Thermochemistry

• Energy


  The capacity to do work
• Work


  Energy change resulting from a process
• Kinetic Energy


  • The energy produced by moving an object, it is one form of energy that is of particular interest to Chemists
• Other forms of energy


  • Radiant energy
  • Thermal Energy
  • Chemical Energy
  • Potential Energy
• When one form of energy disappears some other forms of energy of equal magnitude must appear, and vice versa
• Spontaneous processes


  Processes which take place on their own without the external intervention of any kind
• All the processes that takes place in nature are spontaneous in character, proceeds only in one direction and are, therefore thermodynamically irreversible
• In an exothermic reaction

  The bonds in the products are (overall) stronger than in the reactants
• Thermochemistry


  The study of heat change in chemical reactions
• Types of systems


  • Open system
  • Closed System
  • Isolated system
• Exothermic Process


  Any process that gives out heat to its surrounding
• Endothermic Process


  In this process heat is supplied to the system by the surrounding
• Diathermic boundaries


  Boundaries that permit the transfer of energy as heat
• Adiabatic boundaries


  Boundaries that do not permit the transfer of energy as heat
• Internal Energy (U)


  The total energy of a system, It consists of Potential and kinetic energies
• Molar internal energy (Um)


  The internal energy per mole of a substance
• For a monatomic gas, the total energy is 3/2 RT
• For a linear molecule, the total energy is 5/2 RT
• For a non-linear molecule, the total energy is 3RT
• 1st law of thermodynamics


  Energy can neither be created nor destroyed, it can only be transformed from one form to another
• Expansion work (w)


  Work arising from a change in volume
• Enthalpy (H)


  The sum of a system's internal energy and the work done by or on the system due to changes in pressure and volume
• Enthalpy of reaction


  The heat change that takes place at constant pressure
• For a constant pressure process, q = ∆H
• Constant-volume conditions are not common, most reactions occur under conditions of constant pressure, (usually atmospheric pressure)
• Enthalpy of Reaction


  A heat change that takes place at constant pressure
• Endothermic reaction


  ∆𝐻 > 0
• Exothermic reaction


  ∆𝐻 < 0
• The "per mole" in the unit for ∆𝐻 means that this is the enthalpy change per mole of the reaction
• Standard heat of formation


  The enthalpy change when one mole of the substance is made from its element in their standard state
• Standard state


  1 atm, and 𝟐𝟗𝟖 𝑲
• Standard heat of combustion


  The enthalpy change when one mole of the substance is completely burned in oxygen
• Heat of hydrogenation


  The enthalpy of hydrogenation is the heat change when one mole of an unsaturated compound react with hydrogen and is completely changed into the corresponding saturated compound
• In both cases the product is the same
• Standard enthalpy of formation (∆𝐻𝑟𝑥𝑛

  𝑜

  )

  The importance is that once the values are known, standard enthalpy of formation can easily be calculated
• By convention, the standard enthalpy of formation of any element in its most stable state is zero
• Many compounds cannot be directly synthesized from there elements. In some cases, the reaction proceeds too slowly, or side reactions produce substances other than the desired compound
• ∆𝐻𝑓

  𝑜 𝑐𝑎𝑛 𝑏𝑒 𝑑𝑒𝑡𝑒𝑟𝑚𝑖𝑛𝑒𝑑 𝑏𝑦 𝑎𝑛 𝑖𝑛𝑑𝑖𝑟𝑒𝑐𝑡 𝑎𝑝𝑝𝑟𝑜𝑎𝑐ℎ which is based on Hess's law, it can be stated as follows. When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in in one step or in a series of steps
• Thus the ∆𝐻𝑓

  𝑜 (𝐶𝑂) = −110.5 𝐾𝐽𝑚𝑜𝑙−1
• Calorimetry


  In the laboratory, heat changes in physical and chemical processes are measured with a calorimeter