54) The following graph shows three different stresses on the reaction. For each stress, numbered 1, 2, and 3 on the graph, identify the stress that has been applied to the system and using Le Chateliers principle, explain why the graph changes the way it does.
1) At the point marked #1, all three concentrations gradually increase, showing that heat is added. The addition of heat puts stress on the system, so the equilibrium shifts to the right to partially undo “stress” and produce more products.
2) At the point marked #2, the concentrations of NO, Cl2 and NOCl suddenly increase. We can assume that the container volume decreased and there is an increase in pressure. In response, the chemical system will shift in a direction that will cause …show more content…
The product has been removed and the reaction shifts to the right. Since more products are needed to return to equilibrium, the concentration of product increases and the concentration of reactants decreases.
55) Use Le Chatelier’s principle to describe and explain the shifts in the following equilibrium reactions:
i) As volume is increased, the concentration of the gas decreases, and the pressure decreases. The equilibrium will not shift because there is equal number of moles on each side of the equation. ii) The equilibrium will shift to the right as Pb(NO3)2 is added. Shifting to the right reduces the concentration of excess reactant caused by adding Pb(NO3)2. iii) Decreasing the temperature removes the product (heat) from the system and the equilibrium will shift to the right. iv) Since the catalyst speeds up the rate of forward and backward reaction equally, there is no shift in equilibrium.
b) The Le Chatelier’s principle can explain why haze is more common in the summer. As winter air loses heat (the product), the equilibrium shifts to the right and favours the forward reaction, producing more N2O4 (colourless). On the other hand, as the summer air is composed of more heat (the product), the reverse action is favoured and produces more
1) At the point marked #1, all three concentrations gradually increase, showing that heat is added. The addition of heat puts stress on the system, so the equilibrium shifts to the right to partially undo “stress” and produce more products.
2) At the point marked #2, the concentrations of NO, Cl2 and NOCl suddenly increase. We can assume that the container volume decreased and there is an increase in pressure. In response, the chemical system will shift in a direction that will cause …show more content…
The product has been removed and the reaction shifts to the right. Since more products are needed to return to equilibrium, the concentration of product increases and the concentration of reactants decreases.
55) Use Le Chatelier’s principle to describe and explain the shifts in the following equilibrium reactions:
i) As volume is increased, the concentration of the gas decreases, and the pressure decreases. The equilibrium will not shift because there is equal number of moles on each side of the equation. ii) The equilibrium will shift to the right as Pb(NO3)2 is added. Shifting to the right reduces the concentration of excess reactant caused by adding Pb(NO3)2. iii) Decreasing the temperature removes the product (heat) from the system and the equilibrium will shift to the right. iv) Since the catalyst speeds up the rate of forward and backward reaction equally, there is no shift in equilibrium.
b) The Le Chatelier’s principle can explain why haze is more common in the summer. As winter air loses heat (the product), the equilibrium shifts to the right and favours the forward reaction, producing more N2O4 (colourless). On the other hand, as the summer air is composed of more heat (the product), the reverse action is favoured and produces more