Honors Chemistry
Unit 8 Thermochemistry
Thermochemistry is the study of the changes in energy that
accompany chemical reactions and physical changes.
(Ch. 16)
A. What is energy- energy is the capacity to do work or
produce heat. Think of energy as something that an object
possesses. Think of heat and work as ways that objects
exchange energy. It exists in two basic forms.
1. Chemical Potential Energy- energy due to the
composition of the substance and the type of bonding
between the atoms. Stronger the bonds, than the more
potential energy.
2. Kinetic Energy- energy of motion of the molecules.
Equation: K.E. = ½ mv2. Greater the velocity, than the
greater the K.E. Also, as you already know: 𝐾. 𝐸. ≈
𝑡𝑒𝑚𝑝.
3. All chemical and physical processes are governed by the
Law of Conservation of Energy (1st Law of
Thermodynamics). Which states that in any chemical
reaction or physical process, energy can be converted
from one form to another, but it is neither created nor
destroyed.
B. All chemical reactions involve changes in energy.
1. When we analyze energy changes, we need to focus on
a well-defined part of the universe.
a. System- the part that is singled out for study. In a
reaction usually the reactants and products.
b. Surroundings- everything with which the system can
exchange energy.
2. Breaking of bonds absorbs energy; endothermic process
a. Endothermic reactions absorb energy from their
surroundings and transfer it to the system.
b. The change (∆) in energy is positive.
3. Formation of bonds releases energy; exothermic process
a. Exothermic reactions release energy from the system
to the surroundings.
b. The change (∆) in energy is negative.
4.
Change in energy called “change in enthalpy”— ∆𝐻
a. when energy required to break bonds > energy
released to form new bonds, +∆𝐻 (endothermic)
1. products at a higher energy state than reactants
(weaker bonds)
2. surroundings lose energy (cool down)
b. when energy required to break bonds < energy
released to form new bonds, –∆𝐻 (exothermic)
1. products at a lower energy state than
reactants (stronger bonds)
2. surroundings gain energy (heat up)
1
∆H°rxn= Hproducts−Hreactants
c.
5.
C.
Thermochemical equation
a. Chemical equation with ∆𝐻
1. listed to the right of equation
2. included as reactant (endothermic) or product
(exothermic)
b. ∆𝐻 can be used in dimensional analysis process. The
magnitude is based upon the stoichiometric amounts
of reactants and products as written.
c. Units: SI unit of energy is the joule (J)
1. calorie: 1 cal = 4.184 J
2. Calorie, also called a nutritional calorie:
1 kilocal = 1 Cal
3. 1000 J = 1kJ
Thermochemical changes of state- many processes besides
chemical reactions absorb or release heat.
1. Phase changes for water that are endothermic.
a. Energy absorbed from surroundings to break
hydrogen bonds to go from a solid to a liquid is called
heat of fusion, Hfusion a.k.a melting
 H20(s) → H20(l)
∆Hfusion= 6.01 kJ/mol
b.
2.
Energy absorbed from surroundings to break
hydrogen bonds to go from a liquid to a gas is called
heat of vaporization Hvaporization a.k.a evaporation.
 H20(l) → H20(g) ∆Hvaporization= 40.7 kJ/mol
Phase changes that are exothermic.
a. Energy released into surroundings to form hydrogen
bonds when a gas changes into a liquid is called heat
of condensation Hcondensation
 H20(g) → H20(l) ∆Hcondensation = −40.7 kJ/mol
b.
Energy released into the surroundings to form
hydrogen bonds to go from a liquid to a solid is called
heat of solidification Hsolidification a.k.a freezing
 H20(l) → H20(s) ∆Hsolidification= −6.01 kJ/mol
D. Calorimetry- measurement of heat flow, the transfer of
heat that occurs in chemical and physical processes can be
measured using a calorimeter. Reactants are put in an
insulated container filled with water, where heat is
exchanged between reactants and water, but no heat is
lost.
1. Specific Heat Capacity (Cs)- quantity of heat required to
change the temperature of one gram of a substance by 1
K or (1°C). Units:
𝑱
𝒈∙𝑲
2
𝒒 = 𝑪𝒔 × 𝒎 × ∆𝑻
 Cs= specifc heat capacity of substance;
H2O = 4.18 J/g•K
 q= heat flow
 m= mass in grams
 ∆𝑇 = 𝑇𝑓 − 𝑇𝑖
3. Types of calorimetry
a. constant pressure- used for processes that take
place at constant pressure: coffee-cup calorimeter.
 assuming no heat escapes, heat released by
the system is absorbed by the water.
𝑞𝑠𝑦𝑠𝑡𝑒𝑚 = −𝑞𝑤𝑎𝑡𝑒𝑟
2.

E.
