5E: principles of chemical thermodynamics & kinetics (all GC)

Question 1

Thermodynamics laws: game analogy

• Zeroth Law*
• First Law*
• Second Law*
• Third Law*

Answer 1

• Zeroth law: says temperature exists and it can equilibriate, sets ground rules for the game
• First law: says change in energy always = sum of heat and work, says you can’t win, conservation of energy in thermodynamic processes
• Second law: says temperature and pressure flow downhill from greater to less and that net entropy of the universe is always increasing, says you can’t even break even
• Third law: says absolute zero is untenable since zero energy can’t be achieved, says you can’t even end the game

Question 2

Thermodynamic system-state function & path function

Answer 2

• State function: properties that describe the current state of the system. Not affected by how the systems got to their state, just the properties of their current state (ex: temperature, pressure, volume)
• Path function: properties that depend on the pathway used to achieve a state (ex: work, heat)
• Thermodynamic systems can’t describe systems on a molecular scale. Thermodynamic systems average out all molecular interactions to find an average. Molecular scale means the sample size is too small
• Internal energy: collective energy of molecules measured on a microscopic scale. Many different types!

Question 3

Internal Energy Types

Answer 3

• Vibrational energy: created by back-and-forth motion of atoms
• Rotational energy: created by rotation of a molecule around its center of mass
• Translational energy: created by movement of a molecule’s center of mass
• Electronic energy: potential electric energy created by attractions between electrons and their nuclei
• Intermolecular potential energy: intermolecular dipole forces
• Rest mass energy: described by E = mc^2

Question 4

What are the 2 ways to transfer energy between systems?

Answer 4

• heat & work
• Difference between heat and work: directional collisions vs. random collisions
• Work is done by energy transfer through _ordered molecular collisions_
  
  • ​more constrainsed the molecules are (higher the P, lower the V)=greater capacity to do work
• Heat is done through _random collisions_ between high energy & low energy particles

Question 5

Zeroth Law

Answer 5

• Concept of temperature
  
  • temperature: thermal energy per mol of molecules

Question 6

First Law of Thermodynamics

Answer 6

• conservation of energy in thermodynamics
  
  • total energy of system & surroudings always conserved
  • ΔE = q + w
    • Energy change of a system = heat flow into the system + work done on the system
    • For closed systems, only internal energy change takes place, ΔU is used instead of ΔE
      
      • If no change in volume occurs as well, omit work. ΔU = q
  • Energy transferred out of a system during expansion (ΔE negative)
  • Energy transferred into a system during contraction (ΔE positive)

Question 7

PV Diagram:

Work

Answer 7

• Work done = area under or enclosed by curve
• Work= any energy transfer that’s not clear
• System performs PV work by changing it’s size or shape using energy from the system
  
  • Ex: piston expanding
• Formula for PV work: W= PΔV
  
  • ​Volume must be constant
  • Note how negative work=work done by the system (if V expands)
  • Positive work=work done ON THE SYSTEM
    
    • ​Keep in mind the MCAT might try to define work done by the system as positive
• PV diagram: x-axis=P, y-axis=V
  
  • Work=area under the PV curve
  • Work is a path function, different curve/path results in a different

amount of work

Question 8

Second Law: concept of entropy

Answer 8

• ΔS_system + ΔS_surroundings = ΔS_universe
• Entropy as a measure of “disorder”
  
  • Entropy: energy trying to spread itself evenly throughout the universe
  • Entropy of an isolated system will never decrease. Since the universe is an isolated system, entropy of the universe never decreases
  • Rxn must increase entropy of the universe to proceed
    
    • Entropy increases with number, size, volume, and temperature
  • ​​Relative entropy for gas, liquid, and crystal states

Question 9

Third Law: absolute zero

Answer 9

• Says zero entropy can only take place at absolute zero.
• However, this is unattainable so third law can only be realized in theory

Question 10

• Measurement of heat changes (calorimetry)
• What are the two types of calorimetry?

Answer 10

• Calorimetry: measuring changes in heat flow of rxn by monitoring temp change of a calorimeter coupled to the rxn to find the change in enthalpy
• Two types: constant pressure and constant volume
  
  • Coffee cup calorimeters (constant pressure)
    
    • Rxn takes place in chamber with open top. Constant pressure of local atmosphere dictates pressure
    • Use insulated chamber to prevent heat exchange w/ surroundings
    • Used to measure heats of reaction and enthalpy b/c no heat from rxn is lost to surroundings
  • Bomb calorimeter (constant volume):
    
    • Rxn takes place in rigid, sealed off container
    • Use insulated chamber to prevent heat exchange with

surroundings

    * Measures internal energy change by finding q **from q =**

CΔT

        * *​C of calorimeter is known, T change can be* * measured after the reaction*

* *​​important for calorimeter chamber to be thermally insulated from the surroundings. No heat exchange between system & surroundings* *

Question 11

• Heat Capacity, what gives greater heat capactiy?
• Specific Heat Capacity

Answer 11

• Heat capacity: How much E must be added to a substance to change its temp by 1 C or K
• More bonds in a molecule = greater heat capacity
  
