Chapter 6
Thermochemistry
-study of heat changes that occur during
chemical reactions, study of relationships
between chemistry and energy
energy- ability to do work or supply heat
-unlike matter, energy is weightless, odorless,
tasteless
-only way to detect energy is by its effects
work- result of a force acting through a distance
heat- flow of energy from one object to another
because of a temperature difference
*energy is something that an object has
*heat and work are ways in which objects
exchange energy
kinetic energy(KE)- energy of motion
-faster = more kinetic energy
potential energy(PE)- energy of position
-a rock at the top of a hill has more PE than a
rock at the bottom
thermal energy- energy associated with
temperature of an object
-higher temp = more thermal energy
-type of KE because it arises from motion of
particles
Law of Conservation of Energy
-energy can neither be created nor destroyed
-energy can be transferred from one object to
another and from one form to another
All chemical reactions involve the release or
absorption of heat:
system- what is being studied
surroundings- everything else around the system
in nature
universe- system and surroundings together
*the system either loses energy to or gains
energy from the surroundings
Heat Flow two directions
1) surroundings  system
endothermic- system gains heat as surroundings
cool, system absorbs heat
2) system  surroundings
exothermic- system loses heat as surroundings
heat up, system releases heat
Units of Energy
joule(J) - SI unit for heat/energy
-can use kJ
1kJ = 1000J
calorie(cal) - amount of energy required to raise
the temp of 1g of pure water 1°C
Calorie(Cal) - nutritional unit
**1Cal = 1kcal = 1000cal
**1J = 0.2390 cal
4.184J = 1cal
Convert:
1) 1656.70J cal
2) 483.12cal  J
3) 0.56721Cal  J
thermodynamics- general study of energy and its
interconversions
The First Law of Thermodynamics
-states that the total energy of the universe is
constant
internal energy(E)- sum of the potential and
kinetic energies of all the particles that
compose the system
**E = PE + KE
-internal energy is a state function
-value depends only on the state of the
system (initial and final values) and not on
how the system arrived at that state
∆E = Efinal - Einitial
∆ = change in
-in a chemical reaction:
∆E = Eproducts - Ereactants
*If the reactants have a higher internal energy
than the products, ∆Esys is negative and the
energy flows out of the system into the
surroundings
*If the reactants have a lower internal energy
than the products, ∆Esys is positive and energy
flows into the system from the surroundings
∆E = the amount of heat transferred and the
work done
∆E = q + w
q= heat
w= work
q (heat)
w (work)
∆E (change in
internal energy)
+
system gains
system loses
thermal energy thermal energy
+
work done on
work done by
the system
the system
+
energy flows energy flows out
into the system
of the system
Ex- Identify each of the following energy
exchanges as heat or work and determine
whether the sign of heat or work (relative to
the system) is positive or negative.
1) An ice cube melts and cools the surrounding
beverage. (ice cube is the system)
*heat, sign is positive
2) A metal cylinder is rolled up a ramp. (the
cylinder is the system; assume no friction)
*work, sign is positive
3) Steam condenses on skin, causing a burn.
(the steam is the system)
*heat, sign is negative
Ex- A cylinder and piston assembly (defined as
the system) is warmed by an external flame.
The contents of the cylinder expand, doing
work on the surroundings pushing the piston
outward. If the system absorbs 732J of heat
and does 628J of work, what is the change in
external energy?
∆E = q + w
q= +732J
w= -628J
∆E = +732J + (-628J) = +104.00J
Try page 246- Ex 6.1 and For Practice 6.1
*Notice the distinction between temperature and
heat
-temperature is a measure of thermal energy
-heat is the transfer of thermal energy
thermal equilibrium- when there is no additional
heat transfer because system and surroundings
are same temperature
heat capacity (C) - the quantity of heat required
to change the temperature of a system 1°C
specific heat capacity (Cs) - the amount of heat
needed to raise the temperature of 1g of a
substance by 1°C
-the higher the specific heat capacity, the more
energy is required to raise the temperature
-metals tend to have fairly low specific heats
because they heat up easily
-substances like water have very high specific
heats because they take a long time to heat up
q = (m)(Cs)(ΔT)
q = heat (J)
m = mass (g)
Cs = specific heat capacity J/g°C *will be given
to you if not solving for- page 247
ΔT = temp change Tfinal – Tinitial °C
calorimetry- the accurate and precise
measurement of heat change for chemical and
physical processes
In calorimetry:
heat released = heat absorbed
-to measure changes accurately, must be carried
out in an insulated container (calorimeter)
ex- foam cup
bomb calorimeter- equipment designed to
measure ∆E for combustion reactions, volume
is constant
page 251 figure 6.6
-when a reaction is performed inside a bomb
calorimeter, the following equation is used:
qcal = (Ccal)(∆T)
qcal = heat absorbed by calorimeter
Ccal = heat capacity of entire calorimeter,
usually given
**the amount of heat gained by the calorimeter
must equal the amount of heat released by the
reaction
-equal in magnitude, but opposite in sign
qcal = -qrxn
-since the reaction inside a calorimeter is
occurring under conditions of constant volume,
then:
qrxn = ∆Erxn
page 252
-most chemical reactions and physical changes
occur at constant pressure in a lab
enthalpy (H)- heat content of a system at
constant pressure
H = ∆H = q
-enthalpy change for chemical reaction = ΔHrxn
-called heat of reaction
∆H is + for endothermic reaction
∆H is – for exothermic reaction
Examples:
Identify each as endo or exothermic and tell the
sign on ∆H.
1) sweat evaporating from skin
endothermic, +∆H
2) water freezing in a freezer
exothermic, - ∆H
3) wood burning in a fire
exothermic, - ∆H
Try these!!
1) an ice cube melting
endothermic, +∆H
2) nail polish remover quickly evaporating after
it is accidently spilled on the skin
endothermic, +∆H
3) a chemical hand warmer emitting heat after
the mixing of substances within a small
handheld package
exothermic, -∆H