Second group of anions. Complexation
equilibrium
Among anions II analytical group (Cl-, Br-, I-,
S2-, SCN-, [Fe(CN)6]4-, [Fe(CN)6]3-,
CN-, BrÎ3–,
IO3-, Cl–) is such which play the important
physiological role, and such anions are a part of drugs or are applied in the
analysis of drugs. Therefore it is necessary to know chemical-analytical
properties of the II analytical group of anions.
The Cl-, Br-, I-,
S2-, SCN-, [Fe(CN)6]4-, [Fe(CN)6]3-,
CN-, BrÎ3–,
IO3-, ClΖ anions are
anions of ˛˛
analytical group. These anions with Ag+ ions in weak nitric acid
solutions give precipitate which is insoluble in diluted (2 mol/L) HNO3.
Therefore, a group reagent ˛˛ groups anions are solution of AgNO3
in the presence of diluted HNO3.
Ions S2- are oxidized by nitric acid to
sulphur (or even to SO42–), therefore group reagent on
anions of II analytical group is acidic solution of Silver nitrate.
Barium Salts of anions of ˛˛ groups are
dissolved in water. Therefore, unlike anions of ˛ groups, anions
of ˛˛
groups don’t react with solution of BaCl2. In aquaeous solutions of
anions of ˛˛
groups are colourless (except [Fe(CN)6]4- and [Fe(CN)6]3-).
Chloride
|
Chloride |
|
|
|
|
|
Chloride |
|
|
Identifiers |
|
|
3587171 |
|
|
14910 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
Cl− |
|
|
35.453 g mol-1 |
|
|
Thermochemistry |
|
|
Std enthalpy of |
−167 kJ·mol−1 |
|
Standard molar |
153.36 J K-1
mol-1 |
|
Related
compounds |
|
|
Other anions |
|
|
Except where noted otherwise, data are given for materials in their standard state (at 25 °C,
100 kPa) |
|
The chloride ion is formed when the element
chlorine, a
halogen, gains an electron to form an anion (negatively
charged ion) Cl−. The salts
of hydrochloric acid contain chloride ions and can
also be called chlorides. The chloride ion, and its salts such as sodium
chloride, are very soluble in water. It is an essential electrolyte
located in all body fluids responsible for maintaining acid/base balance,
transmitting nerve impulses and regulating fluid in and out of
cells.
The word chloridism can also form part of the name
of chemical compounds in which one or more chlorine atoms are covalently
bonded. For example, methyl chloride, more commonly called chloromethane,
(CH3Cl) is an organic covalently bonded compound, which does not
contain a chloride ion.
The presence of chlorides, e.g. in seawater,
significantly aggravates the conditions for pitting
corrosion of most metals (including stainless steels and high-alloyed
materials) by enhancing the formation and growth of the pits through an
autocatalytic process.
Chloride is used to form salts that can preserve food
such as sodium chloride. Other salts such as calcium
chloride, magnesium chloride, potassium chloride have varied uses ranging from
medical treatments to cement formation.
An example is table salt, which is sodium
chloride with the chemical formula NaCl. In water,
it dissociates into Na+ and Cl− ions.
Examples of inorganic covalently bonded
chlorides that are used as reactants are:
·
phosphorus trichloride, phosphorus pentachloride, and thionyl
chloride, all three of which reactive chlorinating reagents that
have been used in a laboratory
·
disulfur dichloride (S2Cl2),
used for vulcanization of rubber.
A chloride ion is also the prosthetic group present in
the amylase
enzyme.
Another example is calcium
chloride with the chemical formula CaCl2. Calcium
chloride is a salt that is marketed in pellet form for
removing dampness from rooms. Calcium chloride is also used for maintaining
unpaved roads and for fortifying roadbases for new construction. In addition,
Calcium chloride is widely used as a De-icer since it
is effective in lowering the melting
point when applied to ice.
In the petroleum industry, the chlorides are a closely
monitored constituent of the mud
system. An increase of the chlorides in the mud system may be an indication
of drilling into a high-pressure saltwater formation. Its increase can also
indicate the poor quality of a target sand.
Chloride is also a useful and reliable chemical indicator
of river / groundwater fecal contamination, as chloride is a non-reactive
solute and ubiquitous to sewage & potable water. Many water regulating
companies around the world utilize chloride to check the contamination levels
of the rivers and potable water sources.
Chloride is a chemical the human body needs for metabolism
(the process of turning food into energy). It also helps keep the body's
acid-base balance. The amount of chloride in the blood is carefully controlled
by the kidneys.
Characteristic
reactions Cl- ions.
Potassium bichromate (pharmacopeia’s
reaction) oxidises chloride-ion in the medium
of sulphatic acid to chromil chloride CrÎ2Cl2:
Ńr2O72-+ 4Cl- + 6H+
→ 2CrÎ2Cl2 + 3H2O.
Chromil chloride is possible to detect by a
smell or on change of colouring of a paper (moistened by a solution diphenilcarbazid).
CrÎ2Cl2 + 4H+ + 4e C Cr2+ + 2H2O + 2Cl-.
diphenilcarbazid diphenilcarbazon
Reaction
performance. To
4-5 drops of investigated solution add 2-3 crystals of Potassium bichromate and
4-5 drops of the concentrated sulphatic acid. To the top of test tube bring the
filtering paper moistened by solution diphenilcarbazid. If there are chloride-ions
the paper is painted in violet-red colour.
Silver nitrate
with Cl- ions forms white precipitate AgCl which is not dissolved in
the diluted acids, but it is well dissolved in aqueous solution of NH3:
Ag+ +
Cl- = AgCl¯;
AgCl + 2NH3
= [Ag(NH3)2]+ + Cl–.
Reaction
performance. To
1-2 drops of an investigated solution add 2-3 drops of 2 mol/L nitric acid and
1-2 drops of AgNO3 solution. If there are chloride-ions the white precipitate
forms, it is soluble in aqueous solution of NH3 or 12 % solution of (NH4)2CO3.
To received solution add some drops of nitric acid. The white precipitate of
AgCl forms again.
