How to calculate the rate of reaction of Sodium and Chlorine $$\ce{2Na(s) + Cl2(g)   -&gt;    2NaCl(s)}$$
How exactly do you calculate the reaction rate of the following equation? 

Since reaction rate is usually calculated in moles per $\rm dm^3$ per second (M/s) and the reactants here are solid. I want to have a equation of $$R = K \times\ce{[Na]}^x\ce{[Cl]}^y$$
Rate laws are determined via experimental data. Reactant concentrations are held constant while others are modified in order to determine the effect of each reactant's concentration on a reaction. For this reason, one cannot determine the rate law from just a stoichiometric equation.
$$\ce{2Na + Cl2 -&gt; 2NaCl}$$

We can propose reaction mechanisms to make an educated guess of the rate law:
\begin{align}
\ce{Na + Cl2 &amp;-&gt; NaCl + Cl}&amp;&amp;\text{(slow step)}\\
\ce{Na + Cl &amp;-&gt; NaCl}&amp;&amp;\text{(fast step)}\\
\end{align}

Using Hess' law, we can combine these to make sure it equals the original equation.
$$\require{cancel} 
\begin{align}
\ce{  &amp;&amp; Na + Cl2 &amp;-&gt; NaCl + \cancel{Cl}}\\
\ce{+ &amp;&amp; Na + \cancel{Cl} &amp;-&gt; NaCl} \\\hline
\ce{  &amp;&amp; 2Na + Cl2 &amp;-&gt; 2NaCl} \\
\end{align}$$

The coefficients of the slowest elementary step in reaction mechanism determine the exponents of the rate law. Hence:
$$\text{Rate}=k[\ce{Na}]^{1}[\ce{Cl2}]^{1}$$

Or just
$$\text{Rate}=k[\ce{Na}][\ce{Cl2}].$$