University of Science and Technology of Hanoi
Department of Fundamental and Applied Sciences
GENERAL CHEMISTRY II
Dr. Nguyen Duc Anh
nguyen-duc.anh@usth.edu.vn
0904888819
GC1 vs GC 2
CH4 + Cl2
CH3Cl + HCl
GC 1
- Structure of atoms and
composition of
molecules
- Types of Bonds
- Geometry of molecules
Missing questions
- How does the reaction proceed?
- What is the relation of the reaction and
the environment?
- Does the reaction provide heat or consume
heat?
- How fast is the reaction ?
- Does the reaction happen naturally?
- Why does the reaction go forward but
2
GC 2
24 hours (8 sections of 3
1. Thermochemistry
hours each )
2. Thermodynamics of
reactions
3. Kinetics of reactions
4. Chemical equilibrium
5. Acids and Bases
6. Electrochemistry
Dr. Nguyen Duc Anh
Department of Fundamental and Applied
Sciences
Phone: 0904888819
Email: nguyen-duc.anh@usth.edu.vn
Dr. Le Phuong Thu
Department of General Education
Phone: 0392411761
Email: lephuong.thu@usth.edu.vn
For administration:
ged_dept@usth.edu.vn
3
Rules
Total score’s constitution:
Attendance + Contribution (10%)
Midterm (30%); (3 first chapters)
Final Exam (60%)
* No Personal Stuffs (Music, Facebook, etc) (perhaps Google, ChatGPT, Deepseek… when your
teacher ask you too many questions)
** Actively involve in discussion/exchange
*** Attendance checking: Random ~50 students in class; presence minimum 5/8 sections
(if you absent 4 sections  you then redo this class ... next year)
* Teaching:
Important concepts will be provided/explained
Lecturers instruct  Students work with text-book, text-book, and text-book.
4
Textbook
5
Calm down and have
with GC
2
because GC 2
is
6
University of Science and Technology of Hanoi
Department of Fundamental and Applied Sciences
GENERAL CHEMISTRY II
THERMOCHEMISTRY
ENERGY FLOW AND CHEMICAL CHANGE
Dr. Nguyen Duc Anh
nguyen-duc.anh@usth.edu.vn
0904888819
7
Content
1. Forms of Energy and Their Interconversion
2. Enthalpy: Chemical Change at Constant Pressure
3. Calorimetry: Measuring the Heat of a Chemical or
Physical Change
4. Stoichiometry of Thermochemical Equations
5. Hess's Law: Finding ∆H of Any Reaction
6. Standard Enthalpies of Reaction (∆Horxn)
8
1. Forms of Energy and Their
Interconversion
All changes in matter are accompanied by changes of
energy
+ Thermodynamics is the study of energy and its transformations.
+ Thermochemistry is a branch of thermodynamics that deals with
the heat involved in chemical and physical changes.
9
9
1. Forms of Energy and Their
Interconversion
Nature of energy
Energy: The capacity to supply heat or
• Potential Energy (EP)
The stored energy
(due to position or
composition)
• Kinetic Energy (EK)
The energy of motion
10
1. Forms of Energy and Their
Interconversion
Nature of energy
• When energy is transferred from one object to another, it
appears as work and heat.
• Work (w): Energy transferred when
an object is moved by force
• Heat (q): Energy transferred
between objects due to a
temperature difference
11
11
1. Forms of Energy and Their
Interconversion
Unit of energy
• The SI unit of energy is the joule (J)
1 J = 1 kg m2 s-2
Electronvolt
• The calorie (cal) was once
defined as the quantity of
energy needed to raise the
temperature of 1 g of water
by 1 oC.
• The1 British
Thermal
Unit
cal = 4.184
J
(BTU) was is often used to
rate appliances.
