Thermodynamics Flashcards
System
Part of universe in which observations are made
Surroundings
Remaining universe except system constitutes surroundings
Open System
Exchange of matter and energy between system and surroundings
Closed system
No exchange of matter, but exchange of energy is possible between system and surroundings
State variables
Values depend upon the state of the system and not on how it is reached
Isolated system
No exchange of energy or matter between system and surroundings
```
Internal energy (U)
Represents the total energy of the system
How to change internal energy of system
- Heat passes into or out of system
- Work done by or on system
- Matter enters or leaves system
Adiabatic system
Thermodynamic process without the transfer of heat
q = 0
Sign convention
Work done on system = +
Increase in internal energy = +
Heat is transferred from the surroundings = +
First Law of therm
The energy of an isolated system is constant
delta U = q + W
Isothermal
Temperature remains constant
Isobaric
Pressure remains constant
Isochoric
Volume remains constant
Cyclic process
When a system undergoes different number of processes and finally returns to its initial state
Work done (with sign convention)
W = -PdV
When body is compressed = Work done on system = + sign
However, Vf - Vi = negative
Hence, to maintain sign convention we put a (-) sign in front
Reversible change
Change is brought out in such a way that the process could, at any moment, be reversed by an infinitesimal change
Here, system and surroundings are almost always in equilibrium
Work done in an isothermal process (reversible)
W = -2.303 nRT log Vf / Vi
(-) for sign convention
Free expansion
No work is done during free expansion whether the process is reversible or irreversible
Work done in isothermal irreversible
q = -W = p(Vf - Vi)
Enthalpy
- Sum of internal energy and product of volume and pressure
- delta H = delta U + p(delta V)
Sign convention for enthalpy
(- exothermic)
(+ endothermic)
When is difference between enthalpy and internal energy significant
- For solids and liquids as expansion in volume is very less, H roughly = U
- For gases, it is much more significant as gases change volume on expanding
For gases, enthalpy change
delta H = delta U + delta (Ng) RT
p(delta V) = delta (n) RT
delta n = no. of moles of products - no. of moles of reactants (both of gaseous)
Extensive properties
Value depends upon the quantity / size of matter present
Intensive property
Do not depend upon quantity or size of matter present
Heat capacity
q = C (delta T)
Quantity of heat required to raise temperature of the substance by one degree celsius
Molar Heat Capacity
Quantity of heat required to raise temperature of one mole of the substance by one degree celsius
delta Q = m(C/n) delta T
Specific heat capacity
Quantity of heat required to raise the temperature of one unit mass of a substance by one degree celsius
q = (mc) delta T = C delta T
Relation between Cp and Cv
Cp - Cv = R
How to measure delta U
- Bomb Calorimeter
- No work is done (delta V = 0)
Reaction Enthalpy
Enthalpy change accompanying a reaction
delta rH = sigma (ai H) - sigma (bi H)
Enthalpy change for a reaction when all participating substances are in standard state (1 bar, eg.)
Denoted by (-) superscript
Standard enthalpy of fusion / Molar enthalpy of fusion
Enthalpy change that accompanies melting of one mole of a substance in standard state
Standard enthalpy of vaporization
Amount of heat required to vaporize one mole of a liquid at constant temperature and under standard pressure
Change in enthalpy when one mole of a solid substance sublimes at a constant temperature and standard pressure
Standard enthalpy of formation
For the formation of one mole of a compound from its elements in their most stable states of aggregation.
Symbol: delta f H (f indicates one mole)
Why is:
CaO + CO2 –> CaCO3
NOT standard enthalpy of formation?
- Standard entalpy of formation has one mole formed from its constituent elements
- Here, its formed from substances of other compounds
- Coefficients refer to number of moles of reactants and products
- Numerical value of delta r H refers to the number of moles of substances specified by an equation
- When a chemical reaction is reversed, the sign of delta H is also reversed
Hess’s Law
If a reaction takes place in several steps then its standard reaction enthalpy is the sum of standard enthalpies of intermediate reactions into which the overall reaction may be divided at the same temperature
Why is enthalpy in standard state zero?
