Introduction
Chemical thermodynamics deals with the study of energy changes that occur during chemical reactions and physical transformations. It helps us understand whether a process is possible and how much energy is involved.
Thermodynamics does not tell how fast a reaction occurs, but it explains whether a reaction can occur or not.
Important Terms
System and Surroundings
- System: Part of the universe under study
- Surroundings: Everything outside the system
Universe relation:
Universe = System + Surroundings
Types of Systems
- Open System – Exchange of matter and energy
- Closed System – Exchange of energy only
- Isolated System – No exchange
State of a System
The state of a system is defined by variables like pressure, temperature, and volume.
State function depends only on initial and final state.
Example:
ΔU = Ufinal − Uinitial
Types of Properties
Intensive Properties
Do not depend on amount
Example: Temperature, pressure
Extensive Properties
Depend on amount
Example: Mass, volume
Thermodynamic Processes
Isothermal Process
Temperature remains constant
T = constant
Adiabatic Process
No heat exchange
q = 0
Isobaric Process
Pressure constant
P = constant
Isochoric Process
Volume constant
V = constant
Internal Energy
Internal energy is the total energy of a system.
Change in internal energy:
ΔU = q + w
Heat and Work
Heat (q)
Energy transferred due to temperature difference
Work (w)
Work done by system:
w = −PΔV
First Law of Thermodynamics
Energy cannot be created or destroyed.
Equation:
ΔU = q + w
Enthalpy
Enthalpy is the heat content of a system.
Definition:
H = U + PV
Change in enthalpy:
ΔH = ΔU + PΔV
Enthalpy Changes
Exothermic Reaction
Heat released
ΔH < 0
Endothermic Reaction
Heat absorbed
ΔH > 0
Thermochemical Equations
Example:
CH4 + 2O2 → CO2 + 2H2O
ΔH = −890 kJ
Hess’s Law
Total enthalpy change is independent of pathway.
ΔH = ΔH1 + ΔH2
Enthalpy of Different Processes
Enthalpy of Formation
ΔHf = enthalpy change when 1 mole compound formed
Enthalpy of Combustion
ΔHc = heat released on complete combustion
Enthalpy of Neutralization
ΔH = heat change when acid reacts with base
Bond Enthalpy
Energy required to break bonds
ΔH = Σ bond broken − Σ bond formed
Second Law of Thermodynamics
Entropy measures disorder.
Entropy change:
ΔS = Sfinal − Sinitial
Spontaneity of Reaction
Condition:
ΔG = ΔH − TΔS
If ΔG < 0 → spontaneous
If ΔG > 0 → non-spontaneous
Gibbs Free Energy
Free energy determines feasibility.
ΔG = ΔH − TΔS
Equilibrium Condition
At equilibrium:
ΔG = 0
Third Law of Thermodynamics
Entropy of perfect crystal at 0 K is zero
S = 0 at 0 K
Important Relationships
Heat capacity:
q = m c ΔT
Relation between Cp and Cv:
Cp − Cv = R
Summary of Important Equations
Internal energy:
ΔU = q + w
Work:
w = −PΔV
Enthalpy:
H = U + PV
Gibbs energy:
ΔG = ΔH − TΔS
Entropy:
ΔS = S2 − S1
Heat formula:
q = mcΔT
Common Mistakes
- Sign convention errors
- Confusing ΔH and ΔU
- Ignoring units (J, kJ)
- Wrong application of Hess’s law
Exam Tips
- Learn formulas properly
- Practice numerical questions
- Understand sign conventions
- Focus on derivations
Conclusion
Chemical thermodynamics helps in understanding energy changes and feasibility of reactions. It is a fundamental chapter for higher chemistry studies and competitive exams.
SEO Keywords
- Chemical Thermodynamics Class 11 notes
- Chapter 5 Chemistry notes
- Thermodynamics formulas
- Gibbs free energy explanation
FAQs
Q1. What is first law of thermodynamics?
ΔU = q + w
Q2. What is Gibbs free energy?
ΔG = ΔH − TΔS
Q3. What is entropy?
Measure of disorder