Transcript Chapter 6

Chapter 6
Thermochemistry
Energy
• Is important macroscopically and
microscopically
• Def: the capacity to do work or produce
heat
Law of Conservation of Energy
• Energy can neither be created nor
destroyed, but can be converted from one
form to another.
• The energy of the universe is constant.
• FIRST LAW OF THERMODYNAMICS
Classifications of Energy
• Potential Energy: energy due to position or
composition
– Ex: dam water, attractive/repulsive forces
• Kinetic Energy: energy due to motion,
depends on mass (m) and velocity (v)
– KE = (1/2)mv2
• Energy can be converted between these
Energy Transfer
• Some energy can be lost as heat (ex: frictional
heating), represented by q
• Heat vs. Temperature: TEMPERATURE reflects
movement of particles. HEAT deals with
transfer of energy between two objects due to a
temperature difference.
• Energy can also be transferred through work
(force activing over a distance)
Important Vocabulary
• Pathway determines
how energy changes
to heat/work
– Includes condition of
the surface
• Total energy
transferred will be
constant, amounts of
heat/work will differ
State Function/Property
• Property of the system that depends only on
its present state (not past or future)
• Changes in the state properties when
switching from one state to another is
independent of the particular pathway taken
between the two states.
• Example: energy is a state function, but work
and heat are not.
– Internal energy, pressure, volume, enthalpy
Heat
• Represented by “q”
• Flows from warm to cold
System vs. Surroundings
• System is what you are focusing on,
surroundings is everything else.
• Energy lost/gained by the system = energy
gained/lost by the surroundings
Endothermic
• System gains heat, surroundings
cool
• Heat INTO (“endo”) system
• Example: Ice melting
• “q” is positive
• Heat is a reactant
N2(g) + O2(g) + HEAT(kJ) --> 2NO(g)
Exothermic
•
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•
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System loses heat to surrounding
“Exo” like EXIT
Example: fire
“q” is negative
Heat is a product
CH4(g) + 2O2(g) --> CO2(g) + 2H2O(g) + HEAT(kJ)
Potential Energy
• Energy is stored in chemical bonds as potential
energy. When bonds are broken (requires
energy) and formed (releases energy), it
changes the potential energy.
Units for Energy
Two common units:
–The “calorie” (old school)
–The “joule” (metric)
1 cal = 4.18 J
J = kg•m2
s2
101.3 J = 1 L atm
Energy Stoichiometry
• Energy can be added into stoichiometry
equations…
C6H12O6 + 6O2 --> 6CO2 + 6H2O + 2800kJ
• You can substitute it in as part of the mole to
mole ratio!
• How much heat is given off when 3.72 moles of
oxygen react with glucose?
• Answer: 1736 kJ
Internal Energy
• Represented by E
• Sum of kinetic and potential energy in the
system
∆E = q + w
• ∆E = change in system’s internal energy
• q = heat (usually in J)
• w = work (usually units are L atm which can
be converted to J)
Example
• Calculate the change in energy of
the system if 38.9 J of work is done
by the system with an associated
heat loss of 16.2 J.
• Answer: 55.1 J
Expansion vs. Compression
w = -P∆V
• w = work
• P = pressure
• ∆V = change in volume
Results…
• If a gas is expanding, ∆V is positive
• If a gas is compressed, ∆V is negative
• When w is negative, work if flowing out
of the system (into surroundings)
• When w is positive, work is flowing into
the system
Example
• A piston is compressed from a
volume of 8.3 L to 2.8 L against a
constant pressure of 1.9 atm. In the
process, there is a heat gain by the
system of 350 J. Calculate the
change in energy of the system.
• Answer: 1400 J
Enthalpy
• Represented by “H”
H = E + PV
• E = internal energy
• P = pressure of the system
• V = volume of the system
This means at constant pressure…
• ∆H = q
• Negative ∆H is exothermic
• Positive ∆H is endothermic
∆H = Hproducts - Hreactants
Calorimetry
• Energy can’t be created nor destroyed.
• If one object loses heat, another object
must gain that heat.
• If a cool object is placed into a hot one,
the hot object gives energy to the cool one
until they arrive at the same final temp.
Conservation of Energy
• The amount of heat lost by the
hotter object equals the amount of
heat gained by the cooler object:
qgained = -qlost
(one is system, other surroundings)
Specific Heat Capacity (J/g°C)
• The amount of energy needed to
raise the temp. of 1 gram of an
object by 1°C.
• A high specific heat means the
object requires a lot of energy to
change temp. (pg. 245)
• If unit is J/mol°C or J/molK, then
it’s called molar heat capacity
Constant Pressure Calorimeter
• Also have
constant volume
calorimeters
q = m*c*DT
q = amount of heat gained
(negative if lost)
m = mass of object (grams)
c = specific heat of object
DT = change in temperature (Tfinal
– Tinitial)
q = m*c*DT
• This equation only works
when the temperature is
changing.
• This is used in calorimetry!
Example
• A bar of iron at 21.0°C is heated
to 83.5°C. If the iron’s mass is
551 grams, how much heat was
added?
• Answer: -1.5X104 J
Calorimetry Example
• A 36.9 g sample of metal is heated to 100.0ºC,
and then added to a calorimeter containing
141.5 g of water at 23.1ºC. The temperature of
the water rises to a maximum of 25.2ºC before
cooling back down. What is the specific heat of
the metal?
• Answer: 0.45 J/gºC