In a first order reaction, $75 %$ of the reactants disappeared in 1.386 hr. What is the rate constant?

3.6 \times 10^{- 3} s^{- 1}

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2.7 \times 10^{- 4} s^{- 1}

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72 \times 10^{- 3} s^{- 1}

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1.8 \times 10^{- 3} s^{- 1}

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Solution

The correct option is D $2.7 \times 10^{- 4} s^{- 1}$
$t_{75 %} = 2 t_{1 / 2}$
$∴ t_{1 / 2} = \frac{1.386 h r}{2} = 0.693 h r$
$k = \frac{0.693}{t_{1 / 2}} = \frac{0.693}{0.693} = 1 h r^{- 1} = \frac{1}{3600 s}$

$= 2.7 \times 10^{- 4} s^{- 1}$

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In a first order reaction, 75% of the reactants disappeared in 1.386 hr. What is the rate constant?

The correct option is D 2.7×10−4s−1t75%=2t1/2∴t1/2=1.386hr2=0.693hrk=0.693t1/2=0.6930.693=1hr−1=13600s=2.7×10−4s−1

In a first order reaction, $75 %$ of the reactants disappeared in 1.386 hr. What is the rate constant?

The correct option is D $2.7 \times 10^{- 4} s^{- 1}$
$t_{75 %} = 2 t_{1 / 2}$
$∴ t_{1 / 2} = \frac{1.386 h r}{2} = 0.693 h r$
$k = \frac{0.693}{t_{1 / 2}} = \frac{0.693}{0.693} = 1 h r^{- 1} = \frac{1}{3600 s}$

$= 2.7 \times 10^{- 4} s^{- 1}$