Thermodynamics

• Laws of thermodynamics


  Help us understand why energy flows in certain directions and in certain ways
• The laws of thermodynamics seem like common sense but there is a layer of math beneath the intuitive level that makes them very powerful at describing systems and making predictions
• First law of thermodynamics


  Energy is not created or destroyed, it only changes forms, from potential energy to kinetic energy to heat energy, etc.
• While the first law is untrue on the quantum level, for chemists it does just fine
• Second law of thermodynamics

  Introduces the concept of entropy, which can be described as disorder, and states that the sum of the entropies of a system and its surroundings must always increase
• Within a system, there is a tendency to go towards higher entropy
• Entropy analogy


  • Your bedroom will over time become messy but it won't suddenly become neat
  • Entropy is a measure of how dispersed the energy of the system is amongst the ways that system can contain energy
  • Entropy is like computer code - the solid state requires more information to describe than the liquid state
• Heat will flow from a hot coffee cup into the table or your hand because the heat energy will be more disordered if more dispersed
• Third law of thermodynamics


  A perfectly crystalline solid at absolute zero has an entropy of zero as this is the most ordered state the substance can be in
• Entropy


  Not a measure of energy itself but of how energy is distributed within a system
• Enthalpy


  The thermodynamic quantity that more accurately describes the energy of a system
• Gibbs free energy (G)


  Tells us whether a process will be spontaneous or not
• Change in Gibbs free energy (ΔG)


  Is given by the equation that includes change in enthalpy (ΔH), change in entropy (ΔS), and temperature (T)
• If ΔG is negative, the process is spontaneous. If ΔG is positive, the process is nonspontaneous
• A spontaneous process can be either enthalpically or entropically favorable, or both, but not neither
• Entropically unfavorable processes can be spontaneous at lower temperatures if they are energetically favorable
• Soap

  • Soap molecules have polar heads and long nonpolar tails, which allows them to spontaneously form structures called micelles that trap nonpolar dirt and grime
• Systems can defy entropy on the small scale, but the second law holds true in that the entropy of the universe is always increasing
• Internal energy


  Change in internal energy of a system = q + w
• q


  Amount of heat energy that enters or leaves the system
• If temperature of system is higher than surroundings


  q is negative (exothermic)
• If temperature of surroundings is higher than system


  q is positive (endothermic)
• w


  Work, equal to -p∆v
• When gas expands


  w is negative
• When gas is compressed

  w is positive
• 1 kJ = 1000 J, 1 cal = 4.184 J, 1 Cal = 1000 cal
• 101.3 J = 1 L·atm
• Calculating change in internal energy


  1. Use ∆E = q + w

  2. Plug in values for q and w
• q = mc∆T


  m = mass, c = specific heat capacity, ∆T = change in temperature
• q = m∆H


  ∆H = enthalpy of fusion/vaporization
• Balancing combustion reaction


  1. Balance carbon atoms first

  2. Balance hydrogen atoms

  3. Balance oxygen atoms
• Combustion reaction releases heat

  Enthalpy of reaction is negative (exothermic)
• Calculating enthalpy of reaction


  Sum of enthalpies of products - sum of enthalpies of reactants
• Using Hess's law


  1. Modify equations 1 and 2 to get equation 3

  2. Add enthalpies of equations 1 and 2 to get enthalpy of equation 3