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AP Chem Unit 9 Review | Thermodynamics and Electrochemistry in About 10 Minutes!

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Educational

Main ideas & lessons (Unit 9: Thermodynamics & Electrochemistry)

1) Entropy (S): what increases/disperses

  • Entropy = level of disorder / how dispersed particles are.
  • Lowest S: solids (very compact, not dispersed)
  • More S: liquids
  • More still: aqueous solutions
  • Highest S: gases (particles spread out the most)
  • Raising temperature increases entropy (molecules move faster)
  • Increasing gas volume increases entropy (more space to disperse)

Entropy changes in reactions:

  • Moving from all solid → mixture including gas ⇒ entropy increases
  • If all substances are gases:
    • More gas molecules ⇒ higher entropy
    • Example: going from 3 gas molecules → 2 gas molecules ⇒ entropy decreases

2) Calculating entropy change for reactions (ΔS)

Method:

  • ΔS = (sum of entropies of products) − (sum of entropies of reactants)

(Parallel concept to the earlier idea of calculating enthalpy change, ΔH.)

3) Thermodynamic favorability: Gibbs Free Energy (ΔG)

  • A thermodynamically favored process tends to occur.
  • Gibbs Free Energy = ΔG
    • ΔG < 0 ⇒ favored
    • ΔG > 0 ⇒ not favored

Standard conditions (°):

  • 25°C, 1 atm, and solutions at 1 M (1 mole/L)

4) Ways to calculate ΔG

Method A (sum of Gibbs energies):

  • ΔG = (sum of Gibbs free energies of products) − (sum of Gibbs free energies of reactants)

Method B (relationship with ΔH and ΔS):

  • ΔG = ΔH − TΔS
  • If you know any three, you can solve for the fourth.

Unit warning:

  • Entropy is usually in J, while ΔH and ΔG are usually in kJ → convert consistently.

5) Predicting favorability using ΔH and ΔS across temperatures

General sign-based “direction” rules:

  • Exothermic (ΔH negative) tends to be favored
  • Entropy increase (ΔS positive) tends to be favored

Consequences by sign patterns:

  • ΔH < 0 and ΔS > 0 ⇒ favored at all temperatures
  • ΔH > 0 and ΔS < 0 ⇒ not favored at any temperature
  • ΔH > 0 and ΔS > 0 ⇒ favored only at high temperatures
  • ΔH < 0 and ΔS < 0 ⇒ favored only at low temperatures

6) Kinetics vs thermodynamics: kinetic control / slow reactions

  • Some reactions are thermodynamically favored but occur extremely slowly (example: rusting).
  • This is kinetic control:
    • Often due to very high activation energy
  • Important distinction:
    • It may be not at equilibrium; the reaction is too slow to proceed appreciably.

Relationship between ΔG and equilibrium constant (K)

7) ΔG–K equation

  • ΔG = −RT ln(K)
    • R = 8.314 J/(mol·K)
    • T in Kelvin

Interpretation:

  • ΔG negative ⇒ K is large ⇒ lots of products form
  • ΔG positive ⇒ K is small (<< 1) ⇒ very little product forms

Solubility and dissolution example (KCl in water)

8) Thermodynamic reasoning for dissolving ionic compounds

  • Ionic compounds can dissolve and dissociate in water.
  • Example: KCl
    • ΔG = −5.3 kJ/mol ⇒ thermodynamically favored ⇒ KCl is soluble

Why it’s favored (enthalpy vs entropy):

  • ΔH is positive (endothermic; not favored)
  • ΔS is positive (entropy increase; favored)
  • The entropy increase outweighs the endothermic ΔH

Lesson:

  • Examine ΔH and ΔS for dissolution to predict whether it dissolves and under what conditions.

When unfavorable reactions still happen

9) How to make endergonic processes proceed

Option 1: add external energy

  • Examples: electricity, light
  • Mentioned uses:
    • charging a cellphone battery
    • photosynthesis (CO₂ + water)

Option 2: couple to a favorable reaction

  • Example idea: extracting Cu(s) from CuS (copper(II) sulfide) alone is unfavorable (ΔG positive).
  • But if you react the sulfide with oxygen (a favorable coupled process), the combined ΔG becomes negative, making the overall reaction favorable.

Electrochemistry: Galvanic cells and electrolytic cells

10) Galvanic cell (battery) fundamentals

  • A galvanic cell harnesses a spontaneous redox reaction to produce electrical energy.
  • Half-reactions:
    • Oxidation at the ANODE
    • Reduction at the CATHODE

Memory cues:

  • “RED CAT and AN OX”
  • Electrons flow through the wire from anode → cathode (likened to A/C direction)

Electrode mass change (for metallic electrodes):

  • Cathode increases in mass
  • Cue: “The CAT gets FAT.”

Salt bridge purpose:

  • Equalizes charge as ions react at electrodes
  • Uses inert ions that don’t react with cell contents
  • Common example ions: Na⁺ and NO₃⁻
  • Ion flow directions:
    • Cations → cathode
    • Anions → anode

11) Cell potential (Ecell) and standard reduction potentials

  • Cell voltage reflects thermodynamic favorability:
    • Positive voltage: favored
    • Negative voltage: not favored; needs external power
  • Standard reduction potentials come from a table.
  • Half-reactions are written as reductions.

Compute Ecell:

  • Ecell = Ecathode − Eanode

12) ΔGcell and electron count

  • Relationship:
    • ΔGcell = −(n)(Faraday’s constant)(Ecell)
  • Faraday’s constant:
    • 96,485 C/mol e⁻
  • You can compute one quantity if you have the other(s).
  • Standard conditions are commonly used for cell potential calculations, though real conditions may differ.

13) Nernst Equation (non-standard conditions)

  • For non-standard T and concentrations, use the Nernst Equation.
  • Conceptual guidance using reaction quotient Q:
    • Increase products or decrease reactants ⇒ Q increases ⇒ voltage decreases
    • Decrease products or increase reactants ⇒ Q decreases ⇒ voltage increases

Clarification:

  • This is not Le Châtelier’s Principle, because a galvanic cell is not at equilibrium during operation.
  • At equilibrium, the battery is “dead”:
    • Voltage = 0

14) Electrolytic cells (non-spontaneous)

  • Electrolytic cells are unfavored processes requiring external power.
  • Common use mentioned: plating metals

15) Electrolysis calculations (Faraday’s laws approach as taught)

Current-time-charge relationship:

  • I (amps) = Q (C) / t (s)
    • equivalently, Q = I·t
  • Example:
    • 10 A for 1 hour
    • Convert 1 hour → 3600 s
    • Q = 10 × 3600 = 36,000 C

Convert charge to mass plated:

  1. Convert charge to moles of electrons using Faraday’s constant
  2. Use the balanced half-reaction to relate electrons to metal atoms/ions
    • Example statement: “2 electrons produce one atom of nickel”
  3. Convert moles of Ni to grams using atomic mass

Speakers / sources featured

  • Jeremy Krug (host of the video)
  • UltimateReviewPacket.com (referenced as a resource)

Original video