  • This is b/c E is redirected to stretching these bonds instead of raising T
• More IMF’s between molecules = greater heat capacity
  
  • This is b/c IMF’s must be broken using E to raise temperature. Some E has to be redirected to do this
• T will always increase when E is added to a substance at a constant V and P
• Formula for heat capacity: q = mcΔT
• Specific heat capacity: intrinsic property, heat capacity per unit mass
  
  • Q = mcΔT
• ΔHrxn = -ΔHcalorimeter

Question 12

Heat transfer:

1. Conduction:

2. Convection

3. Radiation

Answer 12

• Conduction: heat transfer through molecular collisions. Requires direct physical contact
  
  • Thermal conductivity (k): an object’s ability to conduct heat, depends on composition and temperature (composition more so)
• Convection: heat transfer through fluid
  
  • Driven by differences in pressure and density, drives warm fluid in direction of cooler fluid
  • Ex. Hot air rises, causing cooler air from ocean to move in
• Radiation: thermal energy transfer via electromagnetic waves
  
  • Newton’s law of cooling: a body’s rate of cooling is proportional to the temperature difference between a body and it’s environment
  • Emissivity: fraction of radiation absorbed by a surface (depends on surface composition)
    
    • Higher emissivity = higher amount of radiation absorbed

Question 13

Endothermic/Exothermic reactions

Enthalpy (ΔH)

Answer 13

• Enthalpy: Used to measure changes in heat. Sum of internal energy and work.
  
  • ΔH=ΔU+PΔV
  • H = enthalpy
  • U = internal energy
  • P = pressure, V = volume
• When only PV work is done at constant pressure and volume, ΔH = q
  • This is b/c no non=PV work means PΔV value is 0

Question 14

• Standard enthalpy of formation (Hf^o):
• standard state
• postive & negative Hf^o change

Answer 14

• Standard enthalpy of formation (Hfo): change in heat for a reaction that creates one mole of that compound from its raw elements in their standard states
• Standard state: reference form of a substance
• Positive Hf^o change = endothermic reaction, heat flows into

system

• Negative Hf^o change = exothermic reaction, heat flows out of

system

Question 15

Hess’ Law of Heat Summation

Answer 15

• Sum of enthalpy changes for each step equals total energy change
• Forward reaction has the opposite change in enthalpy as reverse
• Energy reaction diagram
  
  • Y-axis:“energy” can be enthalpy, Gibbs, or energy
  • x-axis: rxn progresses
  • Difference between initial and final energy states is constant regardless of changes in activation energy

Question 16

Gibbs Free energy: ΔG

Answer 16

• way of finding entropy change in both system and surroundings using only information about the system,
• ΔG = ΔH – TΔS
  
  • ​​ΔG shows how much non-PW work is “free” for a reaction, determines if a reaction is spontaneous or not
  • ΔG, ΔH, and ΔS only refer to changes in the SYSTEM, not the surroundings
• Both ΔS and ΔH are required in determining if a reaction proceeds spontaneously
• G is an extensive property, state function (like enthalpy, entropy)

Question 17

Spontaneous reactions and ΔG^o

Answer 17

• Negative ΔG^o means a reaction proceeds spontaneously
  
  • ΔG^o value depends on both ΔH and ΔS

Question 18

Heat of fusion, heat of vaporization

Answer 18

• Heat of vaporization is usually greater than heat of fusion
  
  • Usually break more bonds from liquid to gas than solid to liquid

Question 19

Phase Diagram: pressure & temperature

Critical Temperature

Critical pressure

Critical Point

Answer 19

• Phase diagram indicates phases of a substance at different pressures and temperatures
• Critical temperature: temperature at which a substance cannot be liquefied regardless of pressure
  
  • Critical pressure: pressure at the critical temperature
  • Critical point: point defined by the intersection of the critical T and P
  • Solids and gases favored at extreme conditions, liquids favored in intermediate conditions
  • Difference in water’s phase diagram
    
    • ​Line between solid and liquid phase has NEGATIVE slope. This allows solid water to be less dense than liquid water
    • Normally, this line has a POSITIVE slope

Question 20

Rate law, rate constant

Answer 20

• Rate law: finds the rate of a reaction by factoring in rate law and concentrations of reactants
• Experimentally determining rate law:
  
  • 1. Compare reaction rates in two trials where concentrations of all reactants except for one stay constant
  • 2. Compare how changing concentration of one reactant affects reaction rate
       * If reactant A is doubled and rate doubles, then order of reactant A is one. If A is doubled and rate quadruples, then order is two

Question 21

Reaction Order

• Zero Order*
• First Order*
• Second/Third Order*

Answer 21

• Rate_forward = k[A]^a[B]^b, rate law = a + b
• Zero order: rxn rate is independent of concentration of any reactant
  
  • Graph: [A] over t, slope of –k is linear and downwards
• First order: rxn rate is directly proportional to concentration of a single reactant
  
