Illustrated Guide to Home Chemistry Experiments

Chapter 16 Laboratory: Electrochemistry 287

LABORATORY 1 6.1:

pRodUCE HydRoGEN ANd oxyGEN By ELECTRoLySIS of wATER

Electrolysis is the process of forcing a redox

reaction to occur by supplying an external

electric current. For example, electrolysis can

be used to plate an iron object with chromium

by reducing chromium ions from a solution,

replacing them with iron ions. Ordinarily, this

reaction could not occur, because chromium

is higher in the activity series of metals than

iron. The application of an external electric

current overcomes the difference in activity

(potentials) between these two metals,

allowing the chromium ions to be reduced to

chromium metal.

RIREEqU d EqUIpmENT ANd SUppLIES

£ goggles, gloves, and protective clothing

£ balance and weighing papers

£ ring stand

£ burette clamp

£ test tube, with stopper (4)

£ test tube rack

£ graduated cylinder, 10 mL

£ beaker, 600 mL

£ 9 v transistor battery

£ marking pen

£ matches, strike–anywhere

£ rubber band (2)

£ magnesium sulfate heptahydrate (~15 g)

£ distilled or deionized water

SBSTITUTIU oNS ANd modIfICATIoNS

• You may substitute any container of similar size for
  the beaker.


• You may substitute wood splints or toothpicks for the
  strike–anywhere matches.


• Magnesium sulfate heptahydrate is sold in drugstores
  as Epsom salts. You may also substitute sodium
  bicarbonate (baking soda) for the magnesium sulfate.

Electrolysis is commonly used industrially and in laboratories
to produce pure hydrogen and oxygen by splitting otherwise
stable water molecules into their component gases, a process
that requires providing an external electrical current at 1.23V or
higher. This voltage is applied to two inert electrodes immersed
in a dilute aqueous solution of an ionic compound. (Pure
water cannot be used, because it is a very poor conductor of
electricity.) In an electrolytic cell (as opposed to a galvanic
cell, where the direction is reversed), the negative electrode,
or cathode, provides the electrons needed to reduce H+ ions
to hydrogen gas. Conversely, the positive electrode, or anode,
accepts the electrons needed to oxidize O– ions to oxygen gas.

We can balance this redox reaction using half-reactions. To do
so, we must identify which species is being oxidized and which
reduced by looking at the change in oxidation states:

2 H+1 2 o–2(l) → 2 H^02 (g) + o^02 (g)

From this balanced equation, we see that hydrogen is being
reduced and oxygen is being oxidized, which allows us to set up
the following two half-reactions:

(oxidation) 2 o–2 → o 2 + 4 e–

eduction) 2 H(r + + 2 e– → H 2

Doubling the reduction half-reaction to put the same number of
electrons on each side—everything, including electrons, must
balance—and then adding the two half-reactions gives us the final
balanced reaction:

[ 2 o–2 → o 2 + 4 e–] + [4 H+ + 4 e– → 2 H 2 ]

= 2 H 2 o2 H → 2 + o 2

In this laboratory session, we’ll use electrolysis to produce

hydrogen and oxygen from water.