Tin(II) chloride, also known as stannous chloride, is a white crystalline solid
with the formula SnCl2. It forms a stable dihydrate, but aqueous solutions
tend to undergo hydrolysis, particularly if hot. SnCl2 is widely used as a
reducing agent, and in electrolytic baths for tin-plating. Tin(II) chloride
should not be confused with the other chloride of tin; tin(IV) chloride or
stannic chloride. Chemical structure
SnCl2 has a lone pair of electrons, such that the molecule in the gas phase is
bent. In the solid state, crystalline SnCl2 forms chains linked via chloride
bridges as shown. The dihydrate is also three-coordinate, with one water
coordinated on to the tin, and a second water coordinated to the first. The main
part of the molecule stacks into double layers in the crystal lattice, with the
"second" water sandwiched between the layers.
Chemical properties Tin(II) chloride can dissolve in less
than its own mass of water without apparent decomposition, but as the
solution is diluted hydrolysis occurs to form an insoluble basic salt:
SnCl2 + H2O Sn(OH)Cl + HCl Therefore if clear solutions of tin(II)
chloride are to be used, it must be dissolved in hydrochloric acid to
maintain the equilibrium towards the left-hand side. Solutions of SnCl2 are
also unstable towards oxidation by the air:
6 SnCl2 + O2 + 2 H2O → 2 SnCl4 + 4 Sn(OH)Cl
This can be prevented by storing the solution over lumps of tin metal.
There are many such cases where tin(II) chloride acts as a reducing agent,
reducing silver and gold salts to the metal, and iron(III) salts to iron(II),
for example: SnCl2 + 2 FeCl3 → SnCl4 + 2 FeCl2
It also reduces copper(II) to copper(I). Solutions of tin(II) chloride can also
serve simply as a source of Sn2+ ions, which can form other tin(II) compounds
via precipitation reactions. For example, reaction with sodium sulfide
produces the brown/black tin(II) sulfide:
SnCl2 + Na2S → SnS + 2 NaCl If alkali is added to a solution of
SnCl2, a white precipitate of hydrated tin(II) oxide forms initially; this then
dissolves in excess base to form a stannite salt such as sodium stannite:
SnCl2(aq) + 2 NaOH → SnO·H2O + 2 NaCl SnO·H2O + NaOH → NaSn(OH)3
Anhydrous SnCl2 can be used to make a variety of interesting tin(II) compounds
in non-aqueous solvents. For example, the lithium salt of
4-methyl-2,6-di-tert-butylphenol reacts with SnCl2 in THF to give the yellow
linear two-coordinate compound Sn(OAr)2. Tin(II) chloride also behaves as a Lewis
acid, forming complexes with ligands such as chloride ion, for example:
SnCl2 + CsCl → CsSnCl3 Most of these complexes are pyramidal,
and since complexes such as SnCl3 have a full octet, there is little tendency to
add more than one ligand. The lone pair of electrons in such complexes is
available for bonding, however, and therefore the complex itself can act as
a Lewis base or ligand. This seen in the ferrocene-related product of the
following reaction : SnCl2 + Fe(η5-C5H5)(CO)2HgCl →
Fe(η5-C5H5)(CO)2SnCl3 + Hg SnCl2 can be used to make a variety of
such compounds containing metal-metal bonds. For example, the reaction with
dicobalt octacarbonyl: SnCl2 + Co2(CO)8 →4Co-(SnCl2)-Co(CO)4
Preparation Anhydrous SnCl2 is prepared by the
action of dry hydrogen chloride gas on tin metal. The dihydrate is made by a
similar reaction, using hydrochloric acid:
Sn + 2 HCl → SnCl2 + H The water then carefully evaporated from
the acidic solution to produce crystals of SnCl2·2H2O. This dihydrate can be
dehydrated to anhydrous using acetic anhydride.
Uses A solution of tin(II) chloride
containing a little hydrochloric acid is used for the tin-plating of steel, in
order to make tin cans. An electric potential is applied, and tin metal is
formed at the cathode via electrolysis. Tin(II) chloride is used as a mordant in
textile dyeing because it gives brighter colours with some dyes e.g. cochineal.
This mordant has also been used alone to increase the weight of silk.
It is used as a catalyst in the production of the plastic polylactic
acid. It also finds a use as a catalyst
between acetone and hydrogen peroxide to form the tetrameric form of acetone
peroxide. Tin(II) chloride also finds wide use as
a reducing agent. This is seen in its use for silvering mirrors, where silver
metal is deposited on the glass: Sn2+ + 2 Ag+ → Sn4+ + 2 Ag
A related reduction was traditionally used as an analytical test for Hg2+(aq).
For example, if SnCl2 is added dropwise into a solution of mercury(II) chloride,
a white precipitate of mercury(I) chloride is first formed; as more SnCl2
is added this turns black as metallic mercury is formed. Stannous chloride can
be used to test for the presence of gold compounds. SnCl2 turns bright purple in
the presence of gold. When mercury is analyzed using atomic
absorption spectroscopy, a cold vapor method must be used, and tin chloride is
typically used as the reductant. In organic chemistry, SnCl2 is mainly
used in the Stephen reduction, whereby a nitrile is reduced to an imine which is
easily hydrolysed to an aldehyde. The reaction usually works best with
aromatic nitriles Aryl-CN. A related reaction starts with an amide, which is
treated with PCl5 to form the imidoyl chloride salt.
The Stephen reduction is less used today, because it has been mostly
superseded by diisobutylaluminium hydride reduction.
Additionally, SnCl2 is used to selectively reduce aromatic nitro groups
to anilines. SnCl2 also reduces quinones to
hydroquinones. Stannous chloride is also added as a
food additive with E number E512 to some canned and bottled foods, where it
serves as a color-retention agent and antioxidant.
SnCl2 is used in Radionuclide angiography to reduce the radioactive
agent technetium-99m-pertechnetate to assist in binding to blood cells.
Finally, aqueous Stannous Chloride is used by many precious metals refining
hobbyists as an indicator of Gold and Platinum group metals in solutions.
Notes References
N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed.,
Butterworth-Heinemann, Oxford, UK, 1997. Handbook of Chemistry and Physics, 71st
edition, CRC Press, Ann Arbor, Michigan, 1990.
The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
A. F. Wells, 'Structural Inorganic Chemistry, 5th ed., Oxford University
Press, Oxford, UK, 1984. J. March, Advanced Organic Chemistry,
4th ed., p. 723, Wiley, New York, 1992.