Electrochemistry

A chemical reaction shuffles electrons around. If we make those electrons flow through a wire instead of jumping directly, that flow is electricity. So a reaction can power a torch — or, run in reverse, electricity can force a reaction to happen.

Galvanic = a battery rolling downhill on its own and giving you energy. Electrolytic = you pushing the ball back uphill using an external power supply. The classic galvanic example is the Daniell cell:

Remember “An–Ox, Red–Cat”: Anode → Oxidation, Reduction → Cathode. This is true for every cell, galvanic or electrolytic.

As zinc dissolves, its side builds up extra positive charge; the copper side builds up extra negative. The salt bridge quietly leaks ions across to cancel this build-up — like a referee passing tokens between two tables so neither runs out of balance. Without it, the cell stalls within seconds.

You can’t measure the height of a single point — only its height above sea level. SHE is chemistry’s “sea level”: we declare E° = 0 for it and measure every other electrode relative to it (SHE conditions: 1 M H⁺, 1 bar H₂ gas, 298 K).

Use the convention “right minus left”: in standard cell notation the cathode is written on the right. Both values are taken as reduction potentials — no flipping signs needed.

Think of E° as a ranking of how badly each species wants to grab electrons. F₂ sits at the very top (greedy oxidiser, +2.87 V); lithium sits at the very bottom (gives electrons away, strongest reducing agent).

E° is the “list price” under perfect lab conditions. The Nernst equation is the discount-or-surcharge that adjusts the price for how much reactant and product you actually have (Q is the reaction quotient — products over reactants). Here n = electrons transferred, F = 96487 C/mol.

Memorise the magic number 0.059 (volts). At room temperature, 2.303RTF = 0.059, so the whole equation becomes a quick mental sum.

A flat battery isn’t “empty” — it has simply reached chemical balance, where forward and backward rates match. Setting E_cell = 0 in the Nernst equation hands you the equilibrium constant directly. Big E°_cell ⇒ huge K_c ⇒ reaction goes almost to completion.

More electrons moved (n) at a higher voltage (E°) means more useful work extracted. The minus sign means a positive emf gives a negative ΔG — exactly the signature of a spontaneous reaction.

A wide, short pipe lets water through easily; a thin, long one resists. Conductivity strips out the pipe’s shape and tells you how well a 1 cm cube of the solution itself carries current. The shape factor G* = l/A is the cell constant.

Conductivity alone is unfair to compare — a strong solution has more ions just because it’s crowded. Molar conductivity asks the fairer question: per mole of salt, how well does it conduct?

Dilute the solution and you have fewer ions in each cm³, so κ (current per volume) drops. But you measure Λ_m per mole, and now those ions have more room to roam (and weak ones split up more), so Λ_m climbs.

A relay team’s total speed is just the sum of each runner’s speed. At infinite dilution the ions are so far apart they ignore each other, so the solution’s conductivity is simply each ion’s personal contribution added up (ν = how many of each ion the formula gives).

A strong electrolyte (like KCl) is already fully split into ions; diluting only spreads them out a little, so the graph against √c is a tidy straight line you can extend back to zero concentration.

A weak electrolyte (like acetic acid) is barely split at normal concentration; only near infinite dilution does it fully ionise, and the curve rockets up too sharply to extrapolate. So we assemble its Λ°_m from strong-electrolyte values that we can measure.

If a weak acid conducts only one-tenth as well as its fully-split ideal, then about one-tenth of it has dissociated, so α = 0.1. Plug that into the K_a formula and you’ve measured the acid’s strength with a conductivity meter.

Electroplating is like paying with electrons: the more current (I) you push, for the more time (t), the more metal you buy. Z (the electrochemical equivalent) is just the “price per coulomb”.

Think of F as the cost of one mole of electrons. A 1+ ion needs 1 mole of electrons (1 F) per mole of metal; a 2+ ion needs 2 F; a 3+ ion needs 3 F.

Give every cell the exact same “electron budget”, and each one spends it according to its own price tag. Equal coulombs through silver and copper cells deposit them in the ratio of their equivalent weights (Ag ≈ 108, Cu ≈ 31.75).

In water there’s always a competition: should the metal ion react, or should water itself? Whoever is “easier” (electrode potential + a real-world nudge called overpotential) wins.

Primary = use-and-throw (dry cell / Leclanché, mercury cell). Secondary = refillable (lead-storage battery, nickel–cadmium) — you push the reaction backwards to recharge.

It’s the workhorse car battery. Discharging coats both plates with PbSO₄; charging (from the alternator) strips it back off, ready to go again.

Unlike a battery (a sealed tank of chemicals), a fuel cell is a generator you keep feeding. The H₂–O₂ cell’s only “exhaust” is pure water — very efficient and pollution-free, which is why spacecraft used them.

A wet iron surface becomes a patchwork of tiny galvanic cells — some spots give up electrons (anode) and dissolve, oxygen is reduced elsewhere, and the iron slowly turns to rust, Fe₂O₃·xH₂O.