Understanding Reversible Systems
For reversible reaction systems, there is ALWAYS (emphasis: ALWAYS!) two processes occurring. There is the process in which reactants turn into products. That's not unusual; every Chemistry student expects that. We refer to this process as the
forward reaction. But at the same time, there is a second process taking place. Product particles are turning into reactant particles. The
reverse reaction is taking place.
When you mix up a batch of reactants and products (Yes. You don't have to start with just reactants ... you can have products in the initial mix as well ... or even only products), the two competing processes typically occur at different rates. The rates depend upon the concentrations of the reactants and the products. You might have studied this in a Kinetics chapter. Rates of the reactions depend upon the concentration of the
reacting particles. So the forward reaction rate depends upon the concentration of reactants and the reverse reaction rate depends upon the concentration of products. As time progresses, the
faster reaction begins to lower its rate because it consumes its
reacting particles much quicker than the competing reaction produces them ... and the concentrations of its reacting particles are decreasing. Meanwhile, the
slower reaction begins to increase its rate because the concentration of its reacting particles is increasing ... and rates depend upon these concentrations.
Eventually, the two competing processes take place at the same rate. When this occurs, we say that
equilibrium has been reached. Balance between the two competing processes - the forward reaction and the reverse reaction - has been established. And because the effects of the two competing processes have been balanced, the amounts of reactants and products will no longer change. The two processes are still taking place - reactants are still turning into products and products are still turning into reactants. It's just that the rate of the two processes is equal. And because reactants are being consumed (by the forward reaction) and the same rate that they are being produced (by the reverse reaction), the amount of reactants is holding steady. And the same can be said of the products. We sometimes refer to this as a
dynamic equilibrium.
To summarize the above, we could say that there are
two conditions that exist when the state of equilibrium has been reached:
- The rate of the forward reaction is equal to the rate of the reverse reaction.
- The amount (or concentrations) of reactants and products are no longer observed to be changing.
What About Those Graphs?
Improving your skill at reading, analyzing, and interpreting graphs (and charts and tables and illustrations) is a major goal of all science instruction. So here's a chance to improve that skill.
There are two graphs with two plotted lines on each graph. Each graph displays how
something changes over time. The graph on the left shows how the concentrations of the reactants and the products change over time. You will ALWAYS observe that these lines curve in opposite directions. Either the reactant concentration decreases while the product concentration increases or vice versa. Here's how you can tell which is which: if a line representing the reactant concentration is
going up (rising), then the reactant concentrations are increasing (and the product concentrations are
going down or decreasing). If this is the case, then you will probably note that on the second graph - the Rate vs. Time graph - the rate of the reverse reaction is higher on the graph than the rate of the forward reaction. The lines are labeled so check it out. Now think about this and try to make sense of it. If the reverse reaction that produces the reactants is occurring at a higher rate (you can say it's
faster) than the forward reaction that consumes the reactants, then you would expect the amount of reactants to be on the increase. This is because those reactant particles are being made (by the reverse reaction) more rapidly than they are being
destroyed (by the forward reaction).
Now let's consider the logic associated with the opposite case - the line representing the reactant concentration is
going down (decreasing). This indicates that the reactant concentrations are decreasing (and the product concentrations are increasing as demonstrated by its rising line on the concentration graph). If this is the case, then you should note that the rate vs. time graph displays the forward reaction as occurring at a greater rate (its higher on the rate-time graph) than the reverse reaction. When the forward reaction takes place more rapidly than the reverse reaction, then reactant particles are being consumed more rapidly than they are being formed. This causes the amounts (and the concentrations) of reactants to decrease and the amounts of products to increase.
What About the Flattening Out of the Curves?
Yes. I'm glad you asked. The lines on both graphs eventually become flat. Remember what was said earlier: when equilibrium is reached, ...
- the rates of the forward and the reverse reaction are equal to each other.
- the concentrations of the reactants and products are no longer changing.
The first bullet point - equal rates - is displayed by the graph on the right (the rate-time graph). Note that the two lines eventually flatten out and do so at the same rate value. You probably want to say it this way - "the lines come together." They do! But interpret the meaning of this ... the rates of the two competing processes become equal to each other.
The second bullet point - non-changing concentrations - is displayed by the graph on the left (the concentration-time graph). Note that both lines eventually become flat; each maintains its own constant level of concentration. That means the concentrations of reactants and products are no longer changing. But also note that these two lines (on the graph at the left) become flat at two different concentration levels ... communicating that while the concentrations of reactants and products are no longer changing, they aren't equal to each other. The final, equilibrium concentration of the reactants is different than the final, equilibrium concentration of the products. Don't be fooled - the rates of the two processes are equal to each other but the concentrations of the two
players are different than one another.