Transcript Chapter 17 - Thermochemistry

Chapter 5
“Thermochemistry”
Mr Daniel : OHS AP Chemistry
Credit to:
Charles Page High School
Stephen L. Cotton
Section 11.1
The Flow of Energy – Heat and Work

OBJECTIVES:
• Explain how energy, heat, and work are
related.
• Classify processes as either exothermic or
endothermic.
• Identify the units used to measure heat
transfer
• Distinguish between heat capacity and
specific heat.
2
Energy Transformations
Thermochemistry – study of heat changes that
occur during chemical and physical changes.
 Energy - capacity for doing work or supplying
heat
• weightless, odorless, tasteless
• Energy stored within chemical substances is
called chemical potential energy

Gasoline contains a significant amount of
chemical potential energy
3
Energy Transformations

Heat (“q”) is energy that transfers from one
object to another, because of a temperature
difference between them.
• only changes of heat can be detected!
• flows from warmer  cooler object
All chemical reactions and changes in physical
state involve either:
a) release of heat, or
b) absorption of heat
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Exothermic and Endothermic Processes
When studying heat changes there are two
important categories:
• the system - the part of the universe on
which you focus your attention (a test tube,
a room, a molecule, the Earth etc.)
• the surroundings - includes everything else
in the universe
 Together, the system and it’s surroundings
constitute the universe

Exothermic and Endothermic Processes
 Thermochemistry
is concerned with the flow
of heat from the system to it’s surroundings,
and vice-versa.
 The Law of Conservation of Energy states
that in any chemical or physical process,
energy is neither created nor destroyed.
• All the energy is accounted
for as work, stored energy,
or heat.
Exothermic and Endothermic Processes
Exothermic and Endothermic Processes
 Exothermic:
Heat flowing out of a system into
it’s surroundings
• defined as negative change
q has a negative value
• The system loses heat
as the surroundings heat up

Exothermic and Endothermic

Every reaction has an energy change
associated with it. For example, the…
** Gummy Bear Sacrifice!! ** (next week!!!)
 Exothermic reactions release energy, usually in
the form of heat.
 Endothermic reactions absorb energy
Chemical Energy is stored
in the form of bonds
between atoms.
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Units for Measuring Heat Flow
1) A calorie is defined as the quantity of heat
needed to raise the temperature of 1 g of pure
water 1 oC.
a Calorie, (written with a capital C), is 1000
calories and refers to the energy in food
(100 Calorie candy bar is 100,000 calories)
1 Calorie = 1 kilocalorie = 1000 calories
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Units for Measuring Heat Flow
2) The Joule is the SI unit of heat and energy
and is related to the calorie:
4.184 J = 1 cal
Named after James Prescott Joule
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Heat Capacity and Specific Heat

Heat Capacity (Cp) - the amount of heat
needed to increase the temperature of an
entire object exactly 1 oC
 Cop = Joules/oC
Heat capacity is
dependent on both
the object’s mass and
its chemical
composition
Heat Capacity and Specific Heat
 Specific
Heat Capacity (c) - the amount of heat
it takes to raise the temperature of 1 gram of
the substance by 1 oC
Often called simply “Specific Heat”
c = Joules/ gram oC
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Table of Specific Heats
Heat Capacity and Specific Heat
Water has a HUGE specific heat value, when is
compared to other chemicals:
C (H2O) = 4.18 J/(g oC), or
C (H O) = 1.00 cal/(g oC)

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• It heats slowly and
• It cools slowly
Consequently, water maintains a relatively
constant temperature!
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Heat Capacity and Specific Heat

Heat calculations for various materials are done
with the formula:
 q = mass (g) x c x T
•  q = change in heat
• c = specific heat
• T = change in temperature
Units are either J/(g oC) or cal/(g oC)
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Section 11.2
Measuring and Expressing
Enthalpy Changes
OBJECTIVES:
• Describe how calorimeters are used to
measure heat flow.
• Construct thermochemical equations.
• Solve for enthalpy changes in chemical
reactions by using heats of reaction.
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Calorimetry

Calorimetry - the precise measurement of the
heat into or out of a system for chemical and
physical processes.

The device used to measure the absorption
or release of heat in chemical or physical
processes is called a Calorimeter
Calorimetry

Foam cups are excellent heat insulators, and
are commonly used as simple calorimeters

For systems at constant pressure, the heat
content is the same as a property called
Enthalpy (H) of the system
A foam cup
calorimeter –
here, two
cups are
nestled
together for
better
insulation
Calorimetry
Changes in heat = Changes in enthalpy = H
 q = H These terms are often used
interchangeably
 Thus, q = H = m x c x T

H is negative for an exothermic process
 H is positive for an endothermic process

Calorimetry

Calorimetry experiments can be performed at a
constant volume using a device called a “bomb
calorimeter” - a closed system
• Used by nutritionists to measure energy
content of food
A bomb calorimeter
A Bomb Calorimeter
Energy
C + O2 → CO2 + 395 kJ
C + O2
395kJ
CO2
Reactants

