Energy

Energy is classified: Kinetic energy – energy of motion Potential energy – energy of position Both energies can be transferred from one object to another. Examples: temperature, energy

Thermodynamics Energy changes that occur during a chemical reaction are due to the making and breaking of chemical bonds. Breaking chemical bonds always requires an input of energy.

System is that small portion of the universe in which we are interested. Surroundings are everything else. Boundary is the separation between the two.

Kinetic Theory of Heat Heat, when it enters a system, produces an increase in the average motion (the average kinetic energy) with which the particles of the system move. All atoms, molecules and ions are in continuous random motion.

Specific Heat The heat needed to raise the temperature of one gram of a substance by one degree Celsius.

Molar Heat Capacity Heat required to raise the temperature of one mole of a substance by one degree Celsius.

State Functions: Extensive properties: mass, volume Intensive properties: don’t depend on size State function if it depends only on the state of the system, not on the path used to get to that state. Examples: temperature, energy, enthalpy

First Law of Thermodynamics Energy is conserved Energy cannot be either created or destroyed. E total = E sys + E surr

Enthalpies of Reaction Exothermic: give off heat to surroundings Endothermic: absorb heat from surroundings Exothermic Reactions: ∆H is negative Endothermic Reactions: ∆H is positive

Chemical Energy Exothermic: CH 4 (g) + 2 O 2 (g) → CO 2 (g) + H 2 O (g) + energy Endotheric: N 2 (g) + O 2 (g) + energy → 2 NO (g)

Internal Energy Internal energy of a system is the sum of the kinetic and potential energies of ALL the particle in the system. ∆E (internal energy) = q (heat) + w (work) If the sign in positive, the heat is flowing into the system; if it is negative, the heat if flowing out of the system.

Example Calculate ∆E for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1.4 kJ of work is done on the system. ∆E = q + w ∆E = 15.6 kJ kJ = 17.0 kJ The system has gained 17.0 kJ of energy

The common type of work with chemical processes is work done by a gas (expansion) or work done to a gas (compression). Pressure is utilized in either case. Pressure is force per unit area, the pressure of a gas is P = F/A

Work is defined as force applied over distance: Work = F x ∆h Since P = F/A, then Work = F x ∆h = P x A x ∆h. (Area x a change in height = volume) Work in a cylinder then is: w = - P∆V

Example Calculate the work with the expansion of a gas from 46 L to 64 L at a constant external pressure of 15 atm. w = - P∆V w = - 15 atm x 18 L = L. atm Since the gas expands, it does work on its surroundings. Energy flows out, so w is negative.

Enthalpy Enthalpy (H) is defined as H = E + PV Internal energy, pressure, and volume are all state functions as well as enthalpy. ∆E = q p + w, which becomes ∆E = q p - P∆V Or q p = ∆E + P∆V So a change in H = (change in E) + (change in PV) At constant pressure, the change in enthalpy of the system is equal to the energy flow as heat.

∆H = H products - H reactants When a mole of methane (CH 4 ) is burned at a constant pressure, 890 kJ of energy is released as heat. Calculate ∆H for a process in which a 5.8 gram sample of methane is burned at constant pressure. Q p = ∆H = kJ/mol (CH4)c, then 5.8 g of methane equals 0.36 mol. ∆H = 0.36 mol (CH4) x – 890 kJ/mol (CH4) = kJ.