part-01: Thermodynamics

Thermodynamics - Review
❑ Thermodynamics is the science of ENERGY
 Properties of matter in terms of energy (micro- or macroscopic)
 The energy exchange in a process (quantity, direction and limit)

Energy is usually defined as the capacity to do work.

Work is directed energy change resulting from a process.

Radiant energy, or solar energy, comes from the sun and is Earth’s
primary energy source.

Thermal energy is the energy associated with the random motion of

atoms and molecules.
Chemical energy is stored within the structural units of chemical
substances; its quantity is determined by the type and arrangement of
constituent atoms.

Potential energy is energy available by virtue of an object’s position.

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part-01: Thermodynamics
Thermodynamics - Review
❑ Object of study in thermodynamics - SYSTEM
 System - a quantity of matter or a region in space chosen
➢ Closed system - no mass crosses system boundary; but energy can be
exchanged with surroundings.
➢ Open system - there is mass or energy exchange with its surroundings
 System contains MATTER, has a BOUNDARY and is always in some STATE
a) Open system. The beaker of hot
coffee transfers energy to the
surroundings; it loses heat as it cools.
Matter is also transferred in the form of
water vapor.

(b) Closed system. The flask of hot

coffee transfers energy (heat) to the
surroundings as it cools. Because the
flask is stoppered, no water vapor
escapes and no matter is transferred.

(c) Isolated system. Hot coffee in an

insulated container approximates an
isolated system. 2
part-01: Thermodynamics
Fundamental Laws of Thermodynamics
❑ First Law of Thermodynamics - ENERGY CONSERVATION
One of many expressions:
Energy cannot be created or destroyed. In a process energy changes
from one form to another form and the total amount of energy
remains constant.
❑ Second Law of Thermodynamics - DIRECTION of change in a process
One of the several expressions states:
Processes occur in a direction of decreasing quality of energy.

❑ Third & Zeroth Laws of Thermodynamics -Temperature definition &

measurement
Third Law - Entropy definition
The entropy of a pure crystalline substance at absolute zero temperature is
zero.
Zeroth Law - Temperature measurement
If two bodies are in thermal equilibrium with a third body, they are also in
thermal equilibrium with each other.
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part-01: Thermodynamics
❑ First Law of Thermodynamics - ENERGY CONSERVATION
Energy cannot be created or destroyed. In a process energy changes from
one form to another form and the total amount of energy remains
constant.

= energy content of 1 mol SO2(g) - energy

content of [1 mol S(s) + 1 mol O2(g)]

Kerja (w) dan Panas (q)

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part-01: Thermodynamics

See example 7-5 (pp 254;

General Chemistry-Petrucci)

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part-01: Thermodynamics
Kerja ekspansi V>0
Kerja kompresi V<0

llustrating work (expansion) during the

chemical reaction:

The oxygen gas that is formed pushes

See example 6.1 (pp back the weight and, in doing so, does
183 Essential work on the surroundings.
Chemistry R change)

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part-01: Thermodynamics
Properties of System
❑ Properties of system - characteristics of system
 Intensive properties - independent of the system size
➢ e.g. Temperature, pressure, density etc. (value does not change
by the division)

 Extensive properties - dependent on the system size

➢ e.g. mass, volume, internal energy etc. (value changes by the
division)

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part-01: Thermodynamics
Basic System Properties [state of system] [1]
❑ State of function: properties that are determined by the state of
the system, regardless of how that condition was achieved.
❑ The magnitude of change in any state function depends only on the
initial and final states of the system and not on how the change is
accomplished

❑ Temperature, T
 T is the measurement of the hotness of a substance/system
 Several scales are in use: oC, F, K and R (less common).
 Is state function

❑ Pressure, P
 P is the force exerts by a fluid to a system per unit area
 absolute pressure and gage pressure
 Several units are commonly used (Pa, bar, mmHg, psi, atm.)
 Is a state function

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part-01: Thermodynamics
Basic System Properties [state of system] [2]
❑Volume, V
 The volume of a system [is a state function]

❑Energy: potential energy (U), entahalpy (H),

Entropy (S), Energy Bebas Gibss (G) are
also state functions

❑Entropy, S
S is a measure of molecular disorder, or
molecular randomness
➢S has an energy unit but is NOT a form of energy;
➢ important: S is not conserved in a process.

