University:
Winston-Salem Forsyth County Schools
Course:
AP Chemistry
Academic year:

2014
Views:

470

Pages:

4
Author:

horselingd3sd

THERMOCHEMISTRY • Kinetic energy is the energy of motion: KE= ½ mv2 • Potential energy is the energy an object possesses by virtue of its location. Potential energy can be converted to kinetic energy. • SI unit is the Joule: 1J= 1kgxm2/s2 • 1 calorie= 4.184 joules • A system is the part of the universe we are studying. The surroundings are everything else. • Force is a push or pull on an object. • Work is the amount of force applied to an object over a distance: w=Fxd • Heat is the energy transferred from a hotter object to a cooler object. • Energy is the capacity to do work or transfer heat. • The first law of thermodynamics states that energy can not be created or destroyed. The total energy of(system+surroundings) is constant. • The total energy of a system(KE+PE) is called the internal energy. • ∆E=Efinal – Einitial _ • When a system undergoes a physical or chemical change, the change in internal energy is equal to the heat added to or liberated from the system plus the work done on or by the system: ∆E= q+w • When heat is transferred from the surroundings to the system q>0 ( system is cold to the touch because it is absorbing heat from your hand) • Work done by the surroundings on the system is positive,w>0. • Endothermic reactions absorb heat from the surroundings. Exothermic reactions liberate heat to the surroundings. • A state function depends on only the initial and final states of the system. E is a state function. q and w are not. • A state function does not depend on how the internal energy is used. • Enthalpy is the heat transferred between a system and the surroundings during a chemical reaction carried out under constant pressure: H • ∆H= Hfinal-Hinitial = qp • Endothermic reactions have a positive enthalpy, exothermic reactions have a negative enthalpy. • Enthalpy of reaction: ∆H= Hproducts- Hreactants • 2H2 + O2→2H2O ∆H = -483.6 kJ What does this thermochemical equation tell you? • Enthalpy is an extensive property. The magnitude is dependent on the amount of reactant. Burning 1 mole of methane produces 890 kJ. Burning 2 moles produces 1780 kJ. • The sign of ∆H depends on the direction of the reaction. The reverse reaction is equal in magnitude but opposite in sign. • Calorimetry is the measure of heat flow. • Heat capacity is the amount of energy required to raise the temperature of an object 1 degree C. Molar heat capacity is the heat capacity of 1 mole of substance. Specific heat is the amount of energy required to raise the temperature of 1 gram of substance 1 degree C. • q=(specific heat) x (grams of substance) x ∆T • ∆H= qp • qsol= ( specific heat of solution)x (grams of solution) x ∆T= -qrxn • Constant volume calorimetry is carried out in a bomb calorimeter. Typically used for combustion reactions. • qrxn= - Ccal x ∆T • Hess’s Law: If a reaction is carried out in steps, ∆H for the reaction is the sum of the ∆H’s for each of the steps. • Example:CH4(g)+2O2(g)→CO2(g)+2H2O(g) ∆H= -802 kJ 2H2o(g)→ 2H2O(l) ∆H=-88 kJ CH4(g)+2O2(g)→CO2(g)+2H2O(l)∆H=-890kJ