$E^{\circ}$ is a promise made under laboratory conditions: every solute at $1\ \mathrm{mol/L}$, every gas at one bar, 25 °C. A real cell leaves those conditions the instant it starts working — it eats its reactants. The Nernst equation is the correction that follows it out: $E = E^{\circ} - \dfrac{RT}{nF}\ln Q$.

Every letter, before any of them works. $E$ is the cell's ACTUAL potential right now, in volts — the thing a voltmeter reads. $E^{\circ}$ is its standard potential in volts, the number you built last lesson. $n$ is the moles of electrons the balanced cell reaction moves — a pure count, no unit. $Q$ is the reaction quotient: products over reactants at this instant, each concentration raised to its coefficient, and also unitless. $R = 8.314\ \mathrm{J/(mol \cdot K)}$ is the gas constant, $T$ the absolute temperature in kelvin, and $F = 96\,500\ \mathrm{C/mol}$ the Faraday constant. You are solving for $E$.

At 25 °C the constants collapse into one number every exam hall knows: $\dfrac{RT}{F}\ln 10 = 0.0592\ \mathrm{V}$, which turns the natural log into the log you can read off a scale — $E = E^{\circ} - \dfrac{0.0592}{n}\log Q$. Read aloud: E equals E-naught minus zero-point-oh-five-nine-two over n, log Q. One decade of $Q$ is worth $59.2\ \mathrm{mV}$ to a one-electron cell — and that is not trivia. It is exactly why a pH meter reads about 59 millivolts per pH unit.

Three sanity rails, free with every calculation. $Q < 1$ means reactant-rich, the cell is eager, and $E$ sits ABOVE $E^{\circ}$. $Q = 1$ is standard, and $E = E^{\circ}$ exactly. And $Q = K$ is equilibrium, where $E = 0$ — which is all a dead battery ever was.