Chapter 6- Thermochemistry

receptivetrucksMechanics

Oct 27, 2013 (3 years and 5 months ago)

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Chapter 6
-

Thermochemistry

Energy


Energy

is the capacity to do work or to
produce heat


Law of conservation of energy
-

energy can be
converted from one form to another but can
be neither created nor destroyed


The energy of the universe is
constant


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Types of Energy


Potential energy
-

the energy due to position
or composition


Potential energy in chemicals is present in the
bonds between atoms


Kinetic energy
-

energy due to the motion of
the object
-

depends on the mass of the object
and its velocity


Equation: KE = ½ mv
2


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Converting Energy


Energy can be converted from one form to
another


Two ways to transfer energy: Through


heat


and

work




Heat
-

involves the transfer of energy between two
objects due to a temperature difference.


Work
-

a force acting over a distance.


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Energy Transfer and Pathway


How an energy transfer is divided between
heat and work depends on the
pathway


Pathway
-

the specific conditions involved

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State Function


State function or State Property
-


A state
function refers to a property of the system
that depends on only its present state.


It does not depend on the system’s past or
future


In other words, it does not depend on

how
the system arrived at its present state


It depends only on the characteristics of
the present state


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Chemical Energy


There are 2 parts to anything we study: the

system



and the

surroundings



System
-

the part of the universe on which we
wish to focus attention


Surroundings
-

everything else in the universe

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Exothermic


Exothermic
-

energy flows

out of


the
system and into the surroundings


The potential energy stored in the chemical
bonds is being converted to thermal energy


The reactants had

more


energy stored in their
bonds than the products.


This means the bonds in the products are
stronger than those of the reactants


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Exothermic


Bonds with more potential energy are LESS
STABLE than bonds with less energy, so the
more reactive something is, the more energy
it can release.


More stable bonds have

less



potential
energy than less stable bonds.

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Endothermic


Endothermic
-

reaction absorbs energy from
the surroundings. Heat flows


into


the system.


The reactants have

less



potential
energy than the products


The reactants are


more stable


than the products

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10

Thermodynamics


Thermodynamics
-

the study of energy and its
interconversions


The first law of thermodynamics
-

the energy of the
universe is constant


Internal energy


the sum of the kinetic and potential
energies of all the particles in the system.


It can be changed by a flow of work, heat, or both


Represented by E


∆E = q + w


q is heat


w is work


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Thermodynamic Quantities


Thermodynamic quantities always consist of 2
parts: a number and a sign


Number
-

gives the magnitude of the change


Sign
-

indicates direction of the flow from the

system’s


point of view


For heat:


Exothermic reactions have a

negative


sign


Endothermic reactions have a

positive


sign


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Thermodynamic Quantities


For work:


if the system does work on the surroundings
then w is

negative





If the surrounding do work on the system, then w
is

positive




If work is done by a gas, this is

expansion



If work is done to a gas, this is

compression





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Work


Equation: w =
-
P∆V


remember that change in volume is final volume


initial volume


If gas expands, change in volume is positive, so
work is negative.



If gas contracts, change in volume is negative, so
work is positive


Note that the pressure is the external pressure.



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6

14

Practice Problems


1
-
4, 8, 20, 22; 26, 28


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15

Enthalpy


Enthalpy: H
H

= E + PV


E is the internal energy of the system


P is the pressure of the system


V is the volume of the system


Since internal energy, pressure, and volume
are state functions, the enthalpy is also a

state function




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Enthalpy


For a process carried out at constant pressure
and where the only work allowed is from a
volume change: ∆H =
q
p



This allows for heat of reaction and change in
enthalpy to be used interchangeably for systems
at constant pressure

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Enthalpy


For any chemical reaction, change is enthalpy can be
determined as follows:


∆H =
H
products



H
reactants




If products have a greater enthalpy than the reactants
than change in enthalpy will be


positive



and the reaction will be

endothermic



If the enthalpy of the products is less than that of the
reactants, change in enthalpy will be

negative


and the reaction will be

exothermic

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Calorimetry


Calorimeter


device used experimentally to
determine the heat associated with a chemical
reaction


Calorimetry
-

the science of measuring heat
-

based on observing the temperature change
when a body absorbs or discharges energy as
heat


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Heat Capacity


Heat capacity
-

measure of the amount of heat
required to raise the temperature of a
substance by 1 degree Celsius


C = heat absorbed/ increase in temperature


C represents heat capacity


This value varies depending on

amount of
the substance


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Heat Capacity


Specific Heat Capacity
-

the heat capacity per
gram of a substance.


