Recall that the **rate law** only focuses on the reactant concentrations and has a general form of:

$\overline{){\mathbf{Rate}}{\mathbf{=}}{\mathbf{}}{\mathbf{k}}{\left[\mathbf{A}\right]}^{{\mathbf{x}}}{\left[\mathbf{B}\right]}^{{\mathbf{y}}}}\phantom{\rule{0ex}{0ex}}\mathbf{Rate}\mathbf{=}\mathbf{}\mathbf{k}{\left[{\mathrm{SO}}_{2}{\mathrm{Cl}}_{2}\right]}^{\mathbf{n}}$

Finding n using experiment 1 and 2:

$\frac{\mathbf{1}\mathbf{.}\mathbf{60}\mathbf{x}{\mathbf{10}}^{\mathbf{-}\mathbf{6}}}{\mathbf{3}\mathbf{.}\mathbf{2}\mathbf{x}{\mathbf{10}}^{\mathbf{-}\mathbf{6}}}\mathbf{=}\frac{\mathbf{k}{(0.1)}^{\mathbf{n}}}{\mathbf{k}{(0.2)}^{\mathbf{n}}}\phantom{\rule{0ex}{0ex}}\frac{\mathbf{1}}{\mathbf{2}}\mathbf{=}{\left(\frac{1}{2}\right)}^{\mathbf{n}}\phantom{\rule{0ex}{0ex}}\mathbf{n}\mathbf{}\mathbf{=}\mathbf{}\mathbf{1}$

Consider this initial-rate data at a certain temperature for the reaction described by

SO_{2}Cl_{2}(g) → SO_{2}(g) + Cl_{2}(g)

Determine the value and units of the rate constant.

[SO | Initial rate (M/s) |

0.1 | 1.60 x 10 |

0.2 | 3.20 x 10 |

0.3 | 4.80 x 10 |

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