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Chem - Gr11 > Thermodynamics > Flashcards

Thermodynamics Flashcards

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1
Q

System

A

Part of universe in which observations are made

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2
Q

Surroundings

A

Remaining universe except system constitutes surroundings

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3
Q

Open System

A

Exchange of matter and energy between system and surroundings

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4
Q

Closed system

A

No exchange of matter, but exchange of energy is possible between system and surroundings

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5
Q

State variables

A

Values depend upon the state of the system and not on how it is reached

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6
Q

Isolated system

A

No exchange of energy or matter between system and surroundings

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7
Q

```

Internal energy (U)

A

Represents the total energy of the system

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8
Q

How to change internal energy of system

A
  1. Heat passes into or out of system
  2. Work done by or on system
  3. Matter enters or leaves system
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9
Q

Adiabatic system

A

Thermodynamic process without the transfer of heat

q = 0

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10
Q

Sign convention

A

Work done on system = +
Increase in internal energy = +
Heat is transferred from the surroundings = +

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11
Q

First Law of therm

A

The energy of an isolated system is constant
delta U = q + W

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12
Q

Isothermal

A

Temperature remains constant

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13
Q

Isobaric

A

Pressure remains constant

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14
Q

Isochoric

A

Volume remains constant

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15
Q

Cyclic process

A

When a system undergoes different number of processes and finally returns to its initial state

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16
Q

Work done (with sign convention)

A

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

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17
Q

Reversible change

A

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

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18
Q

Work done in an isothermal process (reversible)

A

W = -2.303 nRT log Vf / Vi
(-) for sign convention

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19
Q

Free expansion

A

No work is done during free expansion whether the process is reversible or irreversible

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20
Q

Work done in isothermal irreversible

A

q = -W = p(Vf - Vi)

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21
Q

Enthalpy

A
  • Sum of internal energy and product of volume and pressure
  • delta H = delta U + p(delta V)
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22
Q

Sign convention for enthalpy

A

(- exothermic)
(+ endothermic)

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23
Q

When is difference between enthalpy and internal energy significant

A
  • 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
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24
Q

For gases, enthalpy change

A

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)

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25
Q

Extensive properties

A

Value depends upon the quantity / size of matter present

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26
Q

Intensive property

A

Do not depend upon quantity or size of matter present

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27
Q

Heat capacity

A

q = C (delta T)
Quantity of heat required to raise temperature of the substance by one degree celsius

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28
Q

Molar Heat Capacity

A

Quantity of heat required to raise temperature of one mole of the substance by one degree celsius
delta Q = m(C/n) delta T

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29
Q

Specific heat capacity

A

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

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30
Q

Relation between Cp and Cv

A

Cp - Cv = R

31
Q

How to measure delta U

A
  • Bomb Calorimeter
  • No work is done (delta V = 0)
32
Q

Heat of Reaction / Enthalpy of reaction

A

Heat absorbed or evolved at constant pressure

33
Q

Reaction Enthalpy

A

Enthalpy change accompanying a reaction
delta rH = sigma (ai H) - sigma (bi H)

34
Q

Standard enthalpy of reaction

A

Enthalpy change for a reaction when all participating substances are in standard state (1 bar, eg.)
Denoted by (-) superscript

35
Q

Standard enthalpy of fusion / Molar enthalpy of fusion

A

Enthalpy change that accompanies melting of one mole of a substance in standard state

36
Q

Standard enthalpy of vaporization

A

Amount of heat required to vaporize one mole of a liquid at constant temperature and under standard pressure

37
Q

Standard enthalpy of sublimation

A

Change in enthalpy when one mole of a solid substance sublimes at a constant temperature and standard pressure

38
Q

Standard enthalpy of formation

A

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)

39
Q

Why is:
CaO + CO2 –> CaCO3
NOT standard enthalpy of formation?

A
  • Standard entalpy of formation has one mole formed from its constituent elements
  • Here, its formed from substances of other compounds
40
Q

Convention regarding thermochemical equations

A
  • 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
41
Q

Hess’s Law

A

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

42
Q

Why is enthalpy in standard state zero?

A
  1. Already in stable states, hence no need of changing
  2. This serves as a reference point
43
Q

Standard enthalpy of Combustion

(delta cH)

A

Enthalpy change per molecule of a substance

44
Q

Enthalpy of atomization

(delta aH)

A
  • Enthalpy change in breaking one mole of bonds to completely obtain atoms in the gase phase
  • For diatomic, it is bond dissociation
45
Q

Ionization Enthalpy

A

Refers to the enthalpy required to remove an electron from an isolated gaseous atom in its ground state

46
Q

Bond enthalpy

(delta bond H)

A

Change in enthalpy when one mole of covalent bonds of a gaseous covalent compound is broken

47
Q

Mean bond enthalpy

A

(enthalpy of atomization / number of bonds)

48
Q

Formula of standard reaction enthalpy

A

delta r H = sigma (bond enthalpy reactants) - sigma (bond enthalpy products)

49
Q

When to use products - reactants and vice versa

A

Using hess’ law of formation -> products - reactants
Average bond energies to break -> reactant - products

50
Q

Lattice enthalpy

A

One mole of an ionic compound dissociates into its ions in gaseous state

51
Q

Born Haber cycle

A

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.

52
Q

Born Haber Cycle Diagram

A

Refer to TB exercises

53
Q

Enthalpy of solution

A

Enthalpy change when one mole of it dissolves in a specific amount of solvent
enthalpy sol = lattice enthalpy + hydration enthalpy

54
Q

Enthalpy of dilution

A

Refers to the enthalpy change associated with the dilution process of a component in a solution at a constant pressure.

55
Q

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

A

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

56
Q

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

A

90 + 430/2 + 240/2 = 425

57
Q

Spontaneous Reaction

A
  • Having the potential to proceed without the assistance of external agency
  • Irreversible
58
Q

Entropy

A

Measure of disorderness of a system

59
Q

Formula of entropy

A

delta S = Q / T

60
Q

Entropy at equilibrium

A

At equilibrium, entropy reaches maximum and change in entropy, delta S = 0

61
Q

Entropy change on melting vs freezing

A

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

62
Q

What is total entropy?

A

delta S total = delta S system + delta S surroundings > 0

63
Q

For same amount of heat is entropy more at higher or lower temp

A

It is higher at lower temp

64
Q

Formula of Gibbs Energy

A

delta G = delta H - T delta S
enthalpy change - temperature * entropy change

65
Q

When is a reaction spontaneous (Gibbs Energy)

A
  • If delta G is < 0 –> Spontaneous
  • If delta G is >0 –> Non-spontaneous
66
Q

Second Law of Thermodynamics

A

The Second Law of Thermodynamics states that the total entropy (a measure of disorder or randomness) of an isolated system always increases over time

67
Q

Third Law of Thermodynamics

A

Entropy of any pure crystalline substance approaches zero as the temp. approaches absolute zero

68
Q

Gibbs Free Energy Formula reversible

A

delta r G = delta r H - T delta S = -RT ln K
= -2.303 RT log K
where log is base 10

69
Q

Is bond formation exothermic or endothermic?

A

Exothermic (delta H = -ve)

70
Q

delta H value on stability?

A

Positive means unstable
Negative is stable

71
Q

Which does more work: reversible or irreversible

A

Irreversible does more work as it has more area under the curve

72
Q

X requires more heat to vaporize than Y. Which has higher enthalpy of vaporization

A

X since it consumes higher heat energy

73
Q

Why do we use reactants - products in some cases

A

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

74
Q

Relation between entropy of system and surrounding

A

Inversely proportional

Chem - Gr11 (8 decks)

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