What happens to thermal (heat) energy?

							Thermodynamics
 Energy is neither created or
destroyed during chemical or
  physical changes, but it is
 transformed from one form
         to another.

        Euniverse = 0
 TYPES of ENERGY

 Kinetic                       Potential
 Mechanical                    Gravitational
 Thermal                       Electrostatic
 Electrical                    Chemical
 Radiant
Energy Conversion Examples:
       1. dropping a rock
       2. using a flashlight
       3. driving a car
           SYSTEMS and SURROUNDINGS
           System:        The thing under study
           Surroundings:  Everything else in the universe

           Energy transfer between system and surroundings:




Endo: heat added to system                     Exo: heat released by system
          HEAT: What happens to
           thermal (heat) energy?
Three possibilities:
• Warms another object
• Causes a change of state
• Is used in an endothermic reaction
Temperature Changes from Heat Exchange




  Example 1: 5 g wood at 0 oC     +   5 g wood at 100 oC
  Example 2: 10 g wood at 0 oC    +   5 g wood at 100 oC
  Example 3: 5 g copper at 0 oC   +   5 g copper at 100 oC
  Example 4: 5 g wood at 0 oC     +   5 g copper at 100 oC

Choices:
1: 0 oC 2: 33 oC   3: 50 oC   4. 67 oC   5: 100 oC 6: other
What happens to thermal (heat) energy?

  When objects of different temperature meet:

  • Warmer object cools
  • Cooler object warms
  • Thermal energy is transferred

  • qwarmer = -qcooler
Quantitative: Calculating Heat Exchange: Specific Heat Capacity
        Specific Heat Capacity
The energy required to heat one gram of a
 substance by 1 oC.

Usefulness: #J transferred = S.H. x #g x T

How much energy is used to heat 250 g water from
 17 oC to 100 oC?
What happens to thermal (heat) energy?

  When objects of different temperature meet:

  • Warmer object cools
  • Cooler object warms
  • Thermal energy is transferred
  • qwarmer = -qcooler
  specific heat x mass x T = specific heat x mass x T
      warmer object                 cooler object
Heat transfer between substances:


                          q wood =




                          q Cu =




Specific heats: Cu = 0.385 J/goC Wood = 1.8 J/goC
Conceptually Easy Example with Annoying Algebra:

If we mix 250 g H2O at 95 oC with 50 g H2O at 5 oC,
what will the final temperature be?
   Thermal Energy and Phase Changes

First: What happens?
   Thermal Energy and Phase Changes
First: What happens?
   Thermal Energy and Phase Changes
First: What happens?
But what’s really happening?
Warming:
• Molecules move more rapidly
• Kinetic Energy increases
• Temperature increases


Melting/Boiling:
• Molecules do NOT move more rapidly
• Temperature remains constant
• Intermolecular bonds are broken
• Chemical potential energy (enthalpy) increases
Energy and Phase Changes:
           Quantitative Treatment
Melting:

Heat of Fusion (Hfus) for Water: 333 J/g


Boiling:

Heat of Vaporization (Hvap) for Water: 2256 J/g
Total Quantitative Analysis
Convert 40.0 g of ice at –30 oC to steam at 125 oC

Warm ice: (Specific heat = 2.06 J/g-oC)



Melt ice:




Warm water (s.h. = 4.18 J/g-oC)
Total Quantitative Analysis
Convert 40.0 g of ice at –30 oC to steam at 125 oC



Boil water:




Warm steam (s.h. = 1.92 J/g-oC)
Energy and Chemical Reactions

Lots of different types of energy.

We use Enthalpy:

Heat exchanged under constant
pressure.
Energy/Enthalpy Diagrams
 Some Examples of Enthalpy
         Change




2 C(s) + 2 H2(g)  C2H4(g) H = +52 kJ
Enthalpy Change and Chemical Reactions

 H is usually more complicated, due to solvent and
 solid interactions.

 So, we measure H experimentally.

 Calorimetry

 Run reaction in a way that the heat exchanged
 can be measured. Use a “calorimeter.”
Bomb Calorimetry Experiment

 N2H4 + 3 O2  2 NO2 + 2 H2O

 Energy released = E absorbed by water +
                       E absorbed by calorimeter

 Ewater =

 Ecalorimeter =
                                                   0.500 g N2H4
 Total E =
                                                   600 g water
 H = energy/moles =                               420 J/oC
Enthalpy Change and Bond Energies


 H = energy needed to break bonds – energy released forming bonds

 Example: formation of water:




    H = [498 + (2 x 436)] – [4 x 436] kJ = -482 kJ
General Rule:
SO2 + ½ O2  SO3 dH = -98.9 kJ

2 SO3  2 SO2 + O2   dH = ?
Hess’s Law



             Enthalpy is a
             State Function.
Thermochemistry Lab Calculations
Goal: What is H for the formation of MgO from Mg(s) and O2(g)?
         Mg(s) + ½ O2(g)  MgO(s)                H = ? kJ/mol
Data:
From lab measurements:
         Mg(s) + 2 H+(aq)  Mg2+(aq) + H2(g)               H1 = ___________ kJ/mol

         MgO(s) + 2 H+(aq)  Mg2+(aq) + H2O(l)             H2 = ___________ kJ/mol

From a table:     H2(g) + ½ O2(g)  H2O(l)                 H3 = -285.8 kJ/mol

Task: Find a way to add these three reactions to get the desired reaction. Manipulate the
H values as needed, and add them.
Calculating Heat Production
Heat of Formation
Heat of Formation: The general idea
Find the enthalpy change for burning
            ethyl alcohol