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

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  • 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

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