1 BTU ≈ 1055 J = 0,293 kW
Joule
BTU
Calorie
Dyne
Watt-hour
Hartree
Newton-meters
12
1. Forms of Energy and Their
Interconversion
You can find the unit of energy her
13
13
1. Forms of Energy and Their
Interconversion
The system and its surrounding
• In thermodynamic, the universe
is divided into two parts:
+ The system: The part that we are
focusing on.
+ The surroundings: Everything else
14
1. Forms of Energy and Their
Interconversion
Internal Energy (E) and Energy Changes
+ Internal Energy of a system is
the sum of kinetic and potential
energy of all “particles” contained
in it.
+
Energy
Change
(ΔE):
Energy
transfer to or from a System.
- A change in the energy of the system
must be accompanied by an equal and
opposite change in the energy of the
surroundings.
15
1. Forms of Energy and Their
Interconversion
The Law of Energy Conservation
The first law of thermodynamic: the total energy of the
universe is constant.
• A chemical reaction always involves a transfer of energy between
the system and the surroundings.
• Energy is conserved and is neither created or destroyed.
• Energy is transferred in the form of heat (q) and/or work (w).
• Because energy transfer can only occur in these two forms, the
total change in the system’s internal energy (DE) is therefore:
16
1. Forms of Energy and Their
Interconversion
Heat and work
• The sign of q and w can be positive or negative depending
on the change the system undergoes.
• If the system provide heat to the surroundings, q is
negative.
• If the system done work on the surrounding, w is
negative.
17
1. Forms of Energy and Their
Interconversion
Energy transferred as heat only
1
2
1. Heat flowing out from a system to the
surroundings
System releases energy as heat
⟶ q is negative
⟶ Efinal < Einitial
⟶ ΔE is negative
2. Heat flowing into a system from
the surroundings
System absorbs energy as heat
⟶ q is positive
⟶ Efinal > Einitial
⟶ ΔE is positive
18
1. Forms of Energy and Their
Interconversion
Energy transferred as work only
Pressure – Volume
work
(PV work)
1. Work done by a system on the
surroundings
System releases energy as work done by it
⟶ w is negative
⟶ Efinal < Einitial
⟶ ΔE is negative
2. Work done on a system by the
surroundings
System absorbs energy as work done on it
⟶ w is positive
⟶ Efinal > Einitial
⟶ ΔE is positive
19
1. Forms of Energy and Their
Interconversion
Sign convention for q, w and DE
• For q: “+” means system gains (absorbs) heat; “-” means
system releases heat
• For w: “+” means work done on system; “-” means work done by
20
1. Forms of Energy and Their Interconversion
Forms of works
• The most important (common) forms of work are:
+ PV work = the expansion (or contraction) of a gas (e.g. in engines)
+ Electrical work = movement of charged particles (e.g. In battery... )
21
1. Forms of Energy and Their Interconversion
PV work – Expansion work
• Pressure – Volume (PV) work: Work done as
the result of a volume change in the system
Expansion work (volume increase)
1. The volume of a gaseous system will increase
if the temperature is raised (Charles’s law)
2. Or a chemical reaction results in a net
increase in
the number of moles of gas (Avogadro’s law)
22
1. Forms of Energy and Their
Interconversion
PV work – Contraction work (volume decrease)
A gaseous system will decrease
in volume if the temperature
is lowered
Or a reaction results in a net
decrease in the number of
23
23
1. Forms of Energy and Their
Interconversion
Energy Changes – State Functions
• The state of a chemical system is
characterized by a set of
parameters including T
(temperature), P (pressure) and V
P2 V2
P(volume).
V
1
T1
1
DE1
= DE2
T2
• State function: A function of
properties whose value depends
only on the present state, or
condition, of the system, not on
the path used to arrive at that
state.
25
1. Forms of Energy and Their
Interconversion
Energy Changes – State Functions
ΔE does not depend on how the change takes place, but only on
the difference between the final and initial states
- q and w are not state functions because their values do depend on the path t
system takes.