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part-01: Thermodynamics
Substance in a System
❑ Substance
 Pure substance - has a fixed and uniform chemical composition.
 Mixture of pure substances
➔ properties of a mixture depend on the properties of each individual component /
constituent and the amount of each in the mixture

❑ Phases of substances
 solid molecules are arranged in 3-D pattern that is repeated throughout. The
attractive forces between molecules are large and keep the molecules in fixed
positions
liquid The molecular spacing is in the same order as in a solid and molecules
remain ordered structure; but the molecules’ position is not fixed in 3-D structure
gas Molecules are far apart; they move freely and collide each other
 The energies contain in the various phase following the order: gas > liquid > solid.

❑ Property-relation diagrams
 Type of diagrams commonly used to describe property-relations of a
substance
T-V, P-V, P-T, P-V-T, T-S and H-S diagrams.
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part-01: Thermodynamics
Chemical Reaction - The Equation
❑ Reaction equation for A reacting with B forming C and D
A + B → C + D ; general expression of a reaction
vAA + vBB = vCC + vDD ; quantitative representation of a reaction
vAA + vBB D vCC + vDD ; representing an equilibrium controlled
reaction

❑ Reaction stoichiometry
 vA, vB, vC and vD are numbers, called stoichiometry coefficients
 The concept of mole (the number of molecules) in a chemical
reaction
 Determination of stoichiometry coefficients - balancing equation
(equal number of each atom on both sides of the equation)
e.g. 2NO + O2 D 2NO2 we have: 2 N, 4 O on both sides

Q. For the same reaction can we write the equation as

4 NO + 2 O2 D 4NO2 or NO + ½O2 D NO2 ?

…….how is stoichiometric?

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part-01: Thermodynamics
❑ Chemical reactions
For most gas-phase reactions occurring at
elevate T. - the gases can be treated as
ideal gases
Many reactions proceed continuously at a
constant P - can be treated as steady-flow
process
Many reactions are carried out at a
constant T (very common) - these are
isothermal process
Many batch reactions are carried out at
constant volume (T and P may vary)
The reactions take place in a well-insulated
chamber is usually treated as an adiabatic
process.

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part-01: Thermodynamics
Review questions
1. Define these terms: system, surroundings, open system,
closed system, isolated system, thermal energy, chemical
energy, potential energy, kinetic energy, law of
conservation of energy.

2. Explain what is meant by a state function. Give two

examples of quantities that are state functions and two
that are not.

3. A gas expands in volume from 26.7 mL to 89.3 mL at

constant temperature. Calculate the work done (in
joules) if the gas expands (a) against a vacuum, (b)
against a constant pressure of 1.5 atm, and (c) against a
constant pressure of 2.8 atm.
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part-01: Thermodynamics
The Second Law of Thermodynamics - DIRECTION of
change in a process
 Processes occur in a direction of decreasing quality of energy.
 The entropy of the universe increases in a spontaneous process and
remains unchanged in an equilibrium process.

a) An exothermic process transfers heat from the system to the surroundings and results in
an increase in the entropy of the surroundings. (b) An endothermic process absorbs heat
from the surroundings and thereby decreases the entropy of the surroundings. 14
part-01: Thermodynamics
Direction of a Reaction; Spontaneous or not..?

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part-01: Thermodynamics
Thermodynamics of Chemical Reaction Systems
❑ Energies transferred during a process
Sensible energy - P, T
} non-chemical
energy
Latent internal energy - energy associated with phase change
Chemical internal energy - energy associated with destruction and
formation of a chemical bond of molecules

❑ Thermodynamics of chemical reaction systems

Study the energy transfer process in chemical reactions
Predict the direction of a chemical reaction and the energy transfer
associated with the reaction
Predict the final state (equilibrium) of a chemical reaction system,
and
Study factors that affect the characteristics of the final state ( not the
process!) of a reaction system

R. Chang, Chapter 18; pp 630

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part-01: Thermodynamics
Enthalpy of Reaction System[1]
❑ Enthalpy definition
  gas phase (ideal gas) at
  H = U +  nRT  const. T & P

H=U+PV  dH=dU+PdV + VdP  g-l, g-s or g-s,l mixture at
H = U
 const. T & P
liquid or solid (dv=0)

Enthalpy calculation
➢Standard enthalpy of formation, Hof
Assign Hof of all stable substances (O2, N2, CO2 H2O etc.) at 298
K & 1 atm as 0 (not at 0 K as for S) (Ho298 values of common
substances are available in most chemistry/Chem. Eng.
Handbooks.)
➢Enthalpy of combustion, Hc -often used for combustion
process, similar to enthalpy of formation. Enthalpy change
when 1 mole of substance combusted completely (reacted
completely with O2).
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part-01: Thermodynamics
Enthalpy of Reaction System[2]
➢H is a direct measure of the reaction heat associated with a reaction (if only
chemical energy change exists). When H < 0 - the reaction is exothermic and
when H > 0 is the reaction is endothermic.