Calcium chloride
|
Calcium Chloride |
|
|
|
|
|
|
|
|
Calcium chloride |
|
|
Other names Calcium(II) chloride, |
|
|
Identifiers |
|
|
10043-52-4 |
|
|
EV9800000 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
CaCl2 |
|
|
110.98 g/mol
(anhydrous) |
|
|
Appearance |
white powder |
|
odorless |
|
|
2.15 g/cm3
(anhydrous) |
|
|
772 °C (anhydrous) |
|
|
1935 °C (anhydrous) |
|
|
Solubility
in water |
74.5 g/100mL
(20 °C) |
|
soluble in acetone, acetic
acid |
|
|
Acidity (pKa) |
8–9 (anhydrous) |
|
Refractive
index (nD) |
1.52 |
|
Structure |
|
|
Orthorhombic (deformed rutile), oP6 |
|
|
Pnnm, No. 58 |
|
|
octahedral,
6-coordinate |
|
|
Hazards |
|
|
|
|
|
H319, H316, H302 |
|
|
P264, P280, P270,
P305+351+338, P337+313, P301+312, P330, P501 |
|
|
EU Index |
017-013-00-2 |
|
Irritant (Xi) |
|
|
1000 mg/kg (oral,
rat) |
|
|
Related compounds |
|
|
Other anions |
|
|
Other cations |
Beryllium chloride |
|
Phase behaviour |
|
|
|
|
Calcium chloride, CaCl2,
is a salt of calcium and chloride. It
behaves as a typical ionic
halide, and is solid at room
temperature. Common applications include brine for
refrigeration plants, ice
and dust control on roads, and desiccation.
Because of its hygroscopic nature, anhydrous
calcium chloride must be kept in tightly sealed, air-tight containers.
Calcium chloride can serve as a source of calcium ions in a solution, as
calcium chloride is soluble. This property can be useful for displacing ions from
solution. For example, phosphate is displaced from solution by calcium:
3 CaCl2
(aq) + 2 K3PO4 (aq) → Ca3(PO4)2
(s) + 6 KCl (aq)
Molten
calcium chloride can be electrolysed to give calcium metal and
chlorine
gas:
CaCl2
(l) → Ca (s) + Cl2 (g)
Calcium
chloride has a very high enthalpy change of solution.
The anhydrous salt is deliquescent; it can accumulate enough water in its crystal
lattice to form a solution.
Calcium chloride can be produced directly from limestone,
but large amounts are also produced as a byproduct of the Solvay
process. North American consumption in 2002 was 1,687,000 tons (3.7 billion
pounds). A Dow Chemical Company manufacturing facility in Michigan houses about
35% of the total U.S. production capacity for calcium chloride.
Calcium chloride occurs as the rare evaporite
minerals sinjarite (dihydrate) and antarcticite
(hexahydrate). A related mineral chlorocalcite
(potassium calcium chloride, KCaCl3) is also very rare.
Uses
Desiccant
Drying
tubes are frequently packed with calcium chloride. Kelp is dried with calcium
chloride for use producing sodium
carbonate. Adding solid calcium chloride to liquids can remove dissolved
water. Calcium chloride is also used in some air moisture absorbent products.
Anhydrous calcium chloride has been approved by the FDA as a packaging aid to
ensure dryness (CPG 7117.02).
These hygroscopic
properties are also applied to keep a liquid layer on the surface of the
roadway, which holds dust down.
By depressing the freezing point, calcium
chloride is used to usually prevent ice formation and to deice. This is
particularly useful on road surfaces. Calcium chloride dissolution is
exothermic, and is relatively harmless to plants and soil; however, recent
observations in Washington state suggest it may be particularly harsh on
roadside evergreen trees. It is also more effective at lower temperatures than
sodium chloride. When distributed for this use, it usually takes the form of
small, white balls a few millimeters in diameter, called prills. Solutions
of calcium chloride can prevent freezing at temperature as low as
−52 °C (−62 °F), making it ideal for filling agricultural
implement tires as a liquid ballast, aiding traction in cold climates.
Calcium chloride is used to increase the water
hardness in swimming pools. This reduces the erosion of the concrete in the
pool. By Le Chatelier's principle and the common
ion effect, increasing the concentration of calcium in the water will
reduce the dissolution of calcium compounds essential to the structure of
concrete.
In marine aquariums, calcium chloride is added to
introduce bioavailable calcium for calcium carbonate-shelled
animals such as mollusks
and cnidarians.
Calcium hydroxide (kalkwasser mix) or a calcium
reactor can also be used to introduce calcium, however calcium chloride
addition is the fastest method and has minimal impact on pH.
As an ingredient, it is listed as a permitted food
additive in the European Union for use as a sequestrant
and firming
agent with the E number E509, and considered as generally recognized as safe (GRAS) by
the U.S. Food and Drug Administration. The average intake of calcium chloride
as food additives has been estimated to be 160–345 mg/day for individuals.
As a firming agent, calcium chloride is used in canned
vegetables, in firming soybean curds into tofu and in producing a caviar substitute
from vegetable or fruit juices. It is commonly used as an electrolyte
in sports drinks and other beverages, including bottled water. The extremely
salty taste of calcium chloride is used to flavor pickles
while not increasing the food's sodium content. Calcium chloride's freezing-point depression
properties are used to slow the freezing of the caramel in caramel-filled
chocolate bars.
In brewing beer, calcium chloride is sometimes used to correct
mineral deficiencies in the brewing water. It affects flavor and chemical
reactions during the brewing process, and can also affect yeast function during
fermentation. Calcium chloride is sometimes added to processed milk to restore
the natural balance between calcium and protein in casein for the
purposes of making cheeses, such as brie,
Pélardon
and Stilton. Also, it is frequently added to sliced
apples to maintain texture.
Calcium chloride can be injected as intravenous therapy for the treatment of hypocalcaemia.
It can be used for magnesium intoxication. Calcium chloride injection may
antagonize cardiac toxicity as measured by electrocardiogram.
It can help to protect the myocardium from dangerously high levels of serum potassium in hyperkalemia.
Calcium chloride can be used to quickly treat calcium channel blocker toxicity,
from the side effects of drugs such as diltiazem (Cardizem) —
helping avoid potential heart attacks.[11]
Aqueous calcium chloride is used in genetic transformation of cells by increasing the
cell membrane permeability, inducing competence for DNA uptake (allowing DNA
fragments to enter the cell more readily).
