I.    Thermodynamics: the branch of science that deals with heat and its relationship to other forms of energy: mechanical, electrical, chemical, etc. 

II.    Different Types of Energy:

   A.    Electromagnetic Energy: the energy associated with light and electromagnetic waves.

   B.    Chemical Energy: the energy associated with chemical reactions between atoms or molecules. 

   C.    Nuclear Energy: the energy associated with changes in the nucleus (nuclear reactions) such as decay, fission, and fusion. 

   D.    Thermal Energy/Heat Energy: the energy associated with a system’s temperature. 

   E.    Mechanical Energy: the energy associated with the movement or the location of an object. A system’s mechanical energy is the sum of its kinetic and                        potential energy. 

III.    Laws of Thermodynamics:

   A.    First Law: energy can be converted from one form to another, but it can neither be created nor destroyed.  

   B.    Second Law: the entropy of the universe increases in a spontaneous process (S universe > 0), remains unchanged in an equilibrium process, and decreases               during a nonspontaneous process (S universe < 0). As you can see, this law connects entropy to the spontaneity of a reaction. 

   C.    Third Law: the entropy of a perfect crystal at absolute zero (0 K) is zero. At this temperature, molecular motions ceases, and the number of microstates is 1;            there is only one way to arrange the atoms to form a perfect crystal. This data serves as a reference point, allowing us to determine the absolute entropies of            substances other than perfect crystal.

IV.    Thermodynamic Functions/Quantities:

   A.    Enthalpy (H):

   B.    Gibbs Free Energy (G):

   C.    Entropy (S): 

V.    Gibbs Free Energy/Free Energy:

   A.    The energy available to do work. In other words, it can also be defined as the free energy associated with a reaction. 

   B.    If a reaction’s free energy is negative then it released energy out into the universe, and this energy can be used by any process that requires it. This is why               it’s called free energy. On the other hand, if a reaction’s free energy is positive, then its value represents the total amount of energy required for that reaction           to take place.

   C.    A system/reaction’s free energy can be defined mathematically in the following equation (the Gibbs-Helmholtz Equation) at constant temperature and                     pressure: G = H-TS. 

VI.    Enthalpy:

   A.    Heat flow. 

   B.    If a reaction’s change in enthalpy (H) is negative, then the reaction undergoes an exothermic process to produce its products. If it is positive, then heat is                 required for the reaction to take place, so the reaction is endothermic. 

VII.    Entropy: 

   A.    A measure of randomness/disorder. If a system or its surroundings have high disorder, it has a high entropy. 

   B.    Recall that solids are composed of highly ordered species (atoms, molecules) in fixed positions; they generally have a small number of microstates. Also,                the species within liquids have some order and some disorder. Lastly, gases consist of species that move in constant random motion. As a result, gases                      generally have a higher entropy than liquids, and liquids generally have a higher entropy than solids: S gas > S liquid > S solid. Melting allows the particles            within solids to occupy more positions, increasing the number of microstates. This results in an increase in entropy. Similarly, vaporization leads to an                      increase in the entropy of a system. 

   C.    Heavier/Larger/complex molecules have a higher entropy than lighter/smaller/simpler ones because they contain more particles and as a result have more                microstates and more disorder. 

   D.    When comparing species and predicting which has a higher entropy, study the macro picture first and then the micro picture. The macro picture includes                 moles (the coefficients in chemical equations), temperature, phase state. The micro picture includes the number of atoms (subscripts). Example: 3O2 (g)                 most likely has higher entropy than 2O3 (g) because the first species has a higher number of moles. Second example: CH4 (g) most likely has a smaller                   entropy than C6H6 (g) because the second specie contains more particles. 

   E.    Microstates/Microscopic States: all the possible microscopic configurations/arrangements that particles within a system can have. The more microstates, the          more possible ways of distributing particles and energy, the more disorder. It is difficult to determine a system’s entropy by determining how many                          microstates it has. Fortunately, alternative methods exist that are much easier. 

   F.    Standard Entropy: the entropy of a substance at standard conditions (1 atm and 25C). 25C is used when defining a standard state because this is equivalent to          77F (room temperature), and many processes are carried out at room temperature. Standard entropy is usually measured in J/K or J/(K mol) for 1 mole of a              substance. 

   G.    Absolute Entropy: the entropy of a substance at absolute zero. Recall that for a perfect crystal this is 0 K.