Two test tubes, one hot and the other cold, the test tube surrounded by ice contains dinitrogen tetraoxide and the test tube chillin’ in the hot bath contains nitrogen dioxide.
NO2<->N2O4, -58.8 delta H for the production of N2O4. This is a reversible reaction. The production of N2O4 is an exothermic process and the production of NO2 is an endothermic process. Putting test tube in the hot bath added energy to the system favoring the production of NO2 gas giving the orange look. Putting the test tube in an ice bath released energy to the surroundings favoring the production of N2O4 gas giving the colorless look. Both gases are present in both instances but one gas is in higher concentration as it’s being produced at a faster rate. In the instance with the hot test tube the rate of production of NO2 is greater than that of N2O4 hence a greater orange glow, N2O4 is still produced but at a slower rate. This is true for the cold test tube but the reverse with the gases. The N2O4 is at a higher concentration as it is being produced at a higher rate while the NO2 is being produced at a much lower rate. At a neutral temperature neither gas is favored more than the other to be produced as both the delta G’s for their reactions is zero. This is equilibrium when the rates of production equal each other so that products of both gases is relatively the same as the reactants that started the reaction. By changing the temperature of the surroundings and favoring a certain enthalpy change can cause the equilibrium to be disrupted and cause one rate to become higher than the other rate.
Le Chatelier’s principle states that a change to concentration of one of the reactants, pressure of the system, or the temperature of the system will cause a shift in equilibrium to counteract these changes. This principle is observed with the changing rates to counteract the change of temperature of the system.