Calcium chloride dihydrate (20% by weight) dissolved in ethanol (95% ABV)
has been used as a sterilant for male animals. The non surgical procedure
consists of the injection of the solution into the testes of the animal. Within
1 month, necrosis of testicular tissue results in sterilization.
Calcium chloride is also used to increase the Calcium
levels in marine (saltwater) reef aquariums. If calcium chloride, CaCl2,
is used as the source of calcium to a reef aquarium then the following can be
used to make a DIY stock solution. It can also be added as a dry solid, but it
has to be pre-mixed before addition to the aquarium and can be harder to
measure out. The use of a stock solution gets around both of these problems.
The variables used to make the calculations are as
follows:
·
M0 = Mass of calcium chloride required
to make the stock solution (grams)
·
V0 = Volume of stock solution (litres)
·
C0 = Calcium concentration of the stock
solution (ppm)
·
V1 = Volume of stock solution dose to
give calcium concentration rise of C2 (cm3)
·
V2 = System volume (litres)
·
C2 = Calcium concentration rise required
in the system (ppm)
·
Ca = Fraction of the type of calcium
chloride that is calcium by weight: 0.3611(CaCl2), 0.3107(CaCl2.H2O),
0.1829(CaCl2.6H2O)
The value of Ca depends on the type of calcium chloride
that is used to make up the stock solution. In most cases this will be CaCl2,
the dehydrated form, so the valve for Ca in this case is 0.3611.
The concentration of the stock solution, C0, is really
arbitrary, but the higher the concentration the less of the stock solution has
to be added to give a certain system concentration increase. Although the
higher the concentration the easier it would be to overdose and cause problems.
A good number to start with is around C0 = 50,000 ppm, which is the concentration
of some commercially available calcium chloride additives. The mass (M0) of
calcium chloride required to add to the stock solution of volume V0, to give
concentration C0 is given by:
M0 = ( C0 * V0 ) / ( Ca * 1000 )
Therefore to get C0 = 50,000ppm stock solution, using
CaCl2, M0 = 138 grams is required to be added to V0 = 1 litre.
The equation to determine the volume (V1, cm³) of
the stock solution (concentration C0 ppm) required to increase the system
(volume V2, litre) calcium concentration by C2 ppm is as follows:
V1 = ( C2 * V2 * 1,000 ) / C0
Therefore to increase the system of V2 = 100 litre
calcium concentration by C2 = 100 ppm, using a stock solution of
concentration C0 = 50,000ppm need to added V1 = 200 cm³.
Calcium chloride is used in concrete mixes to help speed
up the initial setting, but chloride ions lead to corrosion of steel rebar, so it should
not be used in reinforced concrete. The anhydrous form of
calcium chloride may also be used for this purpose and can provide a measure of
the moisture in concrete.
Calcium chloride is used in swimming
pool water as a pH buffer and to adjust the calcium hardness of the water.
Calcium chloride is included as an additive in plastics
and in fire extinguishers, in wastewater treatment as a
drainage aid, in blast furnaces as an additive to control scaffolding
(clumping and adhesion of materials that prevent the furnace charge from
descending), and in fabric softener as a thinner.
The exothermic dissolution of calcium chloride is used in
self-heating cans and heating
pads.
In the oil industry, calcium chloride is used to increase
the density of solids-free brines. It is also used to provide inhibition of
swelling clays in the water phase of invert emulsion drilling fluids.
CaCl2 acts as flux material (decreasing
melting point) in the Davy process for the industrial production of Sodium
metal, through the electrolysis of molten NaCl.
Calcium chloride is also an ingredient used in ceramic
slipware. It suspends clay particles so that they float within the solution
making it easier to use in a variety of slipcasting techniques.
Calcium chloride can act as an irritant by
desiccating moist skin. Solid calcium chloride dissolves exothermically, and burns
can result in the mouth
and esophagus
if it is ingested. Ingestion of concentrated solutions or solid products may
cause gastrointestinal irritation or ulceration.
Sodium chloride
|
Sodium chloride |
|
|
|
|
|
|
|
|
Sodium chloride |
|
|
Other names Common salt Halite Table salt |
|
|
Identifiers |
|
|
VZ4725000 |
|
|
3534976 |
|
|
13673 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
NaCl |
|
|
58.44 g mol−1 |
|
|
Appearance |
Colorless crystals |
|
Odorless |
|
|
2.165 g cm−3 |
|
|
801 °C, 1074 K, 1474 °F |
|
|
1413 °C, 1686 K, 2575 °F |
|
|
Solubility
in water |
359 g L−1 |
|
21.5 g L−1 |
|
|
14.9 g L−1 |
|
|
Refractive
index (nD) |
1.5442 (at 589 nm) |
|
Structure |
|
|
Cubic |
|
|
Fm3m, No. 225 |
|
|
a = 564.02 pm |
|
|
Octahedral (Na+) |
|
|
Thermochemistry |
|
|
Std enthalpy of |
-411.12 kJ mol−1 |
|
Standard molar |
72.11 J K−1
mol−1 |
|
36.79 J K−1
mol−1 |
|
|
3000–8000 mg/kg
(oral in rats, mice, rabbits)[1] |
|
|
Related compounds |
|
|
Other anions |
|
|
Other cations |
Lithium chloride |
|
Phase behaviour |
|
|
|
|
Sodium chloride, also known as salt,
common salt, table salt or halite, is an ionic
compound with the formula NaCl, representing equal proportions of sodium and chloride.
Sodium chloride is the salt most responsible for the salinity of the ocean and of the extracellular fluid of many multicellular organisms. As
the major ingredient in edible salt, it is commonly used as a condiment and
food preservative.
In solid sodium chloride, each ion is surrounded by six
ions of the opposite charge as expected on electrostatic grounds. The
surrounding ions are located at the vertices of a regular octahedron.
In the language of close-packing, the larger chloride ions are arranged in a
cubic array whereas the smaller sodium ions fill all the cubic gaps (octahedral voids) between
them. This same basic structure is found in many other compounds
and is commonly known as the halite or rock-salt crystal structure. It can be represented
as a face-centered cubic (fcc) lattice with a two-atom basis or as two interpenetrating
face centered cubic lattices. The first atom is located at each lattice point,
and the second atom is located half way between lattice points along the fcc
unit cell edge.
