Illustrated Guide to Home Chemistry Experiments

294 DIY Science: Illustrated Guide to Home Chemistry Experiments

In the first lab of this chapter, you may have

wondered how we knew that the electrolysis

of water would require at least 1.23V. A glance

at the standard reduction potentials listed in

Table 16-3 provides the answer. (Table 16-3

lists standard reduction potentials for only

a handful of half-reactions. Comprehensive

tables that list reduction potentials for

thousands of half-reactions are available

online and in printed reference works.) Of

the half-reactions involved in the electrolysis

of water, the reduction of oxygen gas and

hydrogen ions (protons) to water:

o 2 (g) + 4 H+(aq) + 4 e– → 2 H 2 o(l)

has the greatest reduction potential, +1.23V. Of course,
oxidation and reduction are two sides of the same coin, so the
reverse reaction is equally valid:

2 H 2 ol) ( → o 2 (g) + 4 H+(aq) + 4 e–

So, looking at the oxidation reaction rather than the reduction
reaction, the oxidation of water to oxygen gas and hydrogen ions
has the greatest oxidation potential, which means that external
current must be supplied at a minimum voltage of –1.23V for
the half-reaction to occur.

Half-reactions that appear near the top of the table—those
with high negative values—represent strong reducers.
Those listed near the bottom of the table—those with high
positive values—represent strong oxidizers. It’s important to
understand that a species in isolation cannot be said to be a
reducer or an oxidizer, because the species with which it reacts
determines whether the first species functions as a reducer or
an oxidizer.

For example, consider the pairs of ions and elements, Zn2+/
Zn, Fe2+/Fe, and Cu2+/Cu, which have reduction potentials
of –0.76V, –0.44V and +0.34V, respectively. If you place solid
iron in a solution of Cu2+ ions, those ions are reduced to solid
copper while the solid iron is oxidized to Fe2+ ions. With this
combination of species, iron serves as the reducer and copper
as the oxidizer. Conversely, if you place solid zinc in a solution
of Fe2+ ions, those ions are reduced to solid iron while the solid

LABORATORY 1 6.3:

mEASURE ELECTRodE poTENTIALS

Half-Reaction E^0 (v)

Li+(aq) + e– → Li(s) –3.05

K+(aq) + e– → K(s) –2.93

Ba2+(aq) + 2 e– → Ba(s) –2.90

Ca2+(aq) + 2 e– → Ca(s) –2.87

Na+(aq) + e– → Na(s) –2.71

Mg2+(aq) + 2 e– → Mg(s) –2.37

Al3+(aq) + 3 e– → Al(s) –1.66

2 H 2 O(l) + 2 e– → H 2 (g) + 2 OH-(aq) –0.83

Zn2+(aq) + 2 e– → Zn(s) –0.76

Cr3+(aq) + 3 e– → Cr(s) –0.73

Fe2+(aq) + 2 e– → Fe(s) –0.44

Cd2+(aq) + 2 e– → Cd(s) –0.40

Co2+(aq) + 2 e– → Co(s) –0.28

Ni2+(aq) + 2 e– → Ni(s) –0.25

Sn2+(aq) + 2 e– → Sn(s) –0.14

Pb2+(aq) + 2 e– → Pb(s) –0.13

2 H+(aq) + 2 e– → H 2 (g) 0.00

Cu2+(aq) + 2 e– → Cu(s) +0.34

O 2 (g) + 2 H 2 O(l) + 4 e– → 4 OH–(aq) +0.40

Ag+(aq) + e– → Ag(s) +0.80

Hg2+(aq) + 2 e– → Hg(l) +0.85

O 2 (g) + 4 H+(aq) + 4 e– → 2 H 2 O(l) +1.23

Au3+(aq) + 3 e– → Au(s) +1.50

H 2 O 2 (aq) + 2 H+(aq) + 2 e– → 2 H 2 O(l) +1.78

F 2 (g) + 2 e– → 2 F-(aq) +2.87

TABLE 16-3: Standard reduction potentials at 25°C.