- ΔE, ΔT, ΔP, and ΔV—depend only on the initial and final states
 they are
26
1. Forms of Energy and Their
Interconversion
Applying in Thermodynamic Processes
• Isobaric processes (constant
pressure): P1 = P2
qP = ΔE + P.ΔV
• Isochoric process (constant
volume): V1 = V2 or ΔV = 0
qV = ΔE
• Isothermal process (constant temperature):
ΔE = 0
q = -w
• Adiabatic process (no heat transferred): q = 0
ΔE = w
27
2. Enthalpy: Chemical change at
constant Pressure
Chemical reactions are often performed in vessels opened to the atmosphere
In constant pressure,
• Enthalpy is the energy in a thermodynamic
system that equals the internal energy
plus the product of volume and pressure.
• The
change
heat
in
absorbed
enthalpy
or
(ΔH)
released
equals
at
the
constant
pressure
• Enthalpy is a state function
DH = Hfinal – Hinitial = Hproducts
– Hreactants
28
2. Enthalpy: Chemical change at
constant P
DH vs DE
• ΔH describes the heat associated (absorbed
or released)
in a chemical reaction at
DH ≈ DE
pressure.
• Since
many
chemical
reactions
involve
little (if any) PV work, most of energy
2. Reactions in which the amount
1. Reactions that do not involve
gases.
change
is due
to4(aq)
a transfer
2KOH(aq)
+ H2SO
⟶ K2SO4(aq)of
+ heat.
2H2O(l)
(mol) of gas does not change.
N2(g) + O2(g) ⟶ 2NO(g)
3. Reactions in which the amount (mol) of gas
does change but qP is usually much larger than
ΔH = -483.6
2H2(g) + OP.ΔV
2(g) ⟶ 2H2O(g)
kJ
PΔV = -2.5
29
2. Enthalpy: Chemical change at
constant P
Exothermic vs Endothermic reactions
ΔH < 0
ΔH > 0
ΔH < 0 : heat is released to ΔH > 0 : heat is absorbed from
the surrounding
Exothermic
the surrounding
Endothermic
30
2. Enthalpy: Chemical change at
constant P
Enthalpy associated with physical changes
ΔHsub = - ΔHdep
ΔHvap = - ΔHcon
ΔHmelt = ΔHfree
ΔHsub = ΔHmelt +
ΔHvap
31
2. Enthalpy: Chemical change at
constant P
Problem
In each of the following cases, determine the sign of ΔH, state
whether the reaction is exothermic or endothermic, and draw an
enthalpy diagram:
(a)H2(g) + ½ O2(g) ⟶ H2O(l) + 285.8 kJ
(b)40.7 kJ + H2O(l) ⟶ H2O(g)
32
3. Calorimetry: Measuring the Heat
• The
heat
associated
with
a
chemical
reaction can be determined experimentally
using a device called calorimeter
• Calorimetry,
heat,
is
the
science
based
on
of
measuring
observing
the
temperature change when a body absorbed or
released energy as heat
• Substances response differently to being
Florentine thermometers, 16
heated
• When
a
substance
is
heated,
the
energy
33
3. Calorimetry: Measuring the Heat
Temperature scales
34
3. Calorimetry: Measuring the Heat
Heat capacity
35
3. Calorimetry: Measuring the Heat
Calorimetry and Heat
•capacity
Heat capacity (C): the quantity of heat required to change temperature of
an object or substance by 1 K (J/K)
(or 1 oC (J/oC))
• Specific heat capacity (c): the quantity of heat required to change the
temperature of 1 gram of a substance or material by 1 K (J/g.K)
(or 1 oC
(J/g.oC))
• Molar heat capacity (Cm): the quantity of heat required to change the
temperature of 1 mole of a substance by 1 K (J/mol.K)
(or 1 oC
(J/mol.oC))
36
3. Calorimetry: Measuring the Heat
Problem
A layer of copper welded to the bottom of a skillet
weighs 125 g. How much heat is needed to raise the
temperature of the copper layer from 25°C to 300°C? The
specific heat capacity (c) of Cu is given in Table 6.2.
q = ???
