2.7.2 · D5Redox & Electrochemistry (Intro)
Question bank — Standard electrode potentials — SHE reference, electrochemical series
True or false — justify
True or false: The SHE has a standard potential of exactly because that is its real, physically measured value.
False. It is by convention — an arbitrary chosen zero, like calling sea level "zero altitude." No single electrode's absolute potential is measurable, so we anchor the whole scale to hydrogen.
True or false: A more positive always means a faster reaction.
False. measures thermodynamic driving force (how far, how spontaneous), not kinetics (how fast). A large positive can still pair with a reaction that crawls due to a high activation barrier.
True or false: means zinc cannot be reduced at all.
False. A negative means reduction is non-spontaneous relative to SHE, not impossible. Couple Zn to something with an even lower (e.g. at ) and does get reduced.
True or false: If you double every coefficient in a half-reaction, its doubles too.
False. is an intensive property — a potential per electron's worth of driving force. Scaling the equation scales and together, leaving unchanged.
True or false: The species at the bottom of the reduction-potential table (most positive ) is the strongest oxidizing agent.
True. Most positive = strongest pull on electrons = most readily reduced = strongest oxidizing agent (it oxidizes something else while being reduced itself). That is why at sits at the extreme.
True or false: A metal high in the series (very negative ) is a strong oxidizing agent.
False — it is a strong reducing agent. Very negative means the reduced form (the metal) gives up electrons easily, so it reduces other things while being oxidized itself.
True or false: uses the anode's oxidation potential.
False. Both values in that formula are reduction potentials as tabulated. The subtraction itself accounts for the anode running in reverse — you do not flip the sign of the anode value first and subtract.
Spot the error
"For the Fe³⁺/I⁻ cell I reversed the iodine half-reaction, so I flipped its to , then wrote ." Find the mistake.
The reader flipped the sign and used the subtraction formula — double-counting. Either use with tabulated values, or add oxidationreduction potentials once flipped. Never both.
"Zn²⁺/Zn is and involves 2 electrons, so to compare it with Ag⁺/Ag I should scale it to per-electron: ." Find the mistake.
is already per-electron in the intensive sense — you never divide it by . Half-cells are compared by their tabulated directly, regardless of how many electrons the equation shows.
"The cell gives , so it is non-spontaneous." Find the mistake.
Cathode and anode were swapped. Cu²⁺/Cu (higher ) is the cathode: , which is spontaneous — exactly why the Daniell cell works.
"Since and is huge, a tiny still gives a large , so any positive guarantees a strongly spontaneous reaction." Find the flaw.
The sign is what decides spontaneity, and being large only scales magnitude — but a tiny positive gives a small-magnitude , so it is only weakly spontaneous and easily reversed by concentration (2.7.03-Nernst-equation-and-concentration-effects).
"To measure Cu's I put both the Cu half-cell and a Zn half-cell together and read the meter — that gives me directly." Find the mistake.
A single measurement gives only the difference between two half-cells. To pin an absolute value on Cu you must couple it to the SHE (defined as ), otherwise you have one equation with two unknowns.
"F₂ has the biggest positive , so fluorine gas is the strongest reducing agent." Find the mistake.
Reversed. Biggest positive = strongest oxidizing agent. is the strongest oxidizer in the common table; the strongest reducers sit at the top with very negative (e.g. Li).
Why questions
Why can't we measure the absolute potential of a single electrode?
A potential difference needs two points. A lone electrode has no second terminal to measure against, so every real reading is a difference between two half-cells — hence the need for a chosen zero reference.
Why is platinum used in the SHE rather than a reactive metal?
Pt is inert (it does not dissolve or take part chemically) yet catalyses the H⁺/H₂ interconversion and conducts electrons. It provides a clean surface for the half-reaction without adding its own competing potential.
Why do we insist on 1 M H⁺, 1 bar H₂, and 298 K for the "standard" hydrogen electrode?
Potential depends on concentration, pressure and temperature (via the Nernst relation). Fixing these to reference values makes a reproducible constant that anyone can reproduce and tabulate.
Why does a negative correspond to a good reducing agent?
A negative reduction potential means the forward (gain-electron) direction is unfavourable, so the species prefers to run backwards — giving up electrons, i.e. acting as a reducer.
Why does reversing a half-reaction change the sign of its but leave the same?
Reversing the direction reverses the sign of ; since and are unchanged, only the sign of flips. The magnitude of the driving force is symmetric.
Why does the electrochemical series predict whether but not always how much product forms?
fixes the sign of and hence direction and the equilibrium constant (2.7.05-Gibs-free-energy-and-equilibrium-constant), but real yields also depend on kinetics, side reactions, and non-standard concentrations shifting the potential.
Why is the electrochemical series ordered by reduction potentials specifically, not oxidation potentials?
A single, consistent convention avoids sign chaos. IUPAC fixed reduction potentials as the standard; oxidation potentials are simply their negatives, so one table suffices.
Edge cases
What is if you connect the SHE to another SHE?
Exactly . Both half-cells are identical, so cathode and anode potentials cancel — there is no driving force and no net reaction.
A predicted reaction has exactly. What happens?
, so the system is at equilibrium under standard conditions — no net forward or reverse drive. Any real push comes only from non-standard concentrations.
Under non-standard conditions, can a reaction with negative still become spontaneous?
Yes. Concentration changes shift the actual away from via the Nernst equation (2.7.03-Nernst-equation-and-concentration-effects); a strongly favourable concentration ratio can flip the sign of even when .
If two half-reactions have identical , which is the cathode?
Neither is favoured — , so the pair sits at standard-state equilibrium with no spontaneous direction. Labelling is arbitrary until concentrations break the tie.
Water can be oxidised or reduced. Why does this limit which species you can actually study in aqueous SHE-based cells?
Any half-reaction far outside water's stability window ( very high or very low) will instead oxidise or reduce water itself, so the tabulated aqueous for extreme couples (like Li) is thermodynamic, not directly observable in plain water.
For the reduced form is itself an ion, not a metal. Does the series still apply?
Yes. The series lists any redox couple, not only metal/ion pairs. at compares to every other couple exactly the same way — higher is the better oxidiser.
Recall One-line self-test
Cover and recall: what three things must you fix for "standard," what the SHE's potential is and why, and the one formula using tabulated reduction values. The three standards ::: 1 M ions, 1 bar gas, 298 K. SHE potential and reason ::: , fixed by convention as the reference zero. Sign rule for spontaneity ::: spontaneous.