Thermal conductivity of NaCl as a function of
temperature has a maximum of 2.03 W/(cm K) at 8 K and decreases to 0.069 at 314
K (41 °C). It also decreases with doping.[2]
Aqueous solutions
The attraction between the Na+ and Cl-
ions in the solid is so strong that only highly polar solvents like water
dissolve NaCl well.
|
Solubility of NaCl in various
solvents |
|
|
360 |
|
|
94 |
|
|
83 |
|
|
71 |
|
|
52 |
|
|
30.2 |
|
|
14 |
|
|
0.65 |
|
|
0.4 |
|
|
0.124 |
|
|
0.05 |
|
|
0.05 |
|
|
0.03 |
|
|
0.018 |
|
|
0.003 |
|
|
0.00042 |
|
When dissolved in water, the sodium chloride framework
disintegrates as the Na+ and Cl- ions become surrounded
by the polar water molecules. These solutions consist of metal aquo complex with the formula [Na(H2O)8]+,
with the Na-O distance of 250 pm. The chloride ions are also strongly solvated, each
being surrounded by an average of 6 molecules of water. Solutions of sodium
chloride have very different properties from pure water. The freezing point is −21.12
°C for 23.31 wt% of salt, and the boiling point of saturated salt solution is
near 108.7 °C. From cold solutions, salt crystallises as the dihydrate NaCl·2H2O.
Salt is currently mass-produced
by evaporation
of seawater
or brine from brine wells
and salt lakes. Mining of
rock salt is also a major source. China is the world's main supplier of salt.
In 2010, world production was estimated at 270 million tonnes, the top five
producers (in million tonnes) being China (60.0), United States (45.0), Germany
(16.5), India (15.8) and Canada (14.0). Salt is also a byproduct of potassium
mining.
In addition to the familiar domestic uses of salt, more
dominant applications of the approximately 250 megatons/year production (2008
data) include chemicals and de-icing.
Salt is the source, directly or indirectly, for the
production of many chemicals, which consume most of the world's production.
Potassium chloride
|
Potassium chloride |
|
|
|
|
|
|
|
|
Other names Sylvite |
|
|
Identifiers |
|
|
TS8050000 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
KCl |
|
|
74.5513 g·mol−1 |
|
|
Appearance |
white crystalline
solid |
|
odorless |
|
|
1.984 g/cm3 |
|
|
770 °C |
|
|
1420 °C |
|
|
Solubility
in water |
281 g/L (0°C) |
|
soluble in glycerol, alkalies |
|
|
Acidity (pKa) |
~7 |
|
Refractive
index (nD) |
1.4902 (589 nm) |
|
Structure |
|
|
Thermochemistry |
|
|
Std enthalpy of |
−436 kJ·mol−1[2] |
|
Standard molar |
83 J·mol−1·K−1[2] |
|
2.6 g/kg
(oral/rat), 0.142 g/kg (intravenous/rat)[3] |
|
|
Related compounds |
|
|
Other anions |
|
|
Other cations |
Lithium chloride |
|
Related compounds |
|
|
|
|
The chemical compound potassium chloride (KCl)
is a metal halide salt
composed of potassium
and chlorine.
In its pure state, it is odorless and has a white or colorless vitreous
crystal
appearance, with a crystal structure that cleaves easily in three
directions. Potassium chloride crystals are face-centered cubic. Potassium chloride was
historically known as "muriate of potash". This name is occasionally still
encountered in association with its use as a fertilizer.
Potash varies in
color from pink or red to white depending on the mining and recovery
process used. White potash, sometimes referred to as soluble potash, is usually
higher in analysis and is used primarily for making liquid starter fertilizers.
KCl is used in medicine,
scientific applications, and food
processing. It occurs naturally as the mineral sylvite and in
combination with sodium chloride as sylvinite.
Chemical properties
In chemistry and physics, it is a very commonly used standard, for example as
a calibration
standard solution in measuring electrical conductivity of (ionic)
solutions, since carefully prepared KCl solutions have well-reproducible and
well-repeatable measurable properties.
|
Solubility of KCl in various
solvents |
|
|
360 |
|
|
0.4 |
|
|
0.41 |
|
|
5.3 |
|
|
192 |
|
|
0.04 |
|
|
0.024 |
|
|
0.00091 |
|
|
62 |
|
|
24.5 |
|
|
0.17–0.5 |
|
Potassium chloride can react as a source of chloride ion. As with any other soluble
ionic chloride, it will precipitate insoluble chloride salts when added to a solution of an
appropriate metal ion:
KCl(aq) + AgNO3(aq)
→ AgCl(s) + KNO3(aq)
Although potassium is more electropositive
than sodium, KCl
can be reduced to the metal by reaction with metallic sodium at 850°C because
the potassium is removed by distillation (see Le Chatelier's principle):
KCl(l)
+ Na(l) ⇌ NaCl(l) + K(g)
This method is the main method for producing metallic
potassium. Electrolysis (used for sodium) fails because of the high
solubility of potassium in molten KCl.
As with other compounds containing potassium, KCl in
powdered form gives a lilac flame test result.
Physical properties
Potassium
chloride has a crystalline structure like many other salts. Its structure is
face-centered cubic. Its lattice
constant is roughly 6.3Å. Some other properties are
·
Transmission range: 210 nm to
20 µm
·
Transmittivity
= 92% at 450 nm and rises linearly to 94% at 16 µm
·
Refractive index = 1.456 at 10 µm
·
Reflection Loss = 6.8% at 10 µm
(two surfaces)
·
dN/dT (expansion
coefficient)= −33.2×10−6/°C
·
dL/dT (refractive index
gradient)= 40×10−6/°C
·
Thermal conductivity = 0.036 W/(cm·K)
·
Damage threshold (Newman & Novak): 4
GW/cm2 or 2 J/cm2 (0.5 or 1 ns pulse rate); 4.2 J/cm2
(1.7 ns pulse rate Kovalev & Faizullov)
.jpg)
Sylvite
Sylvinite
Potassium chloride occurs naturally as sylvite, and it
can be extracted from sylvinite. It is also extracted from salt water and
can be manufactured by crystallization from solution, flotation
or electrostatic
separation from suitable minerals. It is a by-product of the making of nitric acid
from potassium nitrate and hydrochloric
acid.
