 min
2. bar2 min–1
3. bar–1 min–2
4. bar–1/2 min–1
The decomposition of NH3 on a platinum surface is a zero-order reaction. The rates of production of N2 and H2 will be respectively:
(given ; k = 2.5 × 10–4 mol–1 L s–1 )
| 1. | 2.5 × 10−4 mol L−1 s−1 and 5.5 × 10−4 mol L−1 s−1
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| 2. | 2.5 × 10−4 mol L−1 s−1 and 7.5 × 10−4 mol L−1 s−1
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| 3. | 1.5 × 10−4 mol L−1 s−1 and 4.5 × 10−4 mol L−1 s−1
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| 4. | 0.5 × 10−4 mol L−1 s−1 and 3.5 × 10−4 mol L−1 s−1 |
At high pressure the following reaction is zero order-
The correct statements among the following is:
| (a) | Rate of reaction = Rate constant |
| (b) | Rate of the reaction depends on the concentration of ammonia |
| (c) | Rate of decomposition of ammonia will remain constant until ammonia disappears completely |
| (d) | Further increase in pressure will change the rate of reaction |
| 1. | (a, b, c) | 2. | (b, c, d) |
| 3. | (a, c, d) | 4. | (a, b, d) |
For a zero-order reaction, the initial amount of reaction is 20 g and half-life is 30 minutes. The amount of reactant left after 60 minutes would be:
1. 5 g
2. 10 g
3. 2.5 g
4. Zero
NO2 required for a reaction is produced by the decomposition of N2O5 in CCl4 as per the equation,
2N2O5(g) 4NO2(g) + O2(g)
The initial concentration of N2O5 is 3.00 mol L–1 and it is 2.75 mol L–1 after 30 minutes. The rate of formation of NO2 is:
1. 2.083 × 10–3 mol L–1 min-1
2. 8.333 × 10–3 mol L–1 min-1
3. 4.167 ×10–3 mol L–1 min-1
4. 1.667 × 10–2 mol L–1 min-1
\(2 \mathrm{C}_2 \mathrm{H}_6(\mathrm{g})+7 \mathrm{O}_2(\mathrm{g}) \rightarrow 4 \mathrm{CO}_2(\mathrm{g})+6 \mathrm{H}_2 \mathrm{O}(l) \),
the rate of disappearance of \(C_2H_6 (g)\):
| 1. | equals the rate of disappearance of \(O_2 (g)\) |
| 2. | is seven times the rate of disappearance of \(O_2 (g)\) |
| 3. | is twice the rate of appearance of \(CO_2 (g)\) |
| 4. | is one-third the rate of appearance of \(H_2O(l)\) |
The units of rate constant and rate of reaction are same for:
1. First order reaction
2. Second order reaction
3. Third order reaction
4. Zero order reaction
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Q. No.
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