Energy Changes in Chemical Reactions

• All chemical reactions exhibit the two fundamental laws: the law of conservation of mass and the law of conservation of energy.
• Energy

  The capacity to do work
• Forms of energy


  • Kinetic energy
  • Thermal energy
  • Chemical energy
  • Potential energy
• Energy can be transformed from one form to another.
• Energy can neither be created nor destroyed.
• The total quantity of energy in the universe is assumed constant.
• Heat


  The transfer of thermal energy between two bodies that are at different temperatures
• Thermochemistry

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

  • Open
  • Closed
  • Isolated
• Open system

  Can exchange mass and energy with its surroundings
• Closed system


  Allows the transfer of energy (heat) but not mass
• Isolated system


  Does not allow the transfer of either mass or energy
• The combustion of hydrogen gas in oxygen


  Releases considerable amount of energy
• Combustion of hydrogen

  2H2(g) + O2(g) à 2H2O(l) + energy
• Exothermic process

  A process that gives off heat
• Decomposition of mercury (II) oxide

  energy + 2HgO(s) à 2Hg(l) + O2(g)
• Endothermic process


  Heat has to be supplied to the system by the surroundings
• Calorimeter


  A closed container designed specifically to measure heat changes
• Calorimetry


  The measurement of heat changes
• Specific heat (s)


  The amount of heat required to raise the temperature of one gram of the substance by one degree Celsius
• Heat capacity (C)


  The amount of heat required to raise the temperature of a given quantity of the substance by one degree Celsius
• Specific heat has the units J/g·°C.
• Heat capacity has the units J/°C.
• Path function

  Heat (q) is a path function, which values are dependent on the path taken
• Calculating heat change

  1. q = msΔt

  2. q = CΔt
• Δt


  The temperature change
• The sign convention for q is positive for endothermic process and negative for exothermic process.
• Heat of combustion is usually measured by placing a known mass of a compound in a steel container called a constant-volume bomb calorimeter.
• The closed bomb is immersed in a known amount of water and the sample is ignited electrically.
• The heat produced by the combustion reaction can be calculated accurately by recording the rise in temperature of the water.
• The heat given off by the sample is absorbed by the water and the bomb.
• qsystem

  The heat change of the system
• qsystem = 0


  qcal + qrxn
• The quantity of Ccal is calibrated by burning a substance with an accurately known heat of combustion.
• Constant-pressure calorimeter is used to determine the heat changes for noncombustion reactions.
• Heat changes for the process (qrxn) is equal to the enthalpy change (ΔH).
• Enthalpy (H)


  Defined by the equation H = E + PV
• ΔH is positive for endothermic processes and negative for exothermic processes.
• ΔH value does not refer to a particular reactant or product
• Quoted ΔH value refers to all the reacting species in molar quantities