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Organic Chemistry 1 · Reaction Mechanisms

Thermodynamics and Kinetics

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On this page 7 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Worked example
  6. Key takeaway
  7. Study tools

In 30 seconds

Thermodynamics answers "how far?": whether a reaction is product-favored, via ΔG = ΔH - TΔS. A negative ΔG (exergonic) means spontaneous; a positive ΔG (endergonic) means nonspontaneous. Kinetics answers "how fast?" and is set by the activation energy Ea, the barrier reactants must climb to reach the . A reaction can be strongly exergonic yet extremely slow because of a high Ea, and product mixtures can be decided by the fastest pathway (kinetic control) or the most stable product (thermodynamic control).

Why this matters

Enzymes catalyze by lowering Ea without changing ΔG or K — the same reaction runs in milliseconds in a cell but years in a beaker. This also explains drug shelf life: a molecule may be thermodynamically unstable (ΔG < 0 for degradation) yet kinetically stable enough to store if its degradation Ea is high; refrigeration slows degradation further by shrinking k. In metabolism, ATP hydrolysis (ΔG° ≈ -30.5 kJ/mol) is coupled to endergonic processes to drive them forward.

The college version

1. Enthalpy and Entropy

Enthalpy (H) is heat content at constant pressure: ΔH < 0 is exothermic (heat released), ΔH > 0 is endothermic (heat absorbed). For organic reactions, ΔH is roughly bonds broken (endothermic) minus bonds formed (exothermic). Entropy (S) measures dispersal of energy and matter; splitting one molecule into two increases ΔS, while joining two into one decreases it.

2. Gibbs Free Energy

Gibbs free energy combines both effects: ΔG = ΔH - TΔS, with T in kelvin. ΔG < 0 is exergonic (spontaneous), ΔG > 0 is endergonic (nonspontaneous), and ΔG = 0 is . It links to the equilibrium constant: ΔG°= -RTlnK, so negative ΔG° gives K > 1. Exergonic ≠ exothermic: an endothermic reaction can still be exergonic if ΔS is positive enough that -TΔS wins, especially at high temperature.

3. Kinetics, Activation Energy, and Reaction Coordinate Diagrams

is set by Ea, not ΔG. The Arrhenius equation k = Ae-Ea/RT shows a higher barrier gives a smaller rate constant. A plots energy versus progress: transition states (‡) are peaks — highest-energy structures with partially broken/formed bonds — and intermediates are valleys — real species with finite lifetimes. ΔG is the reactants-to-products gap; Ea is the reactants-to-transition-state gap.

How it works

  1. Decide by computing ΔG = ΔH - TΔS.
  2. Remember spontaneity says nothing about speed; rate is set by Ea.
  3. Draw the reaction coordinate diagram, marking peaks (transition states) and valleys (intermediates).
  4. Identify the rate-determining step as the highest Ea.
  5. For competing products, ask whether the reaction is under kinetic control (lowest Ea) or thermodynamic control (most stable product).
  6. Confirm that a favorable reaction may still need heat, light, or a catalyst to overcome a large Ea.

Common confusions

Do not confuseWithDifference
ExergonicExothermicΔG < 0 vs. ΔH < 0
SpontaneityReaction rateΔG decides "if"; Ea decides "how fast"
Transition stateIntermediatePeak (no lifetime) vs. valley (real species)
Equilibrium constant KRate constant kK sets final ratio; k sets speed
Kinetic productThermodynamic productFastest to form vs. most stable
FavorableFastLarge driving force can coexist with huge Ea

Memory aids

"Thermo tells you WHERE you end up (ΔG); Kinetics tells you HOW HIGH the hill (Ea)." For sign conventions: EXergonic = energy EXits the free-energy account (spontaneous); EXothermic = heat EXits the system.

Quick review

Topic Recap

Thermodynamics predicts spontaneity through ΔG = ΔH - TΔS, linking to K; kinetics predicts speed through Ea, visualized on reaction coordinate diagrams with transition states (peaks) and intermediates (valleys). The two are independent: favorable ≠ fast, and product mixtures may be kinetically or thermodynamically controlled.

Knowledge Check

  1. A reaction has ΔH = -20 kJ/mol and ΔS = +60 J/(mol·K) at 298 K. Is it spontaneous?
  2. What is the difference between a transition state and an ?
  3. Why can a highly exergonic reaction still be extremely slow?
  4. Define kinetic versus thermodynamic control.
  5. Does a catalyst change ΔG or K? Why or why not?

