182 DIY Science: Illustrated Guide to Home Chemistry Experiments
EvdRE y Ay REdox REACTIoNS
Redox reactions, including some that we literally couldn’t
live without, are common in everyday life. Here are just a
few examples:- Metals are usually mined in the form of an oxide,
carbonate, or other salt of the metal. Refining
metals—whether by smelting, electrolysis, or some
other method—uses redox reactions; specifically,
reduction of the metal in a positive oxidation state to
the neutral metallic form. - Water purification uses redox reactions to oxidize
colored, bad-tasting, or otherwise objectionable
impurities to forms that are safe and acceptable. - Bleaches used in laundry, papermaking, and other
processes all depend upon redox reactions to oxidize
colored stains and impurities to colorless compounds. - Photographic film works by a redox reaction initiated
by light. When photons strike the tiny grains of
silver bromide and other silver halides present in
the emulsion, they cause silver ions to be reduced
to microscopic specks of metallic silver. During
development, those tiny specks of silver catalyze
reduction of the silver halide grains surrounding them.
After washing away the unreduced silver, what remains
are the tiny black grains that make the picture. - Corrosion and rusting are examples of redox
reactions that are familiar to everyone. In particular,
unprotected iron and steel are prone to react with
atmospheric oxygen in the presence of moisture to
form iron oxides, familiarly called rust. Many forms
of protection against corrosion take advantage
of redox reactions as well. For example, plating
iron with zinc (called galvanizing) protects the
underlying iron, because it will not oxidize while in
contact with the more active metal.
Many chemical reactions do not involve any change in oxidation
state. For example, when we reacted aqueous solutions of
copper(II) sulfate and sodium hydroxide to form a precipitate of
copper(II) hydroxide and a solution of sodium sulfate, no changes
in oxidation state occurred:
CuSo 4 (aq) + 2 NaoH(aq) → Cu(oH) 2 (s) + Na 2 So 4 (aq)
or, showing the oxidation states:
Cu2+(aq) + So 4 2–(aq) + 2 Na+(aq) + 2 oH–(aq)
→ Cu2+(oH–) 2 + 2 Na+(aq) + So 4 2–(aq)
On both sides of the equation, copper is in the +2 oxidation state,
sulfate –2, sodium +1, and hydroxide –1.
Oxidation state 0 is the neutral form of any element, such as (at
standard conditions) aluminum in its metallic form or chlorine in its
gaseous form. An element in oxidation state 0 possesses the same
number of negatively charged electrons in its shell as the number
of positively charged protons in its nucleus, leaving a net charge of 0.
If a chlorine atom at oxidation state 0 gains an electron, that
chlorine atom becomes a chlorine ion with a charge of –1, and an
oxidation state of –1. Conversely, if a chlorine atom at oxidation
state 0 loses an electron, that chlorine atom becomes a chlorine
ion with a charge of +1, and an oxidation state of +1. With the
exception of the noble gases—helium, neon, argon, krypton, xenon,
and radon, all in column 18 of the periodic table—all elements are
commonly found in at least one oxidation state other than 0. Some
elements have only one or two known non zero oxidation states.
Other elements are commonly found in many oxidation states. For
example, in its natural form and in its many compounds, carbon
can be found in oxidation states –4, –3, –2, –1, 0, +1, +2, +3, and +4.
Oxidation and reduction also apply to atoms that are in oxidation
states other than 0. For example, iron is found in three oxidation
states. Metallic iron is in oxidation state 0. Metallic iron can be
oxidized to the Fe2+ ion, also called the ferrous ion or the iron(II)
ion. The Fe2+ ion can be further oxidized to the Fe3+ ion—also
called the ferric ion or the iron(III) ion—or it can be reduced to
metallic iron. The Fe3+ ion can be reduced to the Fe2+ ion by gaining
one electron, or all the way to Fe^0 (iron metal) by gaining three
electrons.
Ions of a particular element in different oxidation states can have
very different chemical and physical properties. For example,
solutions that contain vanadium ions in different oxidation states
show a rainbow of colors. Solutions of V5+ ions are yellow; those
of V4+ are blue; V3+, green; and V2+, violet. Solutions of ions in
different oxidations states of chromium, manganese, and other
elements show similar color differences.
It’s important to understand that in a redox reaction oxidation
and reduction are two sides of the same coin. If one atom is
oxidized during the reaction, another must be reduced, and vice
versa. Furthermore, the total number of electrons lost by the
atom or atoms being oxidized must equal the total number of
electrons gained by the atom or atoms that are being reduced.
Charge must be conserved. For example, when aluminum reacts
with the H+ ions from hydrochloric acid to form Al3+ ions, the
three electrons lost when one aluminum atom is oxidized reduce
three H+ ions to H^0 atoms.The laboratory sessions in this chapter explore various aspects of
redox reactions.