## 15.39b

$aR \to bP, Rate = -\frac{1}{a} \frac{d[R]}{dt} = \frac{1}{b}\frac{d[P]}{dt}$

Angela 1K
Posts: 80
Joined: Fri Sep 29, 2017 7:05 am

### 15.39b

The question asks to find the amount of time it would take in the reaction of A --> 2B + C, when [A]0 = 0.15M, for the concentration of B to increase to 0.19M, given that k = 0.0035L/mol*min in the rate law for the loss of A.

I was wondering why the solutions manual solves this equation as a second-order reaction, when it's only a first order reaction (unless I'm mistaken)

Yashaswi Dis 1K
Posts: 56
Joined: Fri Sep 29, 2017 7:04 am

### Re: 15.39b

It's because it is specified in the question that the two following equations will be second-order so you just have to use the second-order integrated rate law because it explicitly says that: 1/[A]t = 1/[A]0 + kt.

Manipulate to solve for t and you should get 3.3 x 103 min.

Hope that helps!

Curtis Tam 1J
Posts: 105
Joined: Thu Jul 13, 2017 3:00 am

### Re: 15.39b

You can't tell if it's first order just by looking at the equation. You can only do so with individual elementary step reactions but not with the total reaction. The order for the net reaction is always experimentally determined which is why it must be given in the problem.

Curtis Tam 1J
Posts: 105
Joined: Thu Jul 13, 2017 3:00 am

### Re: 15.39b

looking at the reaction***

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