Driving Forces

Enthalpy, H, is a term that equates to the heat involved in any chemical changes. A positive enthalpy means an endothermic reaction, or the products end with more heat energy than the reactants. Negative enthalpy is an exothermic reaction, which loses heat as it progresses.

Enthalpy of formation is the energy required to make 1 mole of something from its elements. In standard conditions, it is usually 0. CO₂ however, does not have an enthalpy of formation equal to 0. Its enthalpy of formation is -393.5 kJ/mol, meaning the reaction to create CO₂ is exothermic.

When it comes to different phases, solids require the most energy, and gases require the least. Liquid water has an enthalpy of formation, ΔHᵒ, of -285.8 kJ/mol, whereas the enthalpy of formation of water vapor is -241.8 kJ/mol. This is because the formation of bonds is exothermic, and liquid water bonds H₂O atoms together with hydrogen bonds. Gases are not bound together with IMFs, so they are endothermic and therefore require less energy.

When there is a temperature change, enthalpy also changes. This can be calculated using the equation:

Q = mcΔt

where m is mass, Δt is change in temperature, Q is heat, and c is specific heat, which is the heat involved in changing the temperature of 1 gram of a substance by 1ᵒ C. Specific heat is a constant that is different for every substance. To get enthalpy from Q, you simply divide it by 1000, because Q is in J/mol rather than kJ/mol.

However, if the temperature passes a boiling/melting point and a phase change is involved, the enthalpy changes. For example, raising the temperature of water from 90ᵒ C to 110ᵒ C changes it from liquid to gas. The heat involved in a phase change can be calculated using this equation:

Q = mL

where m is mass, Q is heat, and L is latent heat, which is the heat required to change the state of 1 gram of a substance between gases and liquids. This once again is a constant that changes for every substance.

The first law of thermodynamics can be quantified into

ΔE = Q + w

where ΔE is change in internal energy, q is heat, and w is work. This is a way of getting energy using enthalpy. Work can be subbed in for PΔV, or pressure times change in velocity, as long as pressure is constant.

Entropy, the other driving force is a measure of randomness or disorder. From the 2nd law of thermodynamics, we know that entropy will always increase on a universal scale. Randomness or disorder can be quantified by a number of microstates possible. A microstate is a possible position or energy of something.

For example, dropping a stack of 5 sheets of paper from 1 inch does not have many possible outcomes. There is only a small area where the papers are going to move, if any. There are not many possible microstates, or, in other words, there is not much entropy. However, if you dropped those same 5 sheets of paper from 10 feet up, the papers are going to fly every which way. There are many different possible outcomes, so many possible positions and energies, and therefore, many microstates.

Gases have the most entropy, as there are the same amount of particles but over the largest space. Heavy solids have more entropy than light solids as there are more particles. The same goes for compounds vs elements. Warmer substances have more entropy, because as a substance heats up its particles speed up, creating more kinetic energy and more microstates.

The 2nd law of thermodynamics can be quantified as:

ΔS = ΔQ/T

where ΔS is the change in enthalpy in J/K, ΔQ is the change in heat in J, and T is the temperature in K. This leads to the third law, which states that any solid has zero entropy at 0 K.

All of these rules come to the conclusion that reactions occur so they can proceed to lesser entropy and greater enthalpy. Ice melting is an example of enthalpy lowering. When jello gels after a period of time, that is an example of entropy increasing. Candles burning leads to lower energy and greater entropy.

Spontaneous reactions are defined as reactions that occur naturally without being forced, or without needing any energy inputs. Spontaneity is found using a combination of enthalpy and energy conditions:

ΔG = ΔH - TΔS

where ΔH is enthalpy, T is temperature in K, ΔS is entropy, and ΔG is Gibbs free energy. If ΔG is positive, the reaction is not spontaneous and will not happen without outside energy. If ΔG is negative, the reaction is spontaneous, and it will happen by itself. Exothermic reactions and reactions at higher temperatures tend to be spontaneous.

The actual definition of Gibbs free energy is available energy released by other reactions into the environment. That energy can be used to do work. Unused free energy is turned into entropy. When ΔG is 0, the forward and reverse reactions are both equally likely, and both spontaneous. A reverse reaction is the same reaction but where the products and reactants switch.