37
3. Calorimetry: Measuring the Heat
Calorimetry
Typically, we can determine the magnitude of the heat flow
by measuring the magnitude of the temperature
change the heat flow produces. The measurement of heat
flow is calorimetry; a device used to measure heat flow is
a
calorimeter.
38
3. Calorimetry: Measuring the Heat
Constant pressure
calorimetry
• Direct measuring DH
• Low precision (neglect the
absorption of calorimeter
components)
• Used for most reactions
Constant volume
calorimetry (Bomb
calorimetry)
• Direct measuring DE
• Higher precision
• Used for combustion reactions
40
4. Stoichiometry of thermochemical
equations
• A thermochemical equation
that includes DHrxn
is
balanced
equation
• The sign of DH indicates whether the reaction is
exothermic or endothermic
Molar ratio
from balance
equation
• The magnitude of DH is proportional to the amount
of substances
DH
(kJ/mol)
• The
value
of DH can beequivalent
used into 1/2
a calculation
in
 286
kJ is thermochemically
mol of O2(g)
 286
kJ isway
thermochemically
the
same
as a moleequivalent
ratio. to 1 mol of H2O(l)
How much energy is thermochemical equivalent to 90
grams of water?
42
4. Stoichiometry of thermochemical
equations
Problem: The major source of aluminum in the world is bauxite (mostly aluminum
oxide). Its thermal decomposition can be written as
Al2O3(s) ⟶ 2Al(s) + 3/2 O2(g)
ΔH =
1676 kJ
If aluminum is produced this way, how many grams of aluminum can form when
1.000×103 kJ of heat is transferred?
43
43
5. Hess’ Law: Finding DH of Any
reaction
Enthalpy
is
a
state
function,
the
enthalpy
change in going from initial
state to final state is
independent
of
the
pathway
“In
going
from
a
particular
set
of
reactants to a particular
set
of
products,
the
change in enthalpy is the
whether the reaction takes
place in one step or a
series of steps”
Path (C)
Reactant
s
Path (B)
ΔH
Path
(A)
Products
44
44
5. Hess’ Law: Finding DH of Any
reaction
The overall enthalpy change for a reaction is equal to the
sum of the enthalpy changes for the individual steps in
the reaction.
DH = DH1 + DH2 +
DH3
45
5. Hess’ Law: Finding DH of Any
reaction
Rules for Hess’s Law
• If reaction is reversed, the sign of DH is also
reversed
• If coefficient of reaction is multiplied by an integer, the DH is also
multiplied by the same value
46
5.Hess’ Law: Finding DH of Any
reaction
Using Hess’s Law to calculate an Unknown
DH
Known that:
Reaction
coefficie
nts
DH =
?
Analysis of overall reaction:
1. Reactants (left side) and products
2. Reaction coefficients
(right side)
 Adjust the “component” reactions
 Adjust the direction of
reactions and theirDH
“component”
coefficients and their DH
47
5. Hess’ Law: Finding DH of Any
reaction
Using Hess’s Law to calculate an Unknown
DH
• Step 1: Adjust “component” reactions direction and
their DH
DH3’ = - DH3 = -206.1 kJ
DH =
?
48
5. Hess’ Law: Finding DH of Any
reaction
Using Hess’s Law to calculate an Unknown
DH
• Step 2: Adjust “component” reactions coefficients
and their DH
DH1’ = 2 x DH1 = 2 x 131.3 kJ
2
1
1x
x
1 DH3’ = - DH3 = -206.1 kJ
1x
x
x
DH =
?