𝒒 = 𝑪𝒔 × 𝒎 × ∆𝑻
Hess’s Law - it states that if you can add two or more
thermochemical equations to produce a final equation,
then the sum of the enthalpy changes of each individual
reaction is the enthalpy change for the final reaction.
- Rule #1 If the reaction is reversed, the sign of the ∆H
changes too.
-
Rule #2 If the coefficients are multiplied by a factor,
then the ∆H is multiplied by the same factor.
-
Rule #3 The state of matter (s, l, g, aq) of reactants and
products, affects whether they cancel out or not.
∆𝐻𝑜𝑣𝑒𝑟𝑎𝑙𝑙 = ∆𝐻1 + ∆𝐻2 +. . .
F.
Standard Enthalpy of Formation (∆H°f)- ∆H values for a
reaction in which a compound is formed from its
constituent elements in their standard states, 298K and
1atm.
∆H°rxn = ∑n∆H°f
(products)
− ∑n∆H°f
(reactants)
G. ∆H using bond energy (BE) data
Bond Energies in (kJ/mol)
Single
H C N O S F Cl Br
H 436 413 391 463 339 567 431 366
C
348 293 358 259 485 328 276
N
163 201
272 200 243
O
146
190 203
S
266 327 253 218
F
155 253 237
Cl
242 218
Br
193
I
I
299
240
243
208
175
151
Multiple
C=C 614
C=N 615
C=O 799
N=N 418
N=O 607
O=O 495
O=S 523
S=S 418
CC 839
CN 891
CO 1072
NN 941
3
1.
energy needed to break a bond (i.e. C–H) in a diatomic,
gaseous molecule, which contains the bond type
 is approximately the same for any molecule

gaseous species
 positive value (+ BE) for breaking bonds
 forming bonds (– BE)
∆H = BEbreaking – BEforming
threshold temperature (Tthreshold) when ∆G = 0
Tthreshold = ∆Ho/∆So
d. summary chart for determining when ∆G < 0
-∆H
+∆H
+∆S
∆G < 0 for all T
∆G < 0 when T > ∆Ho/∆So
-∆S ∆G < 0 when T < ∆Ho/∆So
∆G > 0 for all T
c.
H. Second Law of Thermodynamics: predicting spontaneous
change.
1. Spontaneous change is one that occurs without ongoing
outside intervention. The sign of ∆H by itself does not
predict spontaneity.
a. Spontaneous processes with ∆H < 0
3
 2Fe + O2⟶ Fe2O3(s)
∆𝐻°𝑟𝑥𝑛 = −826 𝑘𝐽
2
b.
2.
Spontaneity can be predicted by a function called
Entropy (S). Entropy is a measure of the disorder or
randomness of particles in a system also a measure of
energy dispersal in a system.
a. 2nd Law of Thermodynamics states that for any
spontaneous process, the entropy of the universe
increases. ∆𝑆𝑢𝑛𝑖𝑣 > 0
b. Entropy determines the direction of chemical and
physical change. A chemical system always proceeds
in a direction that increases the entropy of the
universe.
3.
Entropy increases (∆𝑆𝑢𝑛𝑖𝑣 > 0) for the following:
a. phase change: solid ⟶ liquid ⟶ gas
b. diffusion, effusion
c. dissolving of solids
d. increase of temperature
e. increase in number of moles of a gas during a rxn.
f. atomic size within a group; increases entropy.
g. molecular complexity within the same state; CH4(g),
C2H6(g), C3H8(g)
4.
I.
°
 CH4 + 2O2 ⟶ CO2 + 2H2O ∆𝐻𝑟𝑥𝑛
= −802 𝑘𝐽
°
 H2O(l) ⟶ H2O(s)
∆𝐻𝑟𝑥𝑛
= −6.02 𝑘𝐽
Spontaneous processes with ∆H > 0
°
 H2O(l) ⟶ H2O(g)
∆𝐻𝑟𝑥𝑛
= 40.7 𝑘𝐽
°
 H2O(s) ⟶ H2O(l)
∆𝐻𝑟𝑥𝑛 = 6.02 𝑘𝐽
°
 NaCl(s) ⟶ Na+(aq) + Cl-(aq) ∆𝐻𝑠𝑜𝑙𝑛
= 3.9 𝑘𝐽
∆S°rxn = ∑n∆S°
(products)
− ∑n∆S°(reactants)
Gibbs Free Energy (G): Overall Energy State (19.5 to 19.6)
1. spontaneous chemical and physical changes occur
because the products are at a lower energy state
and/or more disordered
a. Gibbs free energy accounts for both ∆H and ∆S
b. change in free energy: ∆G= ∆Ho – T∆So
2. determine if thermodynamically favorable (∆G < 0)
a. lower potential energy (-∆H)—chemical reactions
b. greater disorder (+∆S)—physical changes
4