dat gen chem 8-

• First Law of Thermodynamics, also known as the Law of Conservation of Energy, states that energy cannot be created or destroyed.
• think of energy transfer in terms of the system, the surroundings, and the universe.

  System = the chemical reaction containing reactants and products

  Surroundings = the environment that surrounds the system

  Universe = the sum of the system and its surroundings

  
• Enthalpy (H) refers to the amount of heat energy contained within a system.

  Endothermic processes: +H

  exothermic processes: -H

  
• Endothermic phase changes (heat is absorbed)

  1. Sublimation = solid to gas
  2. Vaporization = liquid to gas
  3. Melting = solid to liquid

  Exothermic phase changes (heat is released)

  1. Deposition = gas to solid
  2. Condensation = gas to liquid
  3. Freezing = liquid to solid
• Endothermic reactions involve molecules moving from being closer together to being further apart... requires an input of energy, because molecules need to gain sufficient kinetic energy to overcome intermolecular attractive forces.

  
• Exothermic reactions involve molecules moving from being far apart to being closer together.. this process requires a release of energy.

  
• 3 Types of Heat Transfers

  • Conduction - transfer of heat via direct contact.
  • Convection - convection currents exist in liquids or gases when hotter, less dense areas rise, and the cooler, more dense areas sink.
  • Radiation - transfer of heat via electromagnetic radiation.
• ∆E = q + w

  ΔE = change in internal energy

  q = heat

  w = work

  w = -P∆V

  w = work

  P = pressure (always positive)

  ΔV = change in volume (Vfinal – Vinitial)
• When there is a decrease in volume (ΔV < 0), work is done onto the system. This results in positive work, since work is done on the gas.

  
• When there is an increase in volume (ΔV > 0), work is done by the system. This results in negative work, since work is done by the gas.

  
• Heat and Enthalpy Relationship

  • Under constant pressure, heat (q) and enthalpy (H) within a system are the same (e.g., calorimetry).

  ∆E = q + w

  +q = endothermic

  -q: exothermic

  ∆E = q + w

  +w = compression

  -w = expansion

  
• Specific Heat Capacity (C): amount of energy required to raise the temperature of 1.0 gram of a substance by 1°Celsius.

  unit: J•g-1•°C-1.

  higher specific heat capacity: takes more heat to warm it up

  
• Specific Heat Capacity of Phases: specific heat capacity for the same molecule differs depending on the molecular state ex water has diff specific heat capacity it its ice or water vapor
• Calculating Heat Absorbed/Released: q = mCΔT

  q = heat absorbed or released by the substance (J)

  m = mass of the substance (g)

  C = specific heat capacity of the substance (J•g-1•°C-1)

  ΔT = change in temperature of the substance (°C)

  
  

  this formula can only be used for calculating the heat for a substance that is not undergoing a phase change.
• heat is transferred between the system and its surroundings. Thus, heat

  absorption and release can be described by the following relationships.

  qsystem + qsurrounding = 0

  qsystem = -qsurrounding
• Calculating Heat for Phase Changes

  q = m∆Hfus/vap

  q = heat absorbed or released by the substance (J)

  m = mass of the substance (kg)

  ∆Hfus = heat of fusion (J•kg-1) = heat required to change a substance from solid to liquid

  ∆Hvap = heat of vaporization (J•kg-1) = heat required to change a substance from liquid to gas.
• Phase Change Diagram

  As heat is added to the substance (x-axis), the temperature of the substance increases (y axis)

  notice during phase transitions, temperature remains constant... because during phase changes, the addition of heat only contributes to altering the substance’s phase, not to changing the temperature.

  
• Bomb Calorimetry: common way to measure the amount of heat produced by a chemical reaction

  • useful in food testing laboratories to determine the amount of heat (calories) in food.
• Standard Enthalpy Change (∆H°): enthalpy change that occurs in a chemical reaction under standard conditions

  STP: 298 K, 1 atm, 1.0 M concentration

  3 ways we can calculate for the standard enthalpy change:

  • Bond Enthalpies: using values for the heat energy absorbed during the breaking of bonds and heat energy released during the formation of bonds
  • Enthalpies of Formation: using values for the enthalpy change when 1 mole of a substance is formed from its constituent elements in standard states, we can calculate the enthalpy change of a more complex reaction.
  • Hess’s Law states that the enthalpy change of a reaction is independent of the pathway taken. Thus, we can add/subtract the enthalpy change of multiple reactions to obtain the enthalpy change of the reaction of interest.
• Breaking Bonds vs. Forming Bonds

  • Breaking bonds requires energy, its endothermic (∆H > 0)
  • Forming bonds releases energy, and is an exothermic process (∆H < 0). forming bonds is a more stable state
  • Different bonds have different bond energies associated with them.

  ∆H°rxn = Σ∆H°bonds broken – Σ∆H°bonds formed

  ex Find the ∆H°rxn for one mole of the following reaction:

  CH4 + Cl2 CH3Cl + HCl

  Bond Energies:

  C–H = 414 kJ•mol-1

  Cl–Cl = 243 kJ•mol-1

  C–Cl = 339 kJ•mol-1

  H–Cl = 431 kJ•mol-1

  = (414 + 243) – (339 + 431) =-113-113kJ