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### 15.39b

Posted: **Mon Mar 05, 2018 9:27 pm**

by **Angela 1K**

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)

### Re: 15.39b

Posted: **Mon Mar 05, 2018 9:34 pm**

by **Yashaswi Dis 1K**

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 10^{3} min.

Hope that helps!

### Re: 15.39b

Posted: **Mon Mar 05, 2018 10:11 pm**

by **Curtis Tam 1J**

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.

### Re: 15.39b

Posted: **Mon Mar 05, 2018 10:12 pm**

by **Curtis Tam 1J**

looking at the reaction***