$\overline{){\mathbf{rate}}{\mathbf{}}{\mathbf{law}}{\mathbf{=}}{\mathbf{k}}{\left[\mathbf{A}\right]}^{{\mathbf{x}}}{\left[\mathbf{B}\right]}^{{\mathbf{y}}}}$

order, x with respect to **A**.

- Use values where
**A changes**while**B remains the same**.

**Solve for x:**

$\frac{\mathbf{rate}\mathbf{}\mathbf{2}}{\mathbf{rate}\mathbf{}\mathbf{1}}\mathbf{=}\frac{{\displaystyle {\left[\mathbf{A}\right]}^{\mathbf{x}}\mathbf{}\mathbf{at}\mathbf{}\mathbf{rate}\mathbf{}\mathbf{2}}}{{\displaystyle {\left[\mathbf{A}\right]}^{\mathbf{x}}\mathbf{}\mathbf{at}\mathbf{}\mathbf{rate}\mathbf{}\mathbf{1}}}$

*x = order*

*(Larger concentration should be on the numerator)*

$\frac{\mathbf{0}\mathbf{.}\mathbf{07}}{\mathbf{0}\mathbf{.}\mathbf{035}}\mathbf{=}\frac{{\left[\mathbf{0}\mathbf{.}\mathbf{10}\right]}^{\mathbf{x}}}{{\left[\mathbf{0}\mathbf{.}\mathbf{05}\right]}^{\mathbf{x}}}\phantom{\rule{0ex}{0ex}}\mathbf{2}\mathbf{=}{\mathbf{2}}^{\mathbf{x}}$

***2 ^{1} = 2*

*x = 1 → 1 ^{st}*

For the reaction 2A + B → C, the initial rate was measured at several different reactant concentrations. From the resulting tabulated data, determine the rate law for the reaction.

[A] (M) | [B] (M) | Initial Rate (M/s) |

0.05 | 0.05 | 0.035 |

0.10 | 0.05 | 0.070 |

0.20 | 0.10 | 0.56 |

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