Reading energy and progress from the diagram
A reaction coordinate diagram plots energy (vertical axis) versus reaction progress (horizontal axis). The leftmost point is the energy of reactants, the rightmost is the energy of products. The curve rises from reactants to a peak (the transition state), then descends to products. The vertical gap between reactants and transition state is the activation energy (Ea), the minimum energy required for the reaction to proceed. The vertical gap between reactants and products is the reaction enthalpy (ΔH or ΔG), which determines whether the reaction is thermodynamically favorable.
Activation energy and catalysts
Reactions with high activation energy are slow at room temperature because few molecules have enough kinetic energy to reach the transition state. A catalyst lowers the activation energy by providing an alternative pathway with a lower transition state energy. Enzymes are biological catalysts that can lower Ea by 10-20 kJ/mol or more, accelerating reactions to rates that sustain life. The catalyst does not change the overall energy difference between reactants and products; it only affects the height of the barrier, not the starting and ending points.
Thermodynamic vs kinetic control
A reaction is thermodynamically favorable if products are lower in energy than reactants (ΔG < 0). But if the activation energy is very high, the reaction proceeds so slowly that it appears not to happen. Conversely, a reaction can be unfavorable thermodynamically (products higher in energy) but still proceed if the activation energy is low. This is why diamond does not spontaneously turn into graphite at room temperature, even though graphite is more stable; the activation energy is prohibitively high. Balancing thermodynamic favorability with kinetic accessibility is central to chemical design and synthesis.