The majority of the potassium chloride produced is used
for making fertilizer,
since the growth of many plants is limited by their potassium intake. As a chemical feedstock, it
is used for the manufacture of potassium hydroxide and potassium
metal. It is also used in medicine, lethal
injections, scientific
applications, food processing, and as a sodium-free substitute
for table
salt (sodium chloride).
It is sometimes used in water as a completion fluid in petroleum and
natural
gas operations, as well as being an alternative to sodium
chloride in household water softener units. KCl is useful as a beta
radiation source for calibration of radiation monitoring equipment,
because natural potassium contains 0.0118% of the isotope 40K.
One kilogram
of KCl yields 16350 becquerels of radiation
consisting of 89.28% beta and 10.72% gamma with
1.46083 MeV. Potassium chloride is used in some deicing products
that are designed to be safer for pets and plants, though these are inferior in
melting quality to calcium chloride (lowest usable temperature 12 °F
(−11 °C) v. −25 °F (−32 °C)). It is also used in
various brands of bottled water, as well as in bulk quantities for fossil fuel
drilling
purposes.
Potassium chloride was once used as a fire
extinguishing agent, used in portable and wheeled fire
extinguishers. Known as Super-K dry chemical, it was more effective than sodium bicarbonate-based dry chemicals and was
compatible with protein foam. This agent fell out of favor with
the introduction of potassium bicarbonate (Purple-K) dry
chemical in the late 1960s, which was much less corrosive and
more effective. It is rated for B and C fires.
Along with sodium
chloride and lithium chloride, potassium chloride is used as a flux
for the gas
welding of aluminium.
Potassium chloride is also an optical crystal with a wide
transmission range from 210 nm to 20 µm. While cheap, KCl crystal is hygroscopic.
This limits its application to protected environments or short term uses such
as prototyping. Exposed to free air, KCl optics will "rot". Whereas
KCl components were formerly used for infrared optics, it has been
entirely replaced by much tougher crystals like zinc
selenide.
Potassium chloride has also been used to create heat packs
which employ exothermic chemical
reactions,[5]
but these are no longer being created due to cheaper and more efficient
methods, such as the oxidation of metals ('Hot Hands', one time use products)
or the crystallization of sodium
acetate (multiple use products).
Potassium chloride is used as a scotophor
with designation P10 in dark-trace
CRTs, e.g. in the Skiatron.
Biological
and medical properties
Potassium is vital in the human body,
and oral potassium chloride is the common means to replenish it, although it
can also be diluted and given intravenously. It can be used as a salt
substitute for food,
but due to its weak, bitter, unsalty flavour, it is
usually mixed with ordinary table salt (sodium chloride) for this purpose to
improve the taste.
The addition of 1 ppm of thaumatin considerably reduces this bitterness. Medically,
it is used in the treatment of hypokalemia
and associated conditions, for digitalis poisoning, and as an electrolyte
replenisher. Brand names include K-Dur, Klor-Con, Micro-K, Slow-K, Sando-K and
Kaon Cl. Side effects can include gastrointestinal discomfort including nausea and vomiting, diarrhea and bleeding of the
digestive tract. Overdoses cause hyperkalemia,
which can lead to paresthesia, cardiac conduction blocks, fibrillation,
arrhythmias, and sclerosis. Prescription potassium
citrate (the potassium naturally found in fruits and vegetables) can be
prescribed as an alternative to potassium chloride. Slow-K is a 1950s development
where the medicine is formulated to enter the bloodstream at delayed intervals.
It was first only prescribed to British military forces to balance their diets
while serving in Korea.
Some cardiac
surgery procedures cannot be carried out on the beating heart. For these
procedures, the surgical team will bypass the heart with a heart-lung machine and inject potassium chloride
into the heart muscle to stop the heartbeat.
The lethal effects of potassium chloride
overdoses have led to its use in lethal
injection, as the third of a three-drug combination. Additionally, KCl is
used (albeit rarely) in fetal intracardiac injections in second- and
third-trimester induced abortions. Jack
Kevorkian's thanatron machine injected a lethal dose of potassium
chloride into the patient, which caused the heart to stop functioning, after a sodium
thiopental-induced coma was achieved.
Orally, potassium chloride is toxic in
excess; the LD50 is around 2.5 g/kg (meaning that
a lethal dose for 50% of people weighing 75 kg (165 lb)
is about 190 g (6.7 ounces)).
Intravenously, this is reduced to just over 30 mg/kg, but of more concern
are its severe effects on the cardiac muscles: high doses
can cause cardiac arrest and rapid death, thus the
aforementioned use as the third and final drug delivered in the lethal
injection process.
Bromide
|
Bromide |
|
|
Bromide[1] |
|
|
Identifiers |
|
|
3587179 |
|
|
14908 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
Br- |
|
|
79.904 g mol-1 |
|
|
Pharmacology |
|
|
12 d |
|
|
Thermochemistry |
|
|
Std enthalpy of |
−121 kJ·mol−1[2] |
|
Standard molar |
82 J·mol−1·K−1[2] |
|
|
|
A bromide is a chemical compound containing a
bromide ion or ligand. This is a bromine atom with
an ionic charge of −1 (Br-); for example,
in caesium bromide, caesium cations (Cs+)
are electrically attracted to bromide anions (Br-) to form the
electrically neutral ionic compound CsBr. The term "bromide"
can also refer to a bromine atom with an oxidation
number of -1 in covalent compounds such as sulfur dibromide (SBr2).
Bromide
is present in typical seawater (35 PSU) with a concentration of around 65 mg/L, which is
around 0.2% of all dissolved salts. Seafoods and deep sea plants generally have high levels
of bromide, while foods derived from land have variable amounts.
One can
test for a bromide ion by adding excess dilute HNO3
followed by dilute aqueous AgNO3 solution. The formation of creamy silver
bromide precipitate confirms the existence of bromides.
