standard hydrogen electrode
/ S-H-E /
You cannot measure the absolute electrical 'push' of a single electrode any more than you can measure how high a place is without agreeing on sea level. A voltmeter always reads the difference between two electrodes. So chemists chose one electrode to be the universal sea level, declared its potential to be exactly zero, and measured everything else against it. That reference is the standard hydrogen electrode.
It is a deliberately reproducible setup: a piece of inert platinum, coated with finely divided platinum black to speed things up, dipped in a solution of acid at exactly 1 molar H+ activity, with hydrogen gas at 1 bar bubbling over it, all at 298 K. The half-reaction it sustains is 2 H+ + 2 e- <-> H2. By international convention its standard electrode potential is defined as 0.000 V at all temperatures. Connect any other half-cell to it under standard conditions, read the voltmeter, and that reading is the other electrode's standard potential.
Every number in a table of standard reduction potentials, and therefore the whole electrochemical series, is anchored to this one arbitrary but fixed choice. The honest point is that the zero is a convention, not a real absence of potential — the true single-electrode potential is unmeasurable, and a different choice of reference (such as the silver/silver chloride electrode used in practice, because the SHE is fiddly) would simply shift every number by a constant without changing any predicted cell voltage.
Couple a zinc electrode (Zn2+/Zn) to the SHE under standard conditions and the voltmeter reads 0.76 V with zinc negative, so the standard reduction potential of Zn2+/Zn is -0.76 V.
Every standard potential is just a voltage measured against this 0 V reference.
The 0 V is assigned by definition, not measured. In the lab the SHE is awkward to maintain, so practical work uses secondary references like Ag/AgCl or the calomel electrode and converts back to the SHE scale afterward.