System . surroundings . universe . thermodynamic state of a system .

1. Heat, Latent Heat, Sensible Heat Energy Internal energy Kinetic Energy Potential Energy Endothermic Exothermic Thermodynamics Thermal Equilibrium System Surroundings Law of Conservation of Energy Law of Conservation of Mass Heat Capacity, Molar Heat Capacity Specific Heat Capacity First Law of Thermodynamics State Function Standard state temperature Standard state pressure Standard states matter Hess’s Law Closed System Open System Isolated System Heat of Fusion Heat of Vaporization System. surroundings. universe. thermodynamic state of a system. • The number of moles and identity of each substance. • The physical states of each substance. • The temperature of the system. • The pressure of the system. Thermochemical standard state conditions • The thermochemical standard T = 298.15 K. • The thermochemical standard P = 1.0000 atm. • Be careful not to confuse these values with STP. Thermochemical standard states of matter • For pure substances in their liquid or solid phase the standard state is the pure liquid or solid. • For gases the standard state is the gas at 1.00 atm of pressure. For gaseous mixtures the partial pressure must be 1.00 atm. • For aqueous solutions the standard state is 1.00 M concentration. State Functions are independent of pathway: • T (temperature), P (pressure), V (volume), E (change in energy), H (change in enthalpy – the transfer of heat), and S (entropy) Examples of non-state functions are: • n (moles), q (heat), w (work) Entropy Enthalpy Gibbs Free Energy

2. Specific heat capacity = heat lost or gained by substance (J) (mass, g) (T change, K) Specific Heat Capacity How much energy is transferred due to Temperature difference? The heat (q) “lost” or “gained” is related to a) sample mass b) change in T and c) specific heat capacity

3. Specific heat capacity (J/(g∙K) = heat lost or gained by system (Joules) mass(grams) DT (Kelvins) ∆H = Hfinal - Hinitial • The stoichiometric coefficients in thermochemical equations must be interpreted as numbers of moles. 1 mol of C5H12 reacts with 8 mol of O2 to produce 5 mol of CO2, 6 mol of H2O, and releasing 3523 kJ is referred to as one mole of reactions. • ∆Horxn =  ∆Hfo(prod) -  ∆Hfo(react) q cP = m(Tf –Ti)

4. heat transfer in (endothermic), +q heat transfer out (exothermic), -q w transfer in (+w) w transfer out (-w) SYSTEM ∆E = q + w

5. There are two basic ideas of importance for thermodynamic systems Law of Conservation of mass and the Law of Conservation of Energy. The second leads to 3 thermodynamic laws. • First Law – the energy of the Universe is constant Chemical systems tend toward a state of minimum potential energy. The first law is also known as the Law of Conservation of Energy. Energy is neither created nor destroyed in chemical reactions and physical changes. Esys +Esurr = EuniverseEsys = KEsys + PEsys • KE – kinetic energy: translational, rotational, vibrational • PE – energy stored in bonds • The second law of thermodynamics states, “In spontaneous changes the universe tends towards a state of greater disorder.” Chemical systems tend toward a state of maximum disorder. The entropy of universe must increase. • Fundamentally, the system must be capable of doing useful work on surroundings for a spontaneous process to occur. In general for a substance in its three states of matter: • Sgas> Sliquid > Ssolid • When: • S > 0 disorder increases (which favors spontaneity). • S < 0 disorder decreases (does not favor spontaneity). • The Third Law of Thermodynamics states, “The entropy of a pure, perfect, crystalline solid at 0 K is zero.”

6. heat transfer in (endothermic), +q heat transfer out (exothermic), -q w transfer in (+w) Compression of system w transfer out (-w) Expansion of system DG = DH-TDS at constant T and P • The relationship describes the spontaneity of a system. • The relationship is a new state function, G, the Gibbs Free Energy. • Sign conventions for G. • G > 0 reaction is nonspontaneous • G = 0 system is at equilibrium • G < 0 reaction is spontaneous SYSTEM ∆E = q + w

7. Free energy has the relationship • G = H -TS. • Because 0 ≤ H ≥ 0 and 0 ≤ S ≥ 0, • there are four possibilities for G. • Forward reaction • HSG spontaneity • < 0 > 0 < 0 at all T’s. • < 0 < 0 T dependent at low T’s. > 0 > 0 T dependent at high T’s. > 0 < 0 > 0 Nonspontaneous at all T’s.

8. Kinetics Four factors that affect the rate of reaction nature of reactant concentration temperature presence of a catalyst 0 order [A] vs t 1/[A] vs t 2nd order ln [A] vs t 1st order Rate of reaction Rate constant Order of reactant Overall order of reaction 1/[A] ln [A] t t General rate expression Initial rate Instantaneous rate Average rate Integrated rate laws ([ ] with respect to time) If zero order [A]0 - [A] = ak t first order second order • Differential rate laws ([ ]with respect to rate) • Rate = k [A]m[B]n[C]p • If zero order Rate = k[A]0 = k • first order Rate = k[A]1 = k[A] • second order Rate = k[A]2 • Change in rate is independent of the change in concentration of a zero order reactant • Change is rate is directly proportional to the change in concentration of a 1st order reactant • Change in rate is directly proportional to the square of the change in concentration of a 2nd order reactant [A] t

9. All equations are shown for a stoichiometric coefficient of 1. for all others use the term akt, or ak in place of kt and k below. Five factors that affect the rate of reaction: nature of reactant Concentration Temperature presence of a catalyst solvent effects Rate of reaction Rate constant Order of reactant Overall order of reaction Substitute ½ [A]o for [A] In integrated rate law equation to get ½ life equation