Reaction Energy Profile
Arrhenius Plot: ln(k) vs 1/T
Interactive visualization of reaction energy profiles - Explore activation energy, transition states, intermediates, and catalysis effects on potential energy surfaces
A reaction coordinate diagram plots the potential energy of the reacting system against the reaction progress. The horizontal axis (reaction coordinate) represents the progress from reactants to products through bond rearrangements. Key features: Reactant energy level, product energy level, activation energy barrier (Ea), transition state (saddle point, ‡), and any reaction intermediates (local minima along the path).
The activation energy is the minimum energy that reactant molecules must possess to undergo the reaction. It creates an energy barrier between reactants and products. For the forward reaction: Ea(fwd) = E(transition state) - E(reactants). For the reverse reaction: Ea(rev) = E(transition state) - E(products). The Arrhenius equation k = A·exp(-Ea/RT) shows how Ea affects the rate constant exponentially. Higher Ea → slower reaction at given temperature.
ΔH = E(products) - E(reactants). Exothermic: ΔH < 0, products lower in energy, energy released to surroundings. Endothermic: ΔH > 0, products higher in energy, energy absorbed from surroundings. Note: ΔH determines thermodynamic favorability (ΔG = ΔH - TΔS), while Ea determines kinetic rate. A reaction can be thermodynamically favorable (ΔG < 0) but kinetically slow (high Ea) — this is why catalysts matter.
The transition state (activated complex) is the highest-energy point along the reaction coordinate. It represents a fleeting molecular configuration where old bonds are partially broken and new bonds are partially formed. Key properties: It is a saddle point on the potential energy surface (maximum along reaction coordinate, minimum in all other directions). It cannot be isolated or observed directly. The Hammond Postulate: the transition state resembles whichever species (reactant or product) is closer in energy — for exothermic reactions the TS is reactant-like (early), for endothermic reactions it is product-like (late).
1. Start from the left (reactants) and follow the curve to the right (products). 2. Peaks are transition states (TS‡) — the highest points represent energy barriers. 3. Valleys between peaks are reaction intermediates — stable or semi-stable species. 4. The height difference from reactants to the first TS gives the forward activation energy Ea(fwd). 5. The height difference from products to the last TS gives the reverse activation energy Ea(rev). 6. The overall height difference between reactants and products gives ΔH.
When a reaction has intermediates, the diagram shows multiple peaks and valleys. Each peak represents a transition state for one elementary step, and each valley represents a relatively stable intermediate. The rate-determining step (RDS) corresponds to the highest energy barrier. The overall reaction rate is determined by the RDS. Intermediates are real chemical species (unlike transition states) that can sometimes be detected experimentally.
Some reactions can produce different products under different conditions. Kinetic product: formed via the lower activation energy pathway (faster, less stable). Thermodynamic product: the more stable product with lower overall energy (slower, forms at higher temperature). At low temperature: kinetic product dominates (reaction stops at first accessible pathway). At high temperature: thermodynamic product dominates (system can overcome all barriers to reach most stable state).
A catalyst provides an alternative reaction pathway with a lower activation energy. It does NOT change the thermodynamics (ΔH is the same with or without catalyst). It does NOT change the equilibrium position (Keq is unchanged). It speeds up BOTH forward and reverse reactions equally. The catalyzed pathway may involve different intermediates and transition states, but reactants and products remain the same.
The Arrhenius equation k = A·exp(-Ea/RT) shows that reducing Ea exponentially increases k. For example, reducing Ea by 20 kJ/mol at 298K increases k by a factor of ~3000. This is why even small reductions in activation energy have dramatic effects on reaction rate. On the Arrhenius plot (ln k vs 1/T), a catalyst lowers the slope (-Ea/R) while the intercept (ln A) may also change due to the different mechanism.
Haber-Bosch process (N₂ + 3H₂ → 2NH₃): Fe catalyst reduces Ea from ~420 kJ/mol to ~150 kJ/mol, enabling ammonia synthesis at practical temperatures. Contact process for H₂SO₄: V₂O₅ catalyst oxidizes SO₂ to SO₃. Catalytic converters: Pt/Pd/Rh reduce NOₓ, CO, and unburned hydrocarbons in automotive exhaust. Cracking: zeolite catalysts break long hydrocarbons into useful shorter chains.
Enzymes are nature's catalysts, typically accelerating reactions by 10⁶ to 10¹² times. Carbonic anhydrase converts CO₂ + H₂O → HCO₃⁻ + H⁺ at 10⁶ reactions/second. DNA polymerase replicates DNA with high fidelity at ~1000 nucleotides/second. ATP synthase couples proton gradient to ATP synthesis with ~100 turnovers/second. Enzyme catalysis uses proximity, orientation, acid-base, covalent, and strain mechanisms.
Fuel cells: Pt catalysts facilitate H₂ oxidation and O₂ reduction with lower activation barriers. Photocatalytic water splitting: TiO₂-based catalysts use sunlight to split water into H₂ and O₂. Carbon capture: catalysts help convert captured CO₂ into useful chemicals. Battery technology: catalytic surfaces improve charge transfer kinetics in lithium-ion and hydrogen fuel cells.
Drug design often targets enzyme active sites to modulate reaction coordinates. Competitive inhibitors raise the effective activation energy. Transition state analogs mimic the TS geometry, binding more tightly than substrate (tight-binding inhibitors). Prodrug design: inactive compound metabolizes to active drug via a favorable energy pathway. Understanding reaction coordinates helps predict metabolic stability, drug-drug interactions, and optimize synthetic routes.