Glossary
Common intermediate
A substance that is produced in one step of a reaction mechanism and then consumed in a subsequent step, linking two or more reactions together in a coupled process. It does not appear in the overall balanced chemical equation.
Example:
In the overall reaction A → C, if A → B is the first step and B → C is the second, then B is the common intermediate that connects the two steps.
Coupled reactions
A process where a thermodynamically unfavorable (nonspontaneous) reaction is made to occur by linking it with a thermodynamically favorable (spontaneous) reaction. The overall process becomes spontaneous if the favorable reaction releases enough energy to overcome the energy requirement of the unfavorable one.
Example:
In biological systems, the synthesis of complex molecules (an unfavorable process) is often achieved through coupled reactions with the hydrolysis of ATP, which is highly spontaneous.
Electricity
A form of energy involving the flow of electrons, often used as an external energy source to drive nonspontaneous chemical reactions.
Example:
When you recharge a phone battery, electricity is used to reverse the spontaneous discharge process, forcing the nonspontaneous reaction to occur.
Electrolytic cells
Electrochemical cells that use an external electrical energy source to drive nonspontaneous redox reactions. They convert electrical energy into chemical energy.
Example:
The industrial production of aluminum from aluminum oxide (Al2O3) uses an electrolytic cell to force the reduction of Al3+ ions, which wouldn't happen spontaneously.
Hess’s Law (for ΔG°)
A principle stating that if a reaction can be expressed as the sum of two or more other reactions, the change in Gibbs free energy (ΔG°) for the overall reaction is the sum of the ΔG° values of the individual reactions.
Example:
To find the ΔG° for the formation of CO2 from C and O2, you can use Hess's Law by adding the ΔG° values of C + 1/2 O2 → CO and CO + 1/2 O2 → CO2.
K (Equilibrium Constant)
A value that expresses the ratio of products to reactants at equilibrium for a reversible reaction, indicating the extent to which a reaction proceeds. For thermodynamically favorable reactions, K > 1, and for unfavorable reactions, K < 1.
Example:
If the K for a reaction is 10^5, it means that at equilibrium, there are significantly more products than reactants, indicating the reaction strongly favors product formation.
Nonspontaneous reactions
Reactions that do not proceed on their own and require a continuous input of energy to occur, typically characterized by a positive change in Gibbs free energy (ΔG > 0).
Example:
Charging a rechargeable battery involves a nonspontaneous reaction that requires electrical energy to reverse the discharge process.
Spontaneous reactions
Reactions that proceed on their own without continuous external energy input once initiated, typically characterized by a negative change in Gibbs free energy (ΔG < 0).
Example:
A piece of iron rusting in the presence of oxygen and water is a spontaneous reaction, though it might be slow.
Thermodynamically favorable reactions
Reactions that occur spontaneously under a given set of conditions without continuous external energy input. They are characterized by a negative change in Gibbs free energy (ΔG° < 0) and an equilibrium constant greater than 1 (K > 1).
Example:
The combustion of methane (CH4 + 2O2 → CO2 + 2H2O) is a thermodynamically favorable reaction that releases a lot of energy, which is why it's used in natural gas stoves.
Thermodynamically unfavorable reactions
Reactions that do not occur spontaneously under a given set of conditions and require a continuous external energy input to proceed. They are characterized by a positive change in Gibbs free energy (ΔG° > 0) and an equilibrium constant less than 1 (K < 1).
Example:
The decomposition of water into hydrogen and oxygen gas (2H2O → 2H2 + O2) is a thermodynamically unfavorable reaction that needs electricity to happen.
ΔG° (Gibbs Free Energy Change)
The change in Gibbs free energy for a reaction under standard conditions (1 atm pressure, 298 K, 1 M concentration). It is a thermodynamic potential that measures the maximum reversible work that a system can perform at constant temperature and pressure.
Example:
A reaction with a ΔG° of -50 kJ/mol indicates it is spontaneous under standard conditions and can do 50 kJ of useful work per mole.