Medical uses
Bromide compounds, especially potassium
bromide, were frequently used as sedatives in the 19th and early 20th
century. Their use in over-the-counter sedatives and headache remedies (such as
Bromo-Seltzer)
in the United States extended to 1975, when bromides were withdrawn as
ingredients, due to chronic toxicity.
This use gave the word
"bromide" its colloquial connotation of a boring cliché,
a bit of conventional wisdom overused as a calming
phrase, or verbal sedative.
The bromide ion is antiepileptic, and
bromide salts are still used as such, particularly in veterinary medicine.
Bromide ion is excreted by the kidneys. The half-life of bromide in the human
body (12 days) is long compared with many pharmaceuticals, making dosing
difficult to adjust (a new dose may require several months to reach
equilibrium). Bromide ion concentrations in the cerebrospinal fluid are about
30% of those in blood, and are strongly influenced by the body's chloride
intake and metabolism.
Since bromide is still used in veterinary medicine
(particularly to treat seizures in dogs) in the United States, veterinary
diagnostic labs can routinely measure blood bromide levels. However, this is
not a conventional test in human medicine in the U.S., since there are no
FDA-approved uses for bromide, and (as noted) it is no longer available in
over-the-counter sedatives. Therapeutic bromide levels are measured in European
countries like Germany,
where bromide is still used therapeutically in human epilepsy.
Chronic toxicity from bromide can result in bromism, a
syndrome with multiple neurological symptoms. Bromide toxicity can also cause a
type of skin eruption. See potassium
bromide.
Lithium bromide was used as a sedative
beginning in the early 1900s, but it fell into disfavor in the 1940s, possibly
when some heart patients died after using a salt substitute (see lithium
chloride). Like lithium carbonate and lithium
chloride it was used as treatment for bipolar
disorder.
In biology
Bromide is needed by eosinophils
(white blood cells of the granulocyte class, specialized for dealing with multi-cellular
parasites), which use it to generate antiparasitic brominating compounds such
as hypobromite,
by the action of eosinophil peroxidase, a haloperoxidase
enzyme which is able to use chloride, but preferentially uses bromide when
available. Despite this use by the body, bromide is not known to be strictly
necessary for animal life, as its functions may generally be replaced (though
in some cases not as well) by chloride. Land plants also do not use bromide.
Bromide salts are also sometimes used in hot tubs and
spas as mild germicidal agents, using the action of an added oxidizing
agent to generate in situ hypobromite,
in a similar fashion to the peroxidase in eosinophils.
Bromide is also not a necessary nutrient for most animals
in the sea, although a few sea animals, such as Murex snails, use
bromide to make organic compounds. However, bromide ion is heavily concentrated
by some species of ocean algae, which construct methyl
bromide and a great number of bromoorganic compounds with it, using the
unusual enzymes called vanadium bromoperoxidases to do these
reactions.
The average concentration of bromide in human blood in
Queensland, Australia is 5.3±1.4 mg/L and varies with age and gender. Much
higher levels may indicate exposure to brominated chemicals (e.g. methyl
bromide). However, since bromide occurs in relatively high concentration in
seawater and many types of seafood, bromide concentrations in the blood are
heavily influenced by seafood contributions to the diet.
Characteristic
reactions Br- ions
Silver nitrate (pharmacopeia’s
reaction) with Br- ions forms pale
yellow precipitate of AgBr:
Br- + Ag+ = AgBr¯.
The precipitate insoluble in mineral acids, but
is slowly dissolved in 25 % a aqueous solution of ammonia.
Reaction
performance. To
2-3 drops of an investigated solution add 1-2 drops of 2 mol/L HNO3
solution and 2-3 drops of AgNO3 solution. If there are Br– ions, the pale
yellow precipitate forms; it is dissolved slowly in 25 % an ammonia solution.
Reaction of oxidation by chloric water (pharmacopeia’s
reaction). Ions Br- in the acidic medium are oxidized
by chloric water (or chlorammine in the presence), which paints a organic
solvent layer (benzene, chloroform) in yellow colour:
2Br- + Cl2 = Br2 + 2Cl–.
Reaction
performance. To
3-4 drops of an investigated solution add to 2-3 drop of 0,5 mL H2SO4
solution, benzene and 2-3 drops of chloric water. A test tube vigorously shake;
if there are Br- ions the benzene layer is painted in yellow colour.
The concentrated sulphatic acid with dry salts of bromide, for example NaBr, gives
HBr:
NaBr +
H2SO4 ®NaHSO4 + HBr.
HBr is oxidized by sulphatic acid to free Br2 (it is brown
gas):
2HBr +
H2SO4 = Br2 + 2H2O + SO2.
Reaction
performance.
5-6 drops of an investigated solution place in a high porcelain crucible and
dry to dry salts. To the dry rest add 2-3 drops concentrated H2SO4
(attention!), close a crucible by glass. If there are bromide-ions observe of formation
brown gas.
Lead (IV) oxide (pharmacopeia’s
reaction) oxidises bromide-ions in the acidic
medium to free bromine:
PbO2 + 2Br- +
4CH3COOH → Pb(CH3COO)2 + Br2
+ 2H2O + 2CH3COO-.
Reaction
performance. To
3-5 drops of an investigated solution add 0,1 g Lead (IV) oxide, 5-6 drops concentrated
acetic acid, some mL of benzene. If there are bromide-ions the organic layer
will be yellow or red (accordingly to amount of bromide ions).
Potassium bromide
|
Potassium bromide |
|
|
|
|
|
|
|
|
Identifiers |
|
|
TS7650000 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
KBr |
|
|
119.002 g/mol |
|
|
Appearance |
white solid |
|
odorless |
|
|
2.74 g/cm3 |
|
|
734 °C, 1007 K, 1353 °F |
|
|
1435 °C, 1708 K, 2615 °F |
|
|
Solubility
in water |
53.5 g/100 mL (0 °C) |
|
very slightly soluble in diethyl ether |
|
|
21.7 g/100 mL |
|
|
4.76 g/100 mL (80
°C) |
|
|
Refractive
index (nD) |
1.559 |
|
Structure |
|
|
10.41 D (gas) |
|
|
Related compounds |
|
|
Other anions |
|
|
Other cations |
Lithium
bromide |
|
|
|
Potassium bromide (KBr) is a salt,
widely used as an anticonvulsant and a sedative in the
late 19th and early 20th centuries, with over-the-counter use extending to 1975
in the US. Its action is due to the bromide ion (sodium
bromide is equally effective). Potassium bromide is used as a veterinary
drug, as an antiepileptic medication for dogs and cats.
