Chapter 14: Entropy & Free Energy Flashcards Preview

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Flashcards in Chapter 14: Entropy & Free Energy Deck (20)
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1
Q

Spontaneous process

A

A process that does occur under a specific set of conditions (e.g. T, P, [& concentration for solutions])

*Often, but not always, exothermic reactions

2
Q

Nonspontaneous process

A

A process that does not occur under a specific set of conditions

3
Q

Entropy

A

entropy = S

The measure of a system’s energy dispersal;
the greater the dispersal, the greater the entropy

4
Q

W = XN

A

W = number of energetically equivalent different ways molecules in a system can be arranged

X = number of cells (“containers”)

N = number of moleucles

e.g. 2 cells & 4 molecules = 24 = 16

5
Q

Quantitative definition of entropy

A

Based on W = XN:

The most probable state is the one with the largest number of possible arrangements –> has the greatest entropy

6
Q

ΔU

A

Internal energy

ΔU = q + w
(heat + work)

**NOT ΔE** (taught wrong?)

7
Q

Standard entropy

A

S° = standard entropy

J/K • mol

The absolute entropy of a substance at 1 atm
*Temperature MUST be specified! NOT standard!
(Typically 25°C, however)

Entropies of substances are always positive, even for elements in standard states (UNLIKE standard enthalpy of formation ΔHf°)

8
Q

Trends of standard entropy (4)

A
  1. solid S° < liquid S° < gas S°
    gas phase of a substance always has highest standard entropy (disperses the most)
  2. monatomic species (e.g. He (g), Ne (g))
    larger molar mass = greater S°
  3. for 2 substances in same phase, with similar molar mass, the more complex substance has the greater S°
    e. g. F2 versus O3 (ozone is more complex)
  4. an element that exists in 2+ allotropic forms
    form with more mobile atoms = form with greater S°
    e.g. graphite has greater S° than diamond
9
Q

Factors that influence entropy of system

A
  1. Volume change
    As volume increases, space between translational energy levels decrease (causing more energy levels to become available, dispersing energy)
  2. Temperature change
    As temperature increases, kinetic energy increases, causing more energy levels to be accessible
  3. Molecular complexity
    The more complex a molecule is, the less restriction of motion there is (i.e. not only translational motion, but also rotational and vibrational)
  4. Molar mass
    The larger the molar mass, the less the space is between energy levels, and the more energy levels within which the system’s energy can be dispersed
  5. Phase change
    The more spaced out the molecules are, the more they can move, rearrange, and the greater the entropy
6. Chemical reaction
When product(s) have more gas molecules than reactant(s), the number of different arrangements (W) increases, and thus so does entropy
10
Q

Determining spontaneity based on ΔSsys & ΔSsurr

A

ΔSsys and -ΔSsurr –> spontaneous process

-ΔSsys and ΔSsurr​ –> spontaneous process

ΔSsys and ΔSsurr​ –> spontaneous process

*BOTH CANNOT BE NEGATIVE for spontaneous processes

11
Q

Relationship between ΔSsurr, -ΔHsys, and T

A

ΔSsurr α -ΔHsys α 1/T

i.e. ΔSsurr = -ΔHsys / T

12
Q

Second law of thermodynamics

A

For a process to be spontaneous as written (in foward direction), ΔSuniv must be positive

ΔSuniv = ΔSsys + ΔSsurr

System’s entropy can decrease so long as surrounding’s increase in entropy makes universe’s positive

13
Q

Equilibrium process

A

One that does not occur spontaneously in either the net forward nor reverse direction, but can be made to occur by the addition or removal of energy to a system at equilibrium

(e.g. melting of ice at 0°C, ice and water are in equilibrium with each other)

ΔSuniv = 0 = reaction is an equilibrium process

14
Q

Third law of thermodynamics

A

Entropy (S) of a perfect crystalline substance = 0 @ 0 K

As temperature increases, molecular motion increases, thus entropy increases

Entropy of any substance at any temperature above 0 K is greater than 0

*Impure/imperfect crystalline substance at 0 K has entropy (S) > 0

*This law enables us to determine experimentally the absolute entropies of substances

15
Q

[Gibbs] free energy (G)

A

free energy = energy available to do work

*state function*

ΔG = ΔH - TΔS

ΔG predicts spontaneity

16
Q

Determining spontaneity using ΔG when ΔG is <, >, or = 0

A

*At constant T & P*

ΔG < 0
SPONTANEOUS in forward direction
non-spontaneous in reverse

ΔG > 0
non-spontaneous in forward direction
spontaneous in reverse

ΔG = 0
system is at equilibrium

17
Q

Predicting the sign of ΔG using the signs of ΔH and ΔS

A

-ΔH and +ΔS = -ΔG
*ALWAYS SPONTANEOUS*
(Only enthalpy is negative.)

ΔH and -ΔS = +ΔG
*Always NONspontaneous*
(Only entropy is negative.)

  • ΔH and -ΔS = -ΔG if TΔS < ΔH
  • *(Both negative, low T = spontaneous.)**
  • ΔH and -ΔS = +ΔG if TΔS > ΔH
  • *(Both negative, high T = NONspontaneous.)**

+ΔH and +ΔS = -ΔG if TΔS > ΔH
(Both positive, high T = spontaneous.)
——
+ΔH and +ΔS = +ΔG if TΔS < ΔH
(Both positve, low T = NONspontaneous.)

18
Q

Standard free-energy change (ΔG°rxn)

A

The free-energy change for a reaction when it occurs under standard-state conditions

i.e. when reactants in their standard states are converted to products in their standard states

19
Q

Standard states of pure substances and solutions for ΔG°rxn

A

Gases: 1 atm pressure

Liquids: pure liquid

Solids: pure solid

Elements: most stable allotropic form at 1 atm & 25°C

Solutions: 1 molar concentration

20
Q

Standard free energy of formation (ΔG°f)

A

Standard free energy of formation of a compound:

Free-energy change that occurs when 1 mole of the compound is synthesized from its constituent elements, each in its standard state

*Value changes with temperature