How does Le Chatelier's principle apply to a gas phase reaction at constant volume?
Let's see your formula (with the difference of enthalpy being negative):

$$\ce{2CO (g) + O2 (g) &lt;=&gt; 2CO2 (g)}\qquad(\Delta H^\circ&lt;0)$$

At constant temperature, for Le Chatêlier purposes,


Adding $\ce{CO}$ produces that it react with $\ce{O2}$ to produce more $\ce{CO2}$ to counteract the effect of the applied change and a new equilibrium is established. 
Removing $\ce{CO2}$ produce that $\ce{CO}$ react with $\ce{O2}$ to produce more $\ce{CO2}$.


When you are adding or removing a reagent or product, and all components are in a gaseous state, you are changing the pressure and you can study the reaction by $K_p=\dfrac{P^2_\ce{CO2}}{P^2_\ce{CO}P_\ce{O2}}$.

We also have $K_p=K_c$ $RT^{Δn}$: 


$K_p$ is the equilibrium constant written with partial pressure of gases.
$K_c$ is the constant but being written by concentration of substances.
$R$ is universal gas constant.
$T$ is the absolute temperature (Kelvin degrees).
$\Delta n$ are moles $(\Delta n=n_{\text{products}} - n_{\text{reagents}})$.


If you change concentration of $\ce{CO}$ you are changing the pressure it, too.

But you can change also the pressure, by changing the volume. (For Boyle's Law: the absolute pressure exerted by a given mass of an ideal gas is inversely proportional to the volume it occupies if the temperature and amount of gas remain unchanged within a closed system). 


For example, if you increase the pressure, volume will decrease and the system will shift towards it will be less molecules to counteract the change. In this case, to products. If you in decrease pressure, volume will increase and the system will shift towards reagents.

Other option is to change the pressure by adding or removing an inert gas, but since it don't intervenes in the reaction, the equilibrium is not affected.