Thermodynamics

1
Thermodynamics

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• Mr. Chan
• Northwestern University

2
Day 7/15

• 45 discuss HW/Lab/Quiz
• 45 QUIZ
• Picture
• 60 Thermochemistry
• Lunch
• 30 Molar Volume Lab
• 15 PreLab/Exit Slips
• 120 Labs

3
Energy and Heat what is energy?

• Energy Capacity for doing work or supplying
  heat
• Work from Physics
• Example Lifting books
• Types of Energy
• Kinetic Energy
• Potential Energy
• Chemical
• Energy stored in chemicals because of their
  composition
• Gravitational (Physics)

4
Energy and Heat what is heat?

• Heat (q)
• Transfer of energy due to a temperature
  difference
• Flows from warmer object to cooler object
• Thermochemistry
• Study of heat changes that occur during chemical
  reactions and changes in state
• Heat capacity
• Amount of heat it takes to change an objects
  temperature by exactly 1 degree celsius
• Water one of the highest, metals low heat
  capacities

5
What are calories?

• Heat needed for 1 g water 1 degree Celsius
• 1 Calorie 1000 calories
• 1 cal 4.186 Joules (SI unit), 1 kcal 4186
  Joules

6
Specific Heat

• Beach scenario
• Sand vs. Water
• Different materials have different capacities to
  absorb or release energy
• Different factors affect an objects heat
  capacity
• Specific Heat (Specific Heat Capacity)
• Amount of heat it takes to raise the temperature
  of 1 gram of the substance 1 degree Celsius
• List of Specific Heats

7
Equation C q/m (delta) T

• Practice
• Temperature of copper with mass 95.4 g changes
  from 25 C to 48 C when metal absorbs 849 Joules
  of heat. What is the specific heat of copper?

8
Calculating Heat Changes

• Calorimetry
• Measurement of heat change for chemical and
  physical processes
• Measuring in Labs
• Calorimeter
• Insulated device that measures absorption or
  release of heat in chemical or physical
  processes.
• What does transfer of heat depend on?

9
Q mC?T

• Mmass of water
• Cspecific heat
• ?Tchange in temperature
• System vs. Surroundings
• Total energy in universe conserved Law of
  conservation of energy
• Energy transfer between system and surroundings

10
Enthalpy

• Heat content of a system at constant pressure
• Heat change ?H
• Q ?H mC?T
• Negative exo, positive endo
• Practice 11-12
• Exothermic and Endothermic Processes
• Endothermic absorbs heat (? H is positive heat
  flows in)
• Exothermic releases heat (? H is negative
  heat flows out)

11
Calorimeter
12
Heats of Reaction

• Heat change for equation exactly as written
• 2NaHCO3 129 kJ ? Na2CO3 H2O CO2
• 2NaHCO3 ? Na2CO3 H2O CO2 ? H 129 kJ
• Standard conditions (1 atm pressure, 25 degrees
  Celsius)
• Heat of combustion
• Heat of reaction for burning of 1 mole of
  substance
• Examples
• Calculate the kJ of heat required to decompose
  2.24 mol of NaHCO3
• C 2S ? CS2 ?H 89.3kJ
• Calculate amount of heat absorbed when 5.66 g of
  CS2 forms

13
Latent Heats

• What about energy gained or released when
  temperature does not change?
• Heat of fusion/solidification
• Energy gained or released when one mole melting
  or when liquid turns to solid
• ? Hfusion - ? Hsolid
• Water fusion (6.01 kJ/mol), solidification
  (-6.01kJ/mol)
• Heat of vaporization/condensation
• Energy gained or released when boiling or when
  gas turns to liquid
• Water condensation -40.7 kJ/mol, vaporization
  40.7 kJ/mol

14
Practice

• How many grams of ice at 0 Celsius could be
  melted by the addition of 2.25 kJ of heat?
• How much heat is absorbed when 24.8 g of H2O (l)
  is converted to steam?

15
Heat of Solution

• Heat change that occurs when one mole of solute
  dissolves in a solvent
• Exo NaOH in water, handwarmers NaOH ? Na OH-
  ? H -445.1 kJ/mol
• Endo NH4NO3, cold packs
• Practice
• How much heat is released when 2.50 mol of NaOH
  is dissolved in water?

16
How do we determine heats of reaction for
equations?

• Many heats of reaction cannot be measured in
  laboratory because they take too slow.
• Use mathematical method
• Hesss Law of heat summation
• Calculating Heats of Reaction by combining
  reactions
• Reverse reaction reverse sign
• Multiply reaction coefficients multiply ?H
• Add Reactions, Add ?Hs
• Practice
• 2H2O2(l) yields 2H2O(l) O2(g)
• H2(g) O2(g) ? H2O2(l) ? H -187.9 kJ
• H2(g) 1/2O2(g) ? H2O(l) ? H -285.8 kJ

17
Determining Heats of Reaction from Standard Heats
of Reactions

• Standard Heat Formation ? Hf
• Change in enthalpy that accompanies the formation
  of one mole of compound from elements in their
  standard states
• Free elements in standard state is zero
• Others in appendix A.6 (note state of matter)
• ?H products reactants
• List of Standard Heats (Table 11.6)
• Practice
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l)
• 2SO2(g) O2(g) ? 2SO3(g)

18
Lab Specific Heat of a Metal (A/W 15)

• Objectives
• Measure specific heat of lead
• Identify an unknown metal from its specific heat
• Techniques
• Calorimetry
• Calculations with Specific Heat
• Identifying an Unknown
• Hypothesis
• Rank in order of increasing specific heat
  water, zinc, aluminum, lead
• Why is a foam coffee cup used instead of a beaker?

19
Lab Energy Content of Foods

• Teach.chem lesson PPT.

20
Exit Slips

• Midterm Self Assessment
• Give yourself 2 grades
• 1) Class grade (estimate using HW, LABS,
  QUIZZES, and TESTS)
• 2) Effort grade
• Write a few sentences justifying your grades and
  any discrepancies..
• Additional feedback regarding class/changes/sugges
  tions