  • Graph: ln[A] over t, slope of –k is linear and downwards
• Second/Third order: rxn rate is proportional to a single reactant’s concentration raised to the second/third power OR the product of
• concentrations of multiple reactants

Question 22

Rate-determining Step

Answer 22

• In multistep reactions, rate of the slowest occurring step is the rate determining step
• Rate-determining step determines the rate law for the reaction

Question 23

Dependence of reaction rate upon temperature

Activation Energy

Answer 23

• Requirements to reach activation E
  
  • 1) Particles must be traveling at a sufficient velocity
  • 2) Particles must collide in the correct orientation
• Activated complex or transition state
• Interpretation of energy profiles showing energies of reactants, products, activation energy, and ΔH for the reaction

Question 24

Dependence of reaction rate upon temperature

Use of Arrhenius Equation

Answer 24

• Used to find rate constant (k). Rate constant dictates rate of reaction (directly proportional)
  
  • If it affects rate constant, it affects rate of reaction in the same way
  • Rate constant is inversely proportional to activation energy, directly proportional to temperature and collision frequency
• p = steric factor, takes into account correct orientation of molecules hitting each other and frequency of collisions

Question 25

What are the 3 factors that affects K in the Arrehius Equation?

Answer 25

• 1) Pressure: higher pressure increases rate constant (relevant for gases)
• 2) Catalysts: presence of catalyst increases rate constant
• 3) Temperature: temperature increases rate constant

Question 26

Kinetic control versus thermodynamic control of a reaction

Answer 26

• Kinetic control: controls speed of a reaction, governed by Ea
  
  • Pertains to top half of a reaction energy diagram (where transition state peak is)
• Thermodynamic control: controls whether or not a reaction will occur, governed by ΔG of the reaction
  
  • Pertains to bottom half of reaction energy diagram

Question 27

Catalysts

What are the two types of catalysts?

Total rate law of catalyzed reactions:

Answer 27

• If concentration of catalyst far outweighs reactant, then rxn order is zero
• Increases rates of forward and reverse reactions
• Can lower Ea, increases steric factor, or both
  
  • However, does not affect overall change in energy, difference between final energies of reactants and products stays the same, only the Ea changes

Two types: heterogeneous and homogenous

• Heterogeneous: different phase than reactants of products. Rate of catalysis depends on strength of attraction between reactant and catalysts (ex: liquid catalyst adsorbs, or sticks to, a solid reactant). Rxn rate enhanced by increasing surface area of catalyst
  
  • Too little attraction: catalyst can’t adsorb to reactant, little effect on reaction rate
  • Too much attraction: catalyst doesn’t want to let go, little effect on reaction rate
  • Want to get just the right amount
• Homogeneous: same phase as reactants and products
• Total rate law of catalyzed reactions = rate law of original reaction + rate law of catalyzed reaction
  
  • Ex: Rate law of acid catalyzed reaction = k₀[A] + k_H+[H+][A]
  • B/c even with catalyst, original reaction will still occur

Question 28

Equilibrium in reversible chemical reactions

Law of Mass Action

Equilibrium Constant K_eq

Reaction Quotient (Q)

proportionals K_eq & Q

Answer 28

• Law of Mass Action
  
  • Rate of any reaction is proportional to the concentrations of the

different species in the reaction (according to the equilibrium constant)

• Equilibrium Constant (see pic attacked for equation)
• Reaction Quotient (Q)
  
  • Same formula as equilibrium constant but reaction isn’t necessarily at equilibrium
  • Value of Q relative to Keq can tell us about _direction rxn will proceed:_
    
    • Q = K: rxn is at equilibrium
    • Q > K: rxn shifts leftwards, towards reactants (too much P)
    • Q < K: rxn shifts rightwards, towards products (too much R)

Question 29

Application of Le Chatelier’s Principle

Answer 29

• Change in concentration, pressure, or temperature causes a system at equilibrium to shift in a direction to reduce the stress
  
  • Change in concentration—-> rxn shifts in direction of less concentrated side
  • Increase in pressure—->rxn shifts towards side of rxn with less moles of gas
  • Increase in temperature—–>rxn shifts away from the side of the equation with more heat (towards gaseous side)

Question 30

Relationship of the equilibrium constant and ΔG^o

What is the equation?

• If Keq = 1, then ΔG^o =?
• If Keq > 1, then ΔG^o = ?
• If Keq < 1, the n ΔG^o = ?

Answer 30

• To determine if a reaction is spontaneous under specific conditions (not just standard conditions), both Q (concentrations of reactant and product) and ΔG must be considered

o ΔG = ΔG^o + RTln(Q)

* **​**Can be rewritten as ΔG<sup>o</sup>= -RTLin (K)
* **If Keq = 1,** then ΔG<sup>o</sup> = 0
* **If Keq \> 1**, then ΔG<sup>o</sup> \< 0
* **If Keq \< 1,** the n ΔG<sup>o</sup> \> 0
* This only determines **spontaneity at a specific temperature** (Keq and ΔG<sup>o</sup> values depend on **_TEMPERATURE)_**