49
5. Hess’ Law: Finding DH of Any
reaction
Using Hess’s Law to calculate an Unknown
DH
• Step 3: Remove the same substances from 2 sides and sum up
reactions
2
DH1’ = 2 x DH1 = 262.6 kJ
2
2
2
1
x
1 DH3’ = - DH3 = -206.1 kJ
x
DH = DH1’ + DH2 + DH3’
= 262.6 + (-41.2) + (206.1) = 15.3 kJ
50
5. Hess’ Law: Finding DH of Any
reaction
51
6. Standard Enthalpy of reactions
(DHorxn)
• Thermodynamic variables, such as DH, vary somewhat with conditions
• A set of specific conditions called standard states is established in order
to study and compare reactions:
Gaseous
substances
1 atm and ideal
behavior
Substance in
solutions
Concentration of 1
M
Pure substances
The most stable form
at
1 atm and the
temperature of
interest (25 oC –
298 K)
• When the enthalpy change of a chemical reaction is measured at the standard
state, it is the standard enthalpy of reactions, DHorxn (so called standard
heat of reaction).
52
6. Standard Enthalpy of reactions
(DHorxn)
Standard Enthalpy of formation
(DHof)
Substance
Substance
• The enthalpy change for the
formation
of
1
mol
of
a
substance in its standard state
from its constituent elements in
their standard states
Standard states
1 mol
• Note: DHof (elements) = 0
53
6. Standard Enthalpy of reactions
(DHorxn)
Determining ΔH°rxn from DHof Values for Reactants and
Products
aA+bB
cC+
dD
DHorxn = (d.DHof, D + c
DHof, C)
(a. DHof, A + b. DHof, B)
The standard enthalpy of reaction is the sum of the standard enthalpies of formation of the
products minus the sum of the standard enthalpies of formation of the reactants
54
54
6. Standard Enthalpy of reactions
(DHorxn)
Problem: Nitrogen dioxide reacts with oxygen to form dinitrogen
pentoxide:
Substance
4NO2(g) + O2(g) ⟶ 2N2O5(s)
Calculate ΔH° rxn from ΔH°f values
55
55
6. Standard Enthalpy of reactions
(DHorxn)
Bond dissociation energies
• Bond dissociation energy (D) is the amount of energy that must be supplied to break
a chemical bond in an isolated molecule in the gaseous state.
• Bond dissociation enthalpy and bond formation enthalpy has the same value but
opposite charges
56
56
6. Standard Enthalpy of reactions
(DHorxn)
Calculation of DHorxn from D
Problem: Find an approximate ∆H° in kilojoules for the industrial synthesis of chloroform by
reaction of methane with Cl2.
57
6. Standard Enthalpy of reactions
(DHorxn)
Heats of combustion
Combustion: A chemical in which a substance
reacts with oxygen gas, releasing energy in
the form of light and heat. Standard enthalpy
of combustion is the heat associated with 1
mol of substance during combustion.
DHcom = -2512.4 kJ
DHorxn = ΣmDHocom, reactants - ΣnDHocom,
• Calculation of DHorxn from DHocom
products
aA+bB
cC+
D b. DHocom, B) - (d.DHocom, D
DHorxn = (a. DHocom, Ad +
+ c DHocom, C)
58
6. Standard Enthalpy of reactions
(DHorxn)
59
59
Key points
• Chemical or physical change is always accompanied by a change in the energy
content of the matter.
• All energy changes occur as heat (q) and/or work (w), transferred either from the
surroundings to the system or from the system to the surroundings.
• The total energy of the universe is constant (law of conservation of energy, or
first law of thermodynamics).
• State function—a property that depends only on the current state of the system
• The change in enthalpy (DH) equals the heat transferred at constant pressure, qP.
Enthalpy (H) is also a state function.
• DHrxn is negative (< 0) if the reaction releases heat (exothermic) and positive (>
0) if it absorbs heat (endothermic).
• Each material has its own capacity for absorbing heat.
• The DH is related stoichiometrically to the amounts (mol) of reactants and
products.
60