## definitions

$\frac{d[R]}{dt}=-k; [R]=-kt + [R]_{0}; t_{\frac{1}{2}}=\frac{[R]_{0}}{2k}$

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405112316
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Joined: Fri Sep 28, 2018 12:17 am

### definitions

Can someone define zero, first, second order reactions? I thought it had to do with how many components are coming together, but that wouldn't make sense for a zero order rxn.

Chem_Mod
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### Re: definitions

It refers to the exponents on the concentrations. We designate these orders usually in terms of each reactant. So, in something like rate = k[A]0, reaction rate is independent of concentration of A itself. Thus, it's zero order with respect to A. Same idea for other orders.

Kevin ODonnell 2B
Posts: 62
Joined: Fri Sep 28, 2018 12:24 am

### Re: definitions

I am also led to believe that these values have to be experimentally found, and are not just based off of something you can look at a chemical equation and tell. For example, if you have 1 mol of x that produces 1 mol of product then you have 2 mol of x that produces 2 mol of product you can see this increases by an exponent of 1 so it is of first-order. If you have 1 mol of y that produces 1 mol of product and then have 2 mol of y that produces 4 mol of product then you can see this increases by an exponent of 2 so it is of second-order. Lastly, if you have 1 mol of z that produces 1 mol of product and then you have 2 mol of z that produces 1 mol of product then you can see this increases by an exponent of 0 since it instead just relies on the rate constat thus it is of zero-order.

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