Products
∆ H= -395 kJ
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Exothermic
 In
an exothermic rxn the products are
lower in energy than the reactants and
energy is released.
 ΔHrxn = -395 kJ
• The negative sign indicates that
energy is released.
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CaCO
→ +CaO
CaCO
CaO
CO2+ CO2
3 + 176
3 → kJ
Energy
CaO + CO2
176 kJ
CaCO3
Reactants

Products
∆H = + 176 kJ
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Endothermic
The products are higher in energy than the
reactants and energy is absorbed.
 ΔHrxn = +176 kJ
• The positive sign indicates that energy is
absorbed

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Heat of Reaction
A
heat of reaction is the heat change for the
equation.
• The physical state of reactants and
products must also be given.
• Standard conditions for the reaction is
101.3 kPa (1 atm.) and 25 oC (different
from STP)
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Thermochemical Equations
An equation that includes energy is called a
thermochemical equation

CH4 + 2O2  CO2 + 2H2O + 802.2 kJ
• 1 mole of CH4 releases 802.2 kJ of energy.
• When you make 802.2 kJ you also make
2 moles of water
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CH4 + 2 O2  CO2 + 2 H2O + 802.2 kJ
 If
10. 3 grams of CH4 are burned
completely, how much heat will be
produced?
10. 3 g CH4
1 mol CH4
16.05 g CH4
802.2 kJ
1 mol CH4
= 514 kJ
Thus, ΔH = - 514 kJ for this reaction
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Summary, so far...
Enthalpy
Enthalpy: The heat content a substance has at
a given temperature and pressure
 Can’t be measured directly because there is no
set starting point
 The reactants start with a heat content
 The products end up with a heat content
 So we can measure how much enthalpy
changes
H (delta H) is the measure of change in heat

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Enthalpy
Enthalpy: Amount of heat in a system (H)
.

If heat is released, H is negative (exothermic)
 If heat is absorbed, H is positive (endothermic)

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Energy
Change is down
ΔH is <0
Exothermic
Reactants 
Products
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Energy
Change is up
ΔH is > 0
Endothermic
Reactants 
Products
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Heat of Reaction
Heat of Reaction: The heat that is released or
absorbed in a chemical reaction
 Equivalent to H
C + O2(g)  CO2(g) + 393.5 kJ
C + O2(g)  CO2(g)
H = -393.5 kJ
 In a thermochemical equation, it is important to
indicate the physical state
 H2(g) + 1/2O2 (g) H2O(g) H = -241.8 kJ
 H2(g) + 1/2O2 (g) H2O(l) H = -285.8 kJ

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Heat of Combustion

The heat from the reaction that completely burns
1 mole of a substance:
C + O2(g)  CO2(g) + 393.5 kJ
C + O2(g)  CO2(g)
H = -393.5 kJ
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Section 17.3
Heat in Changes of State
 OBJECTIVES:
• Classify the enthalpy change
that occurs when a
substance melts, freezes,
boils, condenses, or
dissolves.
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Section 17.3
Heat in Changes of State
 OBJECTIVES:
• Solve for the enthalpy change that
occurs when a substance melts,
freezes, boils condenses or dissolves.
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Heat in Changes of State
 Heat
of Fusion (Hfus) : the heat absorbed
by a substance in melting from a solid to
a liquid
 For H2O: Hfus = 334 J/ g
 Molar Hfus = 6.01 kJ/ mol
 q = mass x Hfus (there is no temperature change)
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 Molar
Heat of Solidification (Hsolid) =
heat lost when one mole of liquid
solidifies (or freezes) to a solid
 For H2O: Hfus = -334 J/ g
 Molar Hfus = -6.01 kJ/ mol
 q = mass x Hsolid (no temperature change)
Heat in Changes of State
Heat
absorbed by a melting
solid is equal to heat lost
when a liquid solidifies
• Thus, Hfus = -Hsolid
Note Table 17.3, page 522
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Heats of Vaporization and Condensation
 When
liquids absorb heat at their boiling
points, they become vapors.
 Molar Heat of Vaporization (Hvap) = the
amount of heat necessary to vaporize
one mole of a given liquid.
 q = mass x Hvap (no temperature change)
For H2O: Hfus = 2260 J/ g
 Molar Hfus = 40.7 kJ/ mol

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Heats of Vaporization and
Condensation
Condensation
is the opposite of
vaporization.
Molar Heat of Condensation (Hcond)
= amount of heat released when
one mole of vapor condenses to a
liquid
 Hvap = - Hcond
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Heats of Vaporization and Condensation
 Note Figure 17.10, page 523
 The large values for water Hvap and
Hcond are the reason hot vapors such
as steam is very dangerous
• You can receive a scalding burn from
steam when the heat of condensation
is released!
H20(g)  H20(l)
Hcond = - 40.7kJ/mol
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Heat of Solution (Ch 11)
Heat
changes can also occur when
a solute dissolves in a solvent.
Molar Heat of Solution (Hsoln) =
heat change caused by dissolution
of one mole of substance
Sodium hydroxide provides a good
example of an exothermic molar
heat of solution:
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Heat of Solution
NaOH(s)
H2O(l)

Na1+(aq) + OH1-(aq)
Hsoln = - 445.1 kJ/mol
The heat is released as the ions
separate (by dissolving) and
interact with water, releasing 445.1
kJ of heat as Hsoln -thus becoming
so hot it steams!
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