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part-01: Thermodynamics
Enthalpy Calculation[1]
Enthalpy of reaction, H, as the difference between the enthalpies
of the products and the enthalpies of the reactants:

(a) Melting 1 mole of ice at 0°C (an endothermic process) results in an enthalpy
increase in the system of 6.01 kJ. (b) Burning 1 mole of methane in oxygen gas
(an exothermic process) results in an enthalpy decrease in the system of 890.4 kJ. 19
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R. Chang. Pp 189
part-01: Thermodynamics
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Figure 6.7 A constant-volume bomb calo-rimeter. The calorimeter is

filled with oxygen gas before it is placed in the bucket. The sample is
ignited electrically, and the heat produced by the reaction can be
accurately determined by measuring the temperature increase in the
known amount of surrounding water. 23
part-01: Thermodynamics
Enthalpy Calculation[2]
When methylhydrazine, CH6N2, a rocket fuel, undergoes
combustion the following reaction occurs:
CH6N2(l) + 5 O2(g) --> 2N2(g) + 2CO2(g) + 6H2O(g)
When 4.00 g of CH6N2(l) is combusted in a bomb
calorimeter, the temperature of the calorimeter increases
from 25oC to 39.5oC. In a separate experiment, the heat
capacity of the calorimeter is measured to be 7.794 kJ/oC.
What is the heat of reaction for the combustion of a mole of
CH6N2(l)?

qrxn = - Ccalorimeter T
= - (7.794 kJ/oC) (14.50oC) = -113.0 kJ
Heat of reaction per mole of CH6N2(l)
= (-113.0 kJ)/(0.0868 mol) = - 1.30 x 103 kJ/mol
part-01: Thermodynamics
Enthalpy Calculation[3]: Calorimetry of Foods
Most of the energy our bodies need
comes from the metabolism of
carbohydrates, fats and proteins.
Carbohydrates decompose into glucose,
C6H12O6. Metabolism of glucose produces
CO2 and H2O and energy
C6H12O6(s) + 6O2 (g) --> 6CO2 (g) + 6H2O(l)
(+2803 kJ)

The combustion of tristearin

C57H110O6, a typical fat:

2C57H110O6 + 163O2 (g) --> 114 CO2(g)

+ 110 H2O(l) (+75,520 kJ)
Calculate Calories from amount
of carbohydrate, protein and fats
On average, the metabolism of
proteins produces ~ 4 Cal/g (45g carb x 4 Cal/g carb) +
carbohydrates produces ~ 4Cal/g
(9g protein x 4 Cal/g protein) +
fats produces about ~ 9 Cal/g
(1 Cal = 1000 cal; I Cal = 4.184 kJ) (3g fat x 9 Cal/g fat) = 243 Cal
part-01: Thermodynamics

Petrucci, pp 250 Chapter 7

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part-01: Thermodynamics

Figure 6.8 A constant-pressure calorimeter made of two Styrofoam coffee cups. The
outer cup helps to insulate the reacting mixture from the surroundings. Two solutions
of known volume containing the reactants at the same temperature are carefully
mixed in the calorimeter. The heat produced or absorbed by the reaction can be
determined by measuring the temperature change. 27
part-01: Thermodynamics

FIGURE 7-3
Determining the specific heat of lead Example 7-2 illustrated (a) A 50.0 g
sample of lead is heated to the temperature of boiling water (100oC. (b) A
50.0 g sample of water is added to a thermally insulated beaker, and its
temperature is found to be 22.0oC. (c)The hot lead is dumped into the cold
water, and the temperature of the final lead water mixture is 28.8°C. 28
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part-01: Thermodynamics
Review questions
1) Define these terms: enthalpy, enthalpy of reaction.

2) Under what condition is the heat of a reaction equal to the

enthalpy change of the same reaction?

3) What is the difference between specific heat and heat

capacity? What are the units for these two quantities? Which
is the intensive property and which is the extensive property?

4) What is meant by the standard enthalpy of a reaction?