Under
standard conditions, potassium bromide is a white crystalline powder. It is
freely soluble in water. In a dilute aqueous solution, potassium bromide tastes
sweet, at higher concentrations it tastes bitter, and at even higher
concentrations it tastes salty. These effects are mainly due to the properties
of the potassium ion—sodium bromide tastes salty at any concentration. In high
concentration, potassium bromide strongly irritates the gastric mucous
membrane, causing nausea and sometimes vomiting (a typical effect of all
soluble potassium salts).
Chemical properties
Potassium bromide, a typical ionic salt,
is fully dissociated and near pH 7 in aqueous solution. It serves as a source of bromide ions. This
reaction is important for the manufacture of silver
bromide for photographic film:
KBr(aq) + AgNO3(aq) → AgBr(s) + KNO3(aq)
Aqueous bromide Br- also forms complexes when reacted with some metal halides
such as copper(II) bromide:
2 KBr(aq) + CuBr2(aq) → K2[CuBr4](aq)
A traditional method for the manufacture of KBr is the
reaction of potassium carbonate with a bromide of iron, Fe3Br8,
made by treating scrap iron under water with excess bromine:
4 K2CO3 + Fe3Br8
→ 8 KBr + Fe3O4 + 4 CO2
Medical and veterinary
The anticonvulsant properties of potassium bromide were
first noted by Sir Charles Locock at a meeting of the Royal Medical and Chirurgical
Society in 1857. Bromide can be regarded as the first effective medication
for epilepsy.
At the time, it was commonly thought that epilepsy was caused by masturbation.
Locock noted that bromide calmed sexual excitement and thought this was
responsible for his success in treating seizures. In the latter half of the
19th century, potassium bromide was used for the calming of seizure and nervous
disorders on an enormous scale, with the use by single hospitals being as much
as several tons a year (the dose for a given person being a few grams per day).
There was not a better epilepsy drug until phenobarbital
in 1912. It was often said the British
Army laced soldiers' tea
with bromide to quell sexual arousal—but that is likely untrue as doing so
would also diminish alertness in battle and similar stories exist about a
number of substances.
Bromide compounds, especially sodium
bromide, remained in over-the-counter sedatives and headache remedies (such
as the original formulation of Bromo-Seltzer)
in the US until 1975, when bromides were outlawed in all over-the-counter
medicines, due to chronic toxicity. Bromide's exceedingly long half life in the
body made it difficult to dose without side effects (see below). Medical use of
bromides in the US was discontinued at this time, as many better and
shorter-acting sedatives were known by then.
Potassium bromide is used in veterinary medicine to treat
epilepsy
in dogs, either as first-line treatment or in addition to phenobarbital,
when seizures are not adequately controlled with phenobarbital alone. Use of
bromide in cats is limited because it carries a substantial risk of causing
lung inflammation (pneumonitis) in them. The use of bromide as a treatment drug
for animals means that veterinary medical diagnostic laboratories are able as a
matter of routine to measure serum levels of bromide on order of a
veterinarian, whereas human medical diagnostic labs in the US do not measure
bromide as a routine test.
Potassium bromide is not approved by the US Food and Drug Administration (FDA) for
use in humans to control seizures. In Germany, it is still approved as an
antiepileptic drug for humans, particularly children and adolescents. These
indications include severe forms of generalized tonic-clonic seizures,
early-childhood-related Grand-Mal-seizures, and also severe myoclonic seizures
during childhood. Adults who have reacted positively to the drug during
childhood/adolescence may continue treatment. Potassium bromide tablets are
sold under the brand name Dibro-Be mono (Rx-only). The drug has almost
complete bioavailability, but the bromide ion has a relatively long half life
of 12 days in the blood, making bromide salts difficult to adjust and dose.
Bromide is not known to interfere with the absorption or excretion of any other
anticonvulsant, though it does have strong interactions with chloride in the
body, the normal body uptake and excretion of which strongly influences
bromide's excretion.
The therapeutic index (ratio of effectiveness to
toxicity) for bromide is small. As with other antiepileptics, sometimes even
therapeutic doses (3 to 5 grams per day, taking 6 to 8 weeks to reach
stable levels) may give rise to intoxication. Often indistinguishable from
'expected' side-effects, these include:
·
Bromism These
are central nervous system reactions. They may include:
depression,
lethargy, somnolence
(from daytime sleepiness to coma)
loss of appetite
and cachexia,
nausea/emesis with exicosis (loss of body fluid)
loss of reflexes or
pathologic reflexes
loss of neural
sensitivity
cerebral
edema with associated headache and papilledema
of the eyes
delirium:
confusion, abnormal speech, loss of concentration and memory, aggressiveness
·
Acne-form dermatitis and other forms of
skin disease may also be seen, as well as mucous hypersecretion in the lungs.
Asthma and rhinitis may worsen. Rarely, tongue disorder, aphten, bad breath,
and obstipation occur.
Potassium bromide is transparent from the near ultraviolet
to long-wave infrared
wavelengths
(0.25-25 µm) and has no significant optical absorption lines in its high
transmission region. It is used widely as infrared optical windows and
components for general spectroscopy because of its wide spectral range. In infrared spectroscopy, samples are analyzed
by grinding with powdered potassium bromide and pressing into a disc.
Alternatively, samples may be analyzed as a liquid film (neat, as a solution,
or in a mull with Nujol)
between two polished potassium bromide discs.
Due to its high solubility and hygroscopic
nature it must be kept in a dry environment. The refractive
index is about 1.55 at 1.0 µm.
In addition to manufacture of silver bromide, potassium
bromide is used as a restrainer in black and white developer formulas. It improves
differentiation between exposed and unexposed crystals of silver halide, and
thus reduces fog.
