Midterm I

• Exothermic: reaction where heat is released to surroundings, products have lower energy than reactants
• Endothermic: reaction where heat is absorbed from surroundings, products have higher energy than reactants
• Activation Energy: The minimum amount of energy needed for a reaction to occur.
• Products/Reactants in endothermic reaction: Products are higher in energy, bonds broken are weaker than bonds formed
• Potential Energy: energy due to physical position or composition
• Kinetic Energy: energy of atomic motion, proportional to temperature
• When beginning to boil water, the bubbles are atmospheric gasses like oxygen nitrogen and carbon dioxide
• Once water is boiling, the bubbles are made of H2O held in liquid by hydrogen bonds
• When liquid H2O becomes a gas, only hydrogen bonds break
• First law of thermodynamics: energy is conserved
• Change in energy of the system is equal to the heat added (q) and work done on the system (w)
• If q is positive, systems gains heat from surroundings
• If work is positive, work is done on system by surroundings
• heat and work are both dependent functions, sum up to a state function (delta E)
• State function: changes are path independents, what happens in between does not matter
• Examples of paths: constant pressure or temperature, constant volume, adiabatic (q=0)
• Work: force times distance, negative for work done on system, w = -Fd
• Pistin shows how the expansion or compression of gas against a constant external pressure, w = force times change in height
• Enthalpy: The total energy of a system before and after a reaction has taken place, state function, H=PV
• At constant temperature or at constant pressure Delta H = q
• Internal Energy vs. Enthalpy INSERT HERE
• If the temperature of a system is constant, kinetic energy can be disregarded. Potential energy has to do with the possibility of bonds breaking or forming
• In constant volume, q=delta E. because work is zero. In constant pressure, Delta H = q because the volume change does work on the surroundings.
• Delta H = Delta E + PV
• The difference between delta E and delta H is the gas volume change. Volume will not equal zero when 1) chemical or physical change to number of moles of gas. 2) heat or cool gas causing temp change
• If the number of moles increases, the volume increases, making the work done by the system via expansion, resulting in a positive volume change. Pv is positive.
• If the number of moles decreases, the volume decreases, making the work done onto the system via compression, resulting in a negative volume change. PV is negative
• When heating ideal gas at constant pressure, some added heat raises the temperature of the gas, some does work on surroundings.
• At constant pressure, DeltaE=qCDeltaT. Molar heat capacity = 3/2 R.
• q at constant pressure for ideal gas = 5/2 RDeltaT
• Heat capacity of polyatomic gases: internal motions and translations like vibrational and rotational modes. Have larger Cv and Cp than monoatomic gases because greater internal kinetic energy. Weakly temperature dependent. Cp = Cv + R still true.
• Under ideal conditions, Cv of gas = 3/2 R, Cp of gas = 5/2 R
• Under constant pressure, some E is lost in vibrations/kinetics so change in temperature is less because gas is doing more effective work.
• Solid and liquid heat capacities more complicated because of interactions between molecules. Different phases --> different heat capacities.
• Calorimetry: experimental measure of heat flow in a system under constant pressure. Water is usually the surroundings. Reaction is the system.
• Adiabatic: a process in which no heat is transferred into or out of a system, and the change in internal energy is only done by work. q=0, internal energy maximized
• Hess' Law: DeltaH for overall process is sum of individual enthalpy changes for each step. Can be used for changes in other path independent state functions,
• Modifications for Hess' Law: if moles are different, multiply value. if the reaction needs to be flipped, change the sign.
• Standard States: when enthalpy change is measured, specify conditions for consistency. Tabulated values assume: gases at 1.00 atm pressure. Solids and liquids: pure substance. Solutes in solution: 1 M concentration. Temperature: 25 C or 298.15 K. Assuming that the volume and pressure are constant in measurements of enthalpy.
• Standard Enthalpy of Formation: the heat given off or absorbed when elements in their most stable state combine to form one mole of a compound. Measured at standard state.