Entropy describes the number of configurations a system can have or, alternatively, how much "order" it possesses.

When water freezes, it forms a crystal lattice, which is more ordered than liquid water. Likewise, collecting playing cards and stacking them in a deck involves a transformation from disorder (cards scattered around the floor) to order (a single deck). Chemically, the precipitation of  from aqueous ions comes from a decrease in entropy, as the free-flowing ions have been combined to form a solid, organized in a lattice. The combination of  and  to give  also results in a decrease in entropy, since the reactants have one mole of gas, and the products contain zero moles of gas.

The disproportionation of  to give two moles of  results in an increase in entropy, because the number of moles of gas increases in going from reactants to products.

The entropy change of the surroundings for an exothermic system must be positive. The release of heat by the exothermic reaction drives the increase in disorder of the environment surrounding the reaction vessel.

As the heat leaves the reaction system, it increases the entropy of the particles in the surroundings.

Entropy decreases when the number of moles of gas decreases during a reaction. In the case of the correct answer, the number of moles of gas decreases from two to one.

When a substance goes from a solid to a gas (sublimation) or from a liquid to a gas (evaporation), entropy increases. Likewise, when a solid dissolves in water, entropy increases. Entropy (i.e. the number of arrangements a system can have) is much greater in a gas than in a liquid or solid. It is also greater for ions solvated in solution than for an un-dissolved solid.

Note that, in general, the entropy of the universe increases. In order for a process to involve a decrease in entropy of the system, there is likely a consequent increase in entropy of the universe.

Entropy is a measure of disorder or randomness. A system with more random motion between molecules has greater entropy. The phases of matter in order of increasing entropy are solid, liquid, then gas. The processes that increase entropy by changing phases will cause a phase transition from lower entropy to higher entropy. These transitions are melting (solid to liquid), vaporization (liquid to gas), and sublimation (solid to gas).

An increase in entropy will result when a solid is dissolved into a liquid because disorder is increased due to the presence of a greater number of particulates (ions). Similarly, any reaction that results in more molecules represents an increase in entropy, as in the example of the decomposition of hydrogen peroxide. A temperature increase does not relate to entropy; it only helps describe the energy of the reaction. An increase in temperature usually signifies an endothermic process, but does not give an indication about entropy.

The only correct answer is the chemical reaction between aluminum and iron (III) oxide, which produces an equal amount of particles and does not involve any phase change.

Even though the dissolution process is exothermic, breaking bonds ALWAYS requires energy input. In the case of an exothermic dissolution, more energy is released when solute-solvent bonds are formed, than was absorbed when solute-solute bonds were broken. The initial step requires an input of activation energy, but the final product allows for a net energy release.

When dealing with the enthalpy of a solution, there are three specific steps, each having their own enthalpy.

1) The breaking of solute-solute

2) The breaking of solvent-solvent bonds

3) The formation of solute-solvent bonds

Remember that breaking bonds always requires an input of energy, and is thus considered endothermic. The formation of new bonds is an exothermic process. When an overall solution is exothermic, it means that the new intermolecular bonds are more stable. A system with less energy is considered more stable.

When a reaction has a positive enthalpy change, it is said to be endothermic. This means that it requires an addition of heat into the system. Because the forward reaction requires more energy than the reverse reaction in order to take place, we see that the activation energy of the endothermic reaction is greater than the activation energy of the reverse reaction.

Essentially, the forward reaction requires a net input of energy, while the reverse reaction results in a net release of energy.

The process of bond formation in chemistry is exothermic. By definition, exothermic reactions have a negative enthalpy change; therefore  must be less than zero in this bond-forming reaction.

Reducing moles of gas will decrease entropy, meaning that . If both entropy and enthalpy are negative, as in this reaction, we cannot determine the value of Gibbs free energy without knowing the temperature.

An exothermic process, by definition, involves a reaction in which the products are lower in energy than the reactants. The reduction in chemical energy results in a release of heat from the reaction.

In the diagram, the path between points is irrelevant. We are simply looking for any instances in which the product point is below the reactant point. Point C has less energy than point B, and point D has less energy than A, B, or C. Transitions from B to C, C to D, or A to D will all result in a reduction of chemical energy, and a release of heat.