Answers and Rationales

  1. Yes. ΔG = -20 - (298)(0.060) = -37.9 kJ/mol, negative, so spontaneous.
  2. A transition state is an energy maximum with partial bonds and no lifetime; an intermediate is an energy minimum that is a real, potentially observable species.
  3. Rate depends on Ea, not ΔG; a large barrier means few molecules react at any instant regardless of favorability.
  4. Kinetic control gives the product formed fastest (lowest Ea); thermodynamic control gives the most stable product (lowest ΔG), usually under reversible, high-temperature conditions.
  5. No. A catalyst lowers Ea equally for forward and reverse, speeding both by the same factor, so ΔG and K are unchanged.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Picture a ball on a hillside. Thermodynamics asks "does the ball end up lower than it started?" — if so, the change is spontaneous. Kinetics asks "how tall is the hill it must roll over?" — even a far-downhill trip is slow if a big ridge blocks the way. Thermodynamics is the map (net altitude change); kinetics is the road profile (how steep the climbs are). A trip can be downhill overall yet blocked by a steep climb.

Where it stops being exact: molecules do not literally roll like balls. Spontaneity is a statistical statement about the free-energy difference between the whole collection of reactants and products, not a single particle — and the "downhill" direction can even flip with temperature when the entropy term dominates.

Simple Example

Methane combustion (CH4 + 2O2 → CO2 + 2H2O) is strongly exergonic, yet methane and oxygen can sit mixed for years with no visible reaction, because the C–H and O=O bonds must first be broken — a very high Ea. A spark supplies that energy.

Worked example

  1. Convert units. Entropies are often tabulated in J/(mol·K) and enthalpies in kJ/mol; divide ΔS by 1000 to share units.
  2. Compute ΔG. With ΔH = -50 kJ/mol and ΔS = -100 J/(mol·K) = -0.100 kJ/(mol·K): ΔG = -50 - (298)(-0.100) = -50 + 29.8 = -20.2 kJ/mol.
  3. Interpret. ΔG < 0, so the reaction is spontaneous; the unfavorable entropy term is outweighed by the exothermic enthalpy.
  4. Read a diagram. On a two-step profile, identify peaks TS1 and TS2 and the intermediate valley between them; the rate-determining step has the larger Ea. Overall ΔG is independent of barrier heights.

Key takeaways

  • High yield: ΔG < 0 is spontaneous (exergonic); ΔG > 0 is nonspontaneous (endergonic).
  • High yield: Exergonic/exothermic refer to ΔG and ΔH — different quantities, never interchanged.
  • High yield: Negative ΔG tells you if a reaction can happen, not how fast — favorable ≠ fast.
  • T is in kelvin; ΔH and TΔS must share the same units.
  • ΔG°= -RTlnK: the more negative ΔG°, the larger K.
  • An endothermic reaction can be spontaneous if ΔS is sufficiently positive.
  • Transition states are maxima; intermediates are minima.
  • A catalyst lowers Ea for both directions, changing neither ΔG nor K.
  • High yield: Kinetic control favors the fastest product (lowest Ea); thermodynamic control favors the most stable (lowest ΔG).

Keep learning

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Practice Organic Chemistry 1

This lesson has no separate scored set. Practice draws from the subject’s question bank.

Study tools & related lessonsYou’ll learn to · Key vocabulary · Related

You’ll learn to

  • Define enthalpy, entropy, and Gibbs free energy, and use ΔG = ΔH - TΔS to predict spontaneity.
  • Distinguish exergonic/endergonic from exothermic/endothermic, and spontaneity from reaction rate.
  • Read a reaction coordinate diagram, locating transition states, intermediates, activation energy, and ΔG.
  • Explain kinetic versus thermodynamic control and why a favorable reaction is not necessarily fast.

Key vocabulary

Enthalpy (Δ H)
Heat change; exothermic negative, endothermic positive
Entropy (Δ S)
Change in disorder/dispersal
Gibbs free energy (Δ G)
ΔH - TΔS; net driving force
Exergonic / endergonic
ΔG < 0 / ΔG > 0
Exothermic / endothermic
ΔH < 0 / ΔH > 0
Spontaneity
Tendency to proceed on its own
Equilibrium
Forward/reverse rates balance; ΔG = 0
Reaction rate
How fast reactants become products
Activation energy (Eₐ)
Barrier to reach the transition state
Reaction coordinate diagram
Energy vs. progress plot
Transition state
Highest-energy peak with partial bonds
Intermediate
Local energy minimum (valley)
Kinetic vs. thermodynamic control
Fastest pathway vs. most stable product

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