Redox Chemistry in Cells—Why Can Electrons Flow Opposite to Standard Potentials? Sometimes cellular redox reactions transfer electrons from a reductant with a higher standard redox potential (E°′) to an oxidant with a lower E°′. Why can this occur inside cells?

Introduction / Context:
Standard redox potentials (E°′) provide a useful reference for the directionality of electron flow. However, the actual driving force in cells depends on the effective potentials, which are functions of concentrations, pH, and coupling to other processes. This explains why some in vivo electron transfers seem to contradict standard tables.

Given Data / Assumptions:

• E°′ values are measured at defined concentrations, temperature, and pH.
• Cellular metabolite levels can differ by orders of magnitude from standard conditions.
• Reactions may be enzyme-coupled or part of multi-step pathways.

Concept / Approach:
The Nernst equation adjusts E based on actual reactant/product activities: E = E°′ + (RT/nF) * ln([oxidized]/[reduced]). If the intracellular ratio strongly favors product formation or if reactions are coupled (e.g., to proton gradients or ATP hydrolysis elsewhere), the net ΔG can become negative, allowing electron flow that would look unfavorable by E°′ alone.

Step-by-Step Solution:

Verification / Alternative check:
Experimental measurements of metabolite ratios (e.g., NADH/NAD+) show large deviations from standard conditions, explaining directionality in central metabolism.

Why Other Options Are Wrong:

• (b) Oxidation–reduction rules apply universally.
• (c) Not always; actual potentials may reverse the driving force.
• (e) Enzymes lower activation energy but do not create energy; they cannot violate thermodynamics.

Common Pitfalls:
Assuming E°′ tables dictate absolute direction in vivo; they are guides that must be corrected for physiological conditions and coupling.

Final Answer:
Redox potentials are defined under standard conditions, while cellular conditions are not standard