Units are:

J/
o
C

*g or J/K* g


Molar Heat Capacity
-

the heat capacity per
mole of a substance



Units are:


J/
o
C

mol or J/K mol


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21

Constant Pressure
Calorimetry


Constant Pressure
Calorimetry
-

measurement
of heat using a simple
calorimetry
-

when
atmospheric pressure remains constant during
the process


Used for determining changes in enthalpy for reactions
that occur in

solution




Memorizing the specific heat capacity of water may
prove beneficial: 4.18 J/
o
C

g

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Constant P
Calorimetry


Energy released by the reaction = energy
absorbed by the solution


Equation: Energy released = s x m x ∆T


s = specific heat capacity of the substance


m = mass of that substance


∆T = Change in temperature


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23

Heat of Reaction


Heat of reaction is an

extensive



property because it depends on the amount of
substances

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24

Constant Volume
Calorimetry



∆E =
q
v


Energy released by reaction = ∆T x
C
calorimeter



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25

Hess’s Law


Hess’s Law
-

enthalpy is a state function
-

when
going from a particular set of reactants to a
particular set of products, the change in
enthalpy is

the same




whether
the reaction takes place in one step or in a series
of steps.

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26

Using Hess’s Law


Add together the reactions and the enthalpies of
those reactions to get the reaction you want and
the enthalpy of that overall reaction.


If a reaction is reversed, the sign of ∆H is

reversed


The magnitude of ∆H is directly proportional to
the quantities of reactants and products in a
reaction. If the coefficients in a balanced reaction
are multiplied by an integer the value of ∆H is

multiplied



by the same integer.


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27

Hints for using Hess’s Law


Hints for using Hess’s Law


Work

backward

from the required reaction,
use the reactants and products to decide how
to manipulate the other reactions


Reverse any reactions as needed to give
required reactants and products


Multiply reactions to give the correct numbers
of reactants and products.


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6

28

Practice Problems


Questions 58, 60, 64


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29

Standard Enthalpy of Formation


Standard enthalpy of formation
-


H
f
o
-

the
change in enthalpy that accompanies the
formation of one mole of a compound from its
elements with all substances in their

standard states



The degree symbol means that the process
occurred under standard conditions


Standard state
-

reference state

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30

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31

Standard States


Compound


For a gas, pressure is exactly 1 atmosphere.


For a solution, concentration is exactly 1 molar.


Pure substance (liquid or solid)


Element


The form [N
2
(
g
), K(
s
)] in which it exists under
conditions of
1
atm

and 25
°
C
.

Enthalpy


Enthalpies of formation are always given

per
mole

of the product with the product in its standard
state


The enthalpy change for a given reaction can be
calculated by subtracting the enthalpies of formation
of the reactants from the enthalpies of formation of
the products


Equation: ∆
H
f
o
reaction

= ∑
n
p
∆H
f
o
products

-


n
r
∆H
f
o
reactants



Elements are not included because elements require
no change in form


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6

32

Practice Problems


66, 70, 72

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6

33

Energy from the Sun


The energy in woody plants, coal, petroleum
and natural gas originally came from

the
sun




We obtain the energy by

burning



the
plants or the decay products of those plants

-

fossil fuels


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34

Petroleum


Petroleum
-

thick dark liquid composed mostly
of compounds called hydrocarbons, which
contain
H and C


Separation of petroleum occurs as different
substances are boiled off in a process called:

distillation




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35

Natural Gas and Coal


Natural gas
-

usually near petroleum deposits
-

consists of mostly methane (CH
4
) and contains
ethane, propane, and butane


Coal
-

will become more important as oil is
used up


Expensive and dangerous to mine


Burning causes air pollution that leads to acid rain


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36

Carbon Dioxide


Molecules in the atmosphere

H
2

O, CO
2

,
CH
4
absorb infrared radiation and radiate it
back toward earth, so the earth is much
warmer than it would be without the
atmosphere


http://www.teachersdomain.org/resource/ph
y03.sci.phys.matter.co2/


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37

Alternative Energy Sources


3 things to consider when finding new energy
sources


Economic, climatic, and supply factors


Potential Sources
-

sun, nuclear processes,
biomass, and synthetic fuels


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38

Alternative Energy


Syngas
-

use coal gasification
-

reduce the size
of molecules by treating coal with oxygen and
steam at high temperatures to break many C
-
C
bonds. The bonds are replaced by C
-
H and C
-
O bonds as coal reacts with water and oxygen


The product is a mixture of CO and H
2


CO and H
2

can also be converted to methanol,
which is used to produce synthetic fibers and
plastics and is used as a fuel


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39

Hydrogen


H
2

as fuel


The heat of combustion for hydrogen is 2.5 times
that of natural gas


The only product of burning hydrogen is water


3 problems: cost of production, storage, transport


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40

Hydrogen as Fuel


Hydrogen is abundant but doesn’t exist as the
free gas, so treating natural gas with steam
produced hydrogen


This reaction is highly endothermic, so it is


not an efficient way

to obtain hydrogen for
fuel


Storage
-

hydrogen decomposes into atoms
when in contact with metal, enter the metal,
and make the metal brittle


http://www.teachersdomain.org/resource/eng06.sci.
engin.systems.fuelcells/



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6

41

Practice Problems


76