- Already in stable states, hence no need of changing
- This serves as a reference point
Standard enthalpy of Combustion
(delta cH)
Enthalpy change per molecule of a substance
Enthalpy of atomization
(delta aH)
- Enthalpy change in breaking one mole of bonds to completely obtain atoms in the gase phase
- For diatomic, it is bond dissociation
Ionization Enthalpy
Refers to the enthalpy required to remove an electron from an isolated gaseous atom in its ground state
Bond enthalpy
(delta bond H)
Change in enthalpy when one mole of covalent bonds of a gaseous covalent compound is broken
Mean bond enthalpy
(enthalpy of atomization / number of bonds)
Formula of standard reaction enthalpy
delta r H = sigma (bond enthalpy reactants) - sigma (bond enthalpy products)
When to use products - reactants and vice versa
Using hess’ law of formation -> products - reactants
Average bond energies to break -> reactant - products
Lattice enthalpy
One mole of an ionic compound dissociates into its ions in gaseous state
Born Haber cycle
The Born Haber cycle is a cycle of enthalpy change of process that leads to the formation of a solid crystalline ionic compound from the elemental atoms in their standard state and of the enthalpy of formation of the solid compound such that the net enthalpy becomes zero.
Born Haber Cycle Diagram
Refer to TB exercises
Enthalpy of solution
Enthalpy change when one mole of it dissolves in a specific amount of solvent
enthalpy sol = lattice enthalpy + hydration enthalpy
Enthalpy of dilution
Refers to the enthalpy change associated with the dilution process of a component in a solution at a constant pressure.
When 2 moles of C2H6 are completely burnt, 3129kj of heat is liberated. Calculate the heat of formation of C2H6. delta Hf for CO2 and H2O are -395kJ and -286kJ respectively
2C2H6 + 7O2 –> 4CO2 + 6H2O
delta heat formation = delta heat (products) - delta heat (reactants)
= [4 * (fH CO2) + 6 * (fH H2O)] - [2 * fH CO + 7 * fH O2] - 3129
= -167
Now for one mole (heat of formation is one mole): -83.5kJ
Given bond enthalpy of H-H and Cl-Cl are 430 kj/mol and 240 kj/mol respectively and delta fH is -90 kj/mol. Find the bond enthalpy of HCl
90 + 430/2 + 240/2 = 425
Spontaneous Reaction
- Having the potential to proceed without the assistance of external agency
- Irreversible
Entropy
Measure of disorderness of a system
Formula of entropy
delta S = Q / T
Entropy at equilibrium
At equilibrium, entropy reaches maximum and change in entropy, delta S = 0
Entropy change on melting vs freezing
Usually on freezing, the particles become closer and less disordered, hence they have lesser entropy. On the other hand, on melting it usually has higher entropy
What is total entropy?
delta S total = delta S system + delta S surroundings > 0
For same amount of heat is entropy more at higher or lower temp
It is higher at lower temp
Formula of Gibbs Energy
delta G = delta H - T delta S
enthalpy change - temperature * entropy change
When is a reaction spontaneous (Gibbs Energy)
- If delta G is < 0 –> Spontaneous
- If delta G is >0 –> Non-spontaneous
Second Law of Thermodynamics
The Second Law of Thermodynamics states that the total entropy (a measure of disorder or randomness) of an isolated system always increases over time
Third Law of Thermodynamics
Entropy of any pure crystalline substance approaches zero as the temp. approaches absolute zero
Gibbs Free Energy Formula reversible
delta r G = delta r H - T delta S = -RT ln K
= -2.303 RT log K
where log is base 10
Is bond formation exothermic or endothermic?
Exothermic (delta H = -ve)
delta H value on stability?
Positive means unstable
Negative is stable
Which does more work: reversible or irreversible
Irreversible does more work as it has more area under the curve
X requires more heat to vaporize than Y. Which has higher enthalpy of vaporization
X since it consumes higher heat energy
Why do we use reactants - products in some cases
That’s cause that’s for bond enthalpy of reactants - bond enthalpy of products.
Otherwise we use enthalpy of formation of products - enthalpy of formation of reactants
Relation between entropy of system and surrounding
Inversely proportional
