4 Lecture 1 Thermo 1

Thermodynamics
Thermodynamics:
→ Describes macroscopic properties of equilibrium systems
→ Entirely Empirical
→ Built on 4 Laws and “simple” mathematics
0th Law ⇒ Defines Temperature (T)
1st Law ⇒ Defines Energy (U)
2nd Law ⇒ Defines Entropy (S)
3rd Law ⇒ Gives Numerical Value to Entropy
These laws are UNIVERSALLY VALID, they cannot be circumvented.
Definitions:
• System: The part of the Universe that we choose to study
• Surroundings: The rest of the Universe
• Boundary: The surface dividing the System from the Surroundings
Systems can be:
• Open: Mass and Energy can transfer between the System and the
Surroundings
• Closed: Energy can transfer between the System and the
Surroundings, but NOT mass
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• Isolated: Neither Mass nor Energy can transfer between the
System and the Surroundings
Describing systems requires:

• A few macroscopic properties: p, T, V, n, m, …
• Knowledge if System is Homogeneous or Heterogeneous
• Knowledge if System is in Equilibrium State
• Knowledge of the number of components
Two classes of Properties:

• Extensive: Depend on the size of the system (n, m, V,…)
• Intensive: Independent of the size of the system (T, p, …)
The State of a System at Equilibrium:

• Defined by the collection of all macroscopic properties that are
described by State variables (p, n, T, V,…) [INDEPENDENT
of the HISTORY of the SYSTEM]
• For a one-component System, all that is required is “n” and 2
variables. All other properties then follow.
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• Notation:
State Functions
• A state function depends only on the initial and final states of a
system.
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Change of State: (Transformations)
• Notation:
3 H2 (g, 5 bar, 100 K) 3 H2 (g, 1 bar, 50 K)
initial state final state
• Path: Sequence of intermediate states
• Process : Describes the Path

- Reversible : (always in Equilibrium)
- Irreversible : (defines direction of time)
- Adiabatic : (no heat transfer between system and surroundings)
- Isobaric : (constant pressure)
- Isothermal: (constant temperature)
- Isochoric: (Constant volume)
 Spontaneous changes
• A spontaneous change is a change that occurs by itself without any
outside assistance
• Spontaneous processes occur in a definite direction
• Nonspontaneous processes require outside assistance
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Zeroth Law of Thermodynamics.
If systems A and B are in thermal equilibrium and systems B
If systems A and B are in thermal equilibrium and systems B and C are in

thermal equilibrium then systems A and C are in thermal equilibrium.
First Law of Thermodynamics
 Energy can neither be created nor destroyed. It can only change forms.
 In any process in an isolated system, the total energy remains the same.
 For a thermodynamic cycle the net heat supplied to the system equals the
net work done by the system.
where: ΔE: is the internal energy

q = the amount of heat
w = work done
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1. Internal energy (E)
• Represented as E
• Sum of all kinetic and potential energies in a system
• Internal energy is a state function (dependent only on the initial and
final states)
• ΔE = Efinal – Einitial
• For a chemical reaction:
ΔE = Eproducts – Ereactants
2. Work (w): w = - PΔV
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3. Heat (q):
 That quantity flowing between the system and the surroundings that
can be used to change the temperature of the system and/or the
surroundings.
 Sign convention: If heat enters the system, then it is positive.
 Heat (q), like w, is a function of path. Not a state function
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 Units of Heat:
1. measured in calories [1 cal = heat needed to raise 1 g H 2O 1°C, from

14.5°C to 15.5°C]
2. The modern unit of heat (and work) is the Joule (1 cal = 4.184 J).
Case of first law equation:
 For Isochoric Process (V = 0) :
 For Isothermal Process (T = 0) :
 For Adiabatic Process (q = 0):
 For Cyclic Process (E = 0) :
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Examples:
1- Assume that the internal energy of a system decreases by 300 J while 200
J of work is done by a gas. What is the value of q? Is heat lost or gained
by the system?
E = -300 J; W = +200 J; E = q + W
q = (-300 J) + (200 J) = - 100 J; Heat Lost: q = -100 J
2- In a thermodynamic process, the internal energy of the system increases

by 500 J. How much work was done by the gas if 800 J of heat is
absorbed?
E = +500 J; q = +800 J; q = E - w
w = E- q= 500 J – 800 J; w = -300 J
3- A system absorbs 200 J of heat as the internal energy increases by 150 J.
What work is done by the gas? [q = +200 J, E = +150 J ]
w = E –q = 150 - 200 = -50 J;
4- During an isobaric expansion a steady pressure of 200 kPa causes the

volume of a gas to change from I L to 3 L. What work is done by the gas?
[ 1 L = 1 x 10-3 m3 ]
Work =- P(Vf – Vi) = (200,000 Pa)(3 x 10-3 m3 – 1 x 10-3 m3)
Work = - 400 J
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Enthalpy (H)
• Chemical changes may involve the release or absorption of heat.
• The heat transferred between the system and surroundings during a
chemical reaction carried out under constant pressure is called enthalpy,
H.
• Again, we can only measure the change in enthalpy, H.
• Mathematically,
H = Hfinal – Hinitial = E + PV
w = –PvV; E = q + w
H = E + PV = (qp + w) – w = qp
• For most reactions PV is small thus H = E
• Heat transferred from surroundings to the system has a positive enthalpy
(i.e., H > 0 for an endothermic reaction).

• Heat transferred from the system to the surroundings has a negative
enthalpy (i.e., H < 0 for an exothermic reaction).

• Enthalpy is a state function.
• qp = H = E + PV
• and E = H – PDV
Enthalpies of Reaction
• For a reaction,
• The enthalpy change that accompanies a reaction is called the enthalpy of
reaction or heat of reaction (Hrxn).
 For the reactions including gases, we can write:H = nRT

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Examples:
1- Calculate the change in internal energy for the reaction
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Thermodynamics

Describing systems requires:

Two classes of Properties:

The State of a System at Equilibrium:

• Path: Sequence of intermediate states

• Process : Describes the Path

If systems A and B are in thermal equilibrium and systems B and C are in

First Law of Thermodynamics

where: ΔE: is the internal energy

 Sign convention: If heat enters the system, then it is positive.

 Heat (q), like w, is a function of path. Not a state function

1. measured in calories [1 cal = heat needed to raise 1 g H 2O 1°C, from

Case of first law equation:

 For Isochoric Process (V = 0) :

 For Isothermal Process (T = 0) :

 For Adiabatic Process (q = 0):

 For Cyclic Process (E = 0) :

q = (-300 J) + (200 J) = - 100 J; Heat Lost: q = -100 J

2- In a thermodynamic process, the internal energy of the system increases

w = E- q= 500 J – 800 J; w = -300 J

3- A system absorbs 200 J of heat as the internal energy increases by 150 J.

What work is done by the gas? [q = +200 J, E = +150 J ]

w = E –q = 150 - 200 = -50 J;

4- During an isobaric expansion a steady pressure of 200 kPa causes the

Work =- P(Vf – Vi) = (200,000 Pa)(3 x 10-3 m3 – 1 x 10-3 m3)

(i.e., H > 0 for an endothermic reaction).

enthalpy (i.e., H < 0 for an exothermic reaction).

reaction or heat of reaction (Hrxn).

 For the reactions including gases, we can write:H = nRT

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