5) Which of the following standard enthalpy of formation values

is not zero at 25°C? Na(s), Ne(g), CH4(g), S8(s), Hg(l), H(g).

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part-01: Thermodynamics
Entropy
❑ Entropy (S) is often described as a measure of how spread
out or dispersed the energy of a system is among the
different possible ways that system can contain energy.
❑ S is state function, hanya ada keadaan awal dan akhir

W = number of microstates
k = called the Boltzmann constant
(1.38 x 10-23 J/K)
Sf = final S; Si = initial S

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part-01: Thermodynamics
part-01: Thermodynamics
Entropy of System
Processes that lead
to an increase
in entropy of the
system:
(a) melting: Sliquid
>Ssolid;
(b) vaporization:
Svapor >Sliquid;
(c) dissolving: Ssoln
>Ssolute + Ssolvent

R. Chang, Chapter 18; pp 633

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part-01: Thermodynamics
Entropy Change
q is heat and T is temperature in K

increase in entropy:
1. Pure liquids or liquid solutions are
formed from solids.
2. Gases are formed from either solids
or liquids.
3. The number of molecules of gas
increases as a result of a chemical
reaction.
4. The temperature of a substance
increases. (Increased temperature
means an increased number of
accessible energy levels for the
increased molecular motion,
whether it be vibrational motion of
atoms or ions in a solid, or
translational and rotational motion
of molecules in a liquid or gas.) 41
part-01: Thermodynamics
Entropy calculation [1]
❑ Entropy calculation
➢ Basis of S calculation: 3rd Law of thermodynamics: S=0 at absolute
temperature=0
298 Cp
= 
0
➢ Standard entropy of a substance (1 mole, at 1 atm. 25C) as S 298 dT
0 T
(So298 value of common substances are available in most chemistry / chem.
engg handbooks.)

This term becomes zero if there is

no phase change during the reaction
❑ The entropy change in a reaction at T, P
dT Qphase
ST0 = ( vi Si0,T )prod −  ( vi Si0,T )reac in which Si0,T = S298
T
0
+  C p ,i +
298 T Tj

➢ If a reaction is adiabatic, it can only proceed when S>0 (rxn reaches equilibrium when
S=0).
➢ Usually S analysis of a reaction system is complicated and less convenient to use.
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part-01: Thermodynamics
part-01: Thermodynamics
Entropy calculation [2]

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part-01: Thermodynamics
Entropy calculation [3]

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part-01: Thermodynamics
Review questions
❑ Explain what is meant by a spontaneous process. Give two
examples each of spontaneous and non-spontaneous
processes.

❑ Which of the following processes are spontaneous and

which are nonspontaneous?
a) dissolving table salt (NaCl) in hot soup;
b) climbing Mt. Everest;
c) spreading fragrance in a room by removing the cap from a
perfume bottle;
d) separating helium and neon from a mixture of the gases.

❑ Statethe second law of thermodynamics in words and

express it mathematically.

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part-01: Thermodynamics
Gibbs Free Energy
❑ Gibbs free energy (also called Gibbs function) definition
G = U + PV -TS = H - TS  G= H - TS at constant T,P
❑ G° is one of the most important thermodynamic properties for a chemical
reaction system.
 It determines the direction of reaction to proceed

for a reaction GT   0 reaction can proceed in the direction specified


at constant T, GT  = 0 reaction at equilibrium (no further change occurs)
P, we have G   0 reaction will NOT proceed (it can proceed backward!)
 T

➢G°< 0 is the pre-condition which MUST be met for any process (not
limited to chemical reaction systems) to occur (spontaneous process).

➢G°<0 indicates a specified reaction has tendency to proceed; however, it

CANNOT tell how fast that reaction will occur - reaction kinetics tell the rxn
rate.

➢A process/reaction proceeds always in the direction of MINIMISING Gibbs

free energy. This is a very important concept.

➢A process/reaction will stop at G°=0, this is called equilibrium state. 47

part-01: Thermodynamics
Gibbs Free Energy Calculation[1]
❑ Standard Gibbs free energy, Go298
Similar to So and Ho, the standard Gibbs free energy of formation of a substance,
Go, is defined at reference state of T=298 K and P=1 atm
(Go298 values of common substances are available in most chemistry/Chem. Eng
Handbooks.)

❑ Gibbs free energy change in a reaction at T, P

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part-01: Thermodynamics
Gibbs Free Energy Calculation[1]

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part-01: Thermodynamics