Chapter 16 Thermochemistry

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Chapter 16Thermochemistry

• Pioneer High School
• Ms. Julia Bermudez

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Chapter 16

• OBJECTIVES
• Explain the units for energy.
• Describe the Difference between Temperature and
  Heat

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Chapter 16

• OBJECTIVES
• Exothermic Endothermic
• Solve heat change problems

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Energy and Heat

• Energy - the ability to do work or produced heat
• It comes in two forms
• Potential Energy the energy due to the
  composition or position of an object
• Kinetic Energy the energy of motion

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Energy and Heat

• Chemical potential energy the energy stored in
  a substance because of its composition
• Sometimes potential energy is released as heat.
• Heat is energy that is in the process of
  flowing from a warm object to a cooler object

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Energy and Heat

• Whats the difference between
• Heat and Temperature
• Heat is the flow of energy from warm to a cold
• Temperature is the measurement of the average
  kinetic energy of an object

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Energy and Heat

• Measurement of Energy
• calorie (cal) the amount of energy required to
  heat up 1.0 g of water 1.0 oC
• Calorie (Cal) 1000 cal 1 Cal
• This is the food Calories
• Joules (J) 4.184 J 1cal
• This is the metric unit for heat

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Energy and Heat

• Conversions
• What you need to know!
• 1000 cal 1 Cal
• 4.184 J 1 cal
• Table 16-1 pg. 491

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Energy and Heat

• Example 1
• A Snickers bar is 180 Cal how many Joules is in
  a Snickers?
• 180 Cal

1000 cal
4.184 J
1 Cal
1 cal
10

• 753,120 Joules
• in
• a Snickers!

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Energy and Heat

• Example 2
• You burn a Three Musketeers Bar to heat up some
  water. If the water absorbs 45,000J from the
  bar, how many Calories is a Three Musketeers Bar?
• 45,000J

1 cal
1 Cal
4.184 J
1000 cal
12

• 10.76Cal
• in a
• 3 Musketeers!

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Practice Problems!

• Pg. 493 1-3
• Homework!
• Pg. 525 75-77

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Exothermic and Endothermic Processes

• All chemical reactions and changes in physical
  state involve either
• Release of heat - Exothermic
• Absorption of heat - Endothermic

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Exothermic and Endothermic Processes

• In studying heat changes, think of defining these
  two parts
• The System - the part of the universe on which
  you focus your attention
• The Surroundings - includes everything else in
  the universe

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Demo Time!
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Exothermic and Endothermic Processes

• During Exothermic Reactions energy is Exiting
  from the system.
• ex Burning methane
• During Endothermic Reactions energy is taken In
  from the surroundings.
• ex Melting of ice

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Specific Heat

• Specific Heat (Cp) - the amount of energy it
  takes to raise 1.0 gram of a substance by 1 oC

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Specific Heat

• Water has a HUGE value compared to other
  chemicals
• Metals have low values
• Because of these differences, Metal heats up
  quicker than water and cools down faster too.

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Specific Heat

• Cp of Water 4.184 J/(g oC)
• Thus, for water
• it takes a long time to heat up,
• it takes a long time to cool off!
• 100oC water is hotter than 100oC alcohol because
  it hols more energy

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Specific Heat

• To calculate heat, use the formula
• q Cp m T
• q heat in Joules (J)
• Cp Specific Heat J/(g oC)
• m mass (g)
• T temperature (oC)

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Specific Heat

• To calculate a change in heat, either cooling
  down or heating up use the formula
• q Cp m DT
• DT the change in temperature
• DT (Tf Ti)
• Tf final temp. Ti initial temp.

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Specific Heat

• Example 1
• Find how much heat is in a 100g glass of 25oC
  water.
• q m Cp T
• m 100 g
  
  
• Cp of H2O 4.184 J/g oC
• T 25 oC

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Specific Heat

• Example 1
• Find how much heat is in a 100g glass of 25oC
  water.
• q Cp m T
• q (4.184 J/g oC)(100 g)(25 oC)
• q 10,460 J

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Specific Heat

• Example 2
• Find how much heat is lost when a 100g glass of
  25oC water cools down to 10oC.
• q m Cp ?T
• Cp of H2O 4.184 J/goC
• m 100 g
• Ti 25 oC
• Tf 10 oC

