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Thermodynamics

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Title: 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..
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    tions
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