Direction of Reversible Reactions

A double arrow in an equation means the reaction is not a one-way trip. Reactants form products, but products also turn back into reactants, and both of these happen simultaneously from the first instant, not one after the other. Early on, if you start with pure reactants, the forward reaction is fast because there is plenty to react, and the reverse reaction is slow because almost no product exists yet. As products pile up, the reverse rate climbs while the forward rate falls, since reactants are being used up. Equilibrium is the moment these two rates become equal. Reactions have not stopped; they are still running both ways, just at matching speeds, so the amounts of everything stop changing.

This is why starting from pure products lands you in the same place as starting from pure reactants. The equilibrium position depends only on the reaction and the temperature, through the equilibrium constant K, not on which side you began. Start with only products, and the reverse reaction dominates at first, building up reactants until the rates match. Start with only reactants, or with some mixture of both, and the system still shifts whichever direction it needs to until forward and reverse rates equalize. Plotted as concentration versus time, both setups produce curves that start at different points but bend toward and meet at the same equilibrium concentrations. This sets up the next idea: once you know K and a starting mixture, an ICE table lets you calculate exactly where that resting point will be.