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Specific Heat

• Example 2
• Find how much heat is lost when a 100g glass of
  25oC water cools down to 10oC.
• q Cp m DT
• q (4.184J/goC)(100g)(10oC25oC)
• q - 6276 J

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Homework

• Chemical Thermodynamics Worksheet 1 1-6 only

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Energy and Change of State

• OBJECTIVES
• Describe Melting and Boiling
• Solve Latent Heat of Fusion and Latent Heat of
  Vaporization problems
• Solve change of state problems

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Heat and Change of State

• During all changes in state energy will either be
  gained or lost
• Two Changes of State -
• Melting/Freezing the phase change between solid
  and liquid states.
• Vaporization/Condensation the phase change
  between liquid and gas states.

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Heat and Change of State

• Melting
• - As energy is gained by a solid, the particles
  gain kinetic energy.
• - The bonds within the solid will break allowing
  the particles to move around more making a liquid
• Freezing
• - Energy is released by the liquid which slows
  down the particles to a vibration

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Heat and Change of State

• - The energy involved in melting/freezing is
  called the Latent Heat of Fusion (DHfus)
• - Water has its own value
• ?Hfus 334 J/g
• During melting/freezing there is no change in
  temp. All the energy is spent to break bonds or
  slow down particles.

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Heat and Change of State

• To calculate heat required for melting/freezing
  use the formula -
• q m DHfus
• m mass
• ?Hfus Latent Heat of fusion
• Changes for each substance

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Energy and Change of State

• Example 1
• Find how much heat is required to melt 65g of
  water?
• q m DHfus
• q (65g)(334 J/g)
• q 21,710 J

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Energy and Change of State

• Vaporization
• As energy is gained by a liquid the particles
  gain kinetic energy.
• The particles gain so much energy they break the
  water surface and change to a gas
• Condensation
• - Energy is released and the gas particles slow
  down and group together creating a liquid

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Energy and Change of State

• The energy involved in vaporization/condensation
  is called the Latent Heat of Vaporization (DHvap)
• - Water has its own value
• ?Hvap 2260 J/g
• Again during this process all the energy is spent
  to break bonds or slow down particles.

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Heat and Change of State

• To calculate heat required for boiling, use the
  formula -
• q m DHvap
• m mass
• ?Hvap Latent Heat of vaporization

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Energy and Change of State

• Example 2
• Find how much heat is required to boil 65g of
  water?
• q m DHvap
• q (65g)(2260 J/g)
• q 146,900 J

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• Classwork
• How much Heat is required to vaporize 343 g of
  Ethanol?
• ?Hvap 38.6 kJ/g
• How much heat is evolved when 1255 g of water
  condenses to a liquid at 100C?
• A sample of ammonia liberates 5.66kJ of heat as
  it solidifies at its melting point. What is the
  mass? ?Hfus 5.66 kJ/g

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Energy and Change of State

• It is possible to calculate the total heat of a
  substance starting at a solid state all the way
  to a gaseous state.

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Energy and Change of State

• Lets take a look at Ice Melting
• As energy is added the temp of the ice rises

TEMPERATURE
ENERGY
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Energy and Change of State

• Soon the ice will begin to melt and the temp does
  not change

TEMPERATURE
ENERGY
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Energy and Change of State

• Once the ice completely melts we get a liquid
  that begins to rise in temp

TEMPERATURE
ENERGY
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Energy and Change of State

• Eventually the liquid will gain energy and begin
  the boil. The temp will not change

TEMPERATURE
ENERGY
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Energy and Change of State

• In a sealed contained the gas can rise in temp.

TEMPERATURE
ENERGY
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Energy and Change of State

• On Each Diagonal line we have a different phase
  and the temp changes.
• We use the Equation
• q m Cp ?T

TEMPERATURE
GAS
LIQUID
SOLID
ENERGY
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Energy and Change of State

• On each flat line energy is used on the particles
  to speed them up or slow them down. The temp
  does not change
• We use the Equation
• q m ?Hvap
• or
• q m ?Hfus

TEMPERATURE
GAS
VAPOR/COND
LIQUID
MELT/FREEZE
SOLID
ENERGY
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Energy and Change of State

• Example 3
• Calculate the total amount of energy needed to
  melt 10 g of ice at -10C to water at 25C.

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Homework

• Chemical Thermodynamics Worksheet 2 - all
  problems