Iodide
|
Iodide |
|
|
Iodide[1] |
|
|
Identifiers |
|
|
3587184 |
|
|
14912 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
I− |
|
|
126.90447 g mol-1 |
|
|
Related compounds |
|
|
Other anions |
|
|
Except where noted
otherwise, data are given for materials in their standard
state (at 25 °C, 100 kPa) |
|
An iodide ion is the ion I−.
Compounds with iodine
in formal oxidation state −1 are called iodides.
This page is for the iodide ion and its salts, not organoiodine compounds. In everyday life,
iodide is most commonly encountered as a component of iodized
salt, which many governments mandate. Worldwide, iodine
deficiency affects two billion people and is the leading preventable cause
of mental retardation.
Iodide is one of the largest monoatomic anions. It is
assigned a radius of around 206 picometers. For comparison, the lighter halides are
considerably smaller: bromide (196 pm), chloride (181 pm), and fluoride (133 pm).
In part because of its size, iodide forms relatively weak bonds with most
elements.
Most iodide salts are soluble in water, but often less so
than the related chlorides and bromides. Iodide, being large, is less
hydrophilic than are the smaller anions. One consequence of this is that sodium
iodide is highly soluble in acetone, whereas sodium chloride is not. The low
solubility of silver iodide and lead iodide
reflects the covalent character of these metal iodides. A test for the presence
of iodide ions is the formation of yellow precipitates of these compounds upon
treatment of a solution of silver nitrate or lead(II)
nitrate.
Aqueous solutions of iodide salts dissolve iodine better
than pure water. This effect is due to the formation of the triiodide
ion, which is brown:
I− + I2 ⇌ I3−
Redox, including antioxidant
properties
Iodide salts are mild reducing
agents and many react with oxygen to give iodine. A reducing agent is a
chemical term for an antioxidant. Its antioxidant properties can be expressed
quantitatively as a redox potential :
I−
⇌ 1/2 I2 + e− (electrons)
= - 0.54 Volt vs SHE
Because iodide is easily oxidized, some enzymes readily
convert it into electrophilic iodinating agents, as required for the biosynthesis
of myriad iodide-containing natural
products. Iodide can function as an antioxidant reducing species that
can destroy reactive oxygen species such as hydrogen
peroxide:
2 I− + Peroxidase + H2O2
+ tyrosine, histidine, lipid, etc. → iodo-Compounds + H2O + 2
e− (antioxidants).
Representative iodides
|
Compound |
Formula |
Appearance |
Use
or occurrence |
|
KI |
white crystals |
iodine component of iodized salt |
|
|
HI |
colourless solution |
strong mineral acid |
|
|
AgI |
yellow powder that darkens in light |
photoactive component of silver-based
photographic film |
|
|
Thyroxine |
C15H11I4NO4 |
pale yellow solid |
hormone essential for human health |
Characteristic
reactions I- ions
Silver nitrate (pharmacopeia’s
reaction) with ions I- forms yellow precipitate
Ag˛ which is not
dissolved in HNO3 and in a aqueous solution of ammonia (unlike AgCl
and AgBr):
Ag+ + I- = Ag˛¯.
Precipitate Ag˛ reacts with
metal zinc in the presence of 1 mol/L H2SO4 solution:
2Ag˛ + Zn = Zn2+ + 2I- + 2Ag¯.
Reaction
performance. To
1-2 drops of an investigated solution add 2-3 drops of 6 mol/L HNO3 solution
and 1-2 drops of AgNO3 solution. If there are bromide-ions the yellow
precipitate forms. It is not dissolved in aqueous solution of NH3. To
precipitate add some drops of 1 mol/L H2SO4 and metal
zinc.
Oxidation reaction. Chloric (bromic) water oxidises ions I- in the acidic medium
to I2:
2I- + Cl2 = I2
+ 2Cl–.
Iodine will dissolve in organic solvents and paints its in red-violet
colour.
Reaction
performance. To
1-2 drops of an investigated solution add 2-3 drops of 1 mol/L H2SO4
solution and some drops chloric (or bromic) water, some drops of benzene,
the test tube is shaked. If there are I- ions, the benzene layer is painted
in red-violet colour.
Salts of Lead with
ions I- form yellow or golden precipitate Pb˛2:
Pb2+
+ 2I- = Pb˛2¯.
This precipitate is dissolved in hot water, and then solution cooling yellow-golden
crystals of Pb˛2
form.
Reaction
performance. To 2-3 drops of an investigated solution add 2
drops of a solution of Pb2+ salt. If there are ions I-
yellow precipitate Pb˛2 forms.
Potassium bichromate (pharmacopeia’s
reaction) oxidises jodide ions in the medium
of sulphatic acid to free iodine:
Ńr2O72-+ 6˛- + 14H+
→ 2Cr3+ + 7H2O + 3˛2.
Reaction
performance. To
3-4 drops of investigated substance add 4-5 drops of 2 mol/L sulphatic acid
solution, 3-4 drops of Potassium bichromate solution, 2 mL water, 2 mL of
chloroform and shake. If there are iodide ions the chloroformic layer colours
in violet or violet-red colour.
The concentrated sulphatic acid with dry iodide
ions forms HI. HI is oxidized by sulphatic acid to free I2 (it is
violet gas):
2ĘI + H2SO4 = 2HI + Ę2SO4,
8I- + SO42- + 10H+
= 4I2¯ +
4H2O + H2S.
Iodine which is formed, allocated in the form
of a dark grey precipitate with characteristic metal shine or paints a solution
in brown colour; at small quantities I- ions the solution gets orange-red
colouring. By heating violet gas of iodine forms.
Reaction performance. As in case of Br- ions
(characteristic reactions on Br- ions see).
Potassium iodide
|
Potassium iodide |
|
|
|
|
|
|
|
|
Potassium iodide |
|
|
Identifiers |
|
|
TT2975000 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
KI |
|
|
166.0028 g/mol |
|
|
Appearance |
white crystalline
solid |
|
3.123 g/cm3 |
|
|
681 °C, 954 K, 1258 °F |
|
|
1330 °C, 1603 K, 2426 °F |
|
|
Solubility
in water |
128 g/100 ml (0 °C) |
|
2 g/100 mL (ethanol) |
|
|
Refractive
index (nD) |
1.677 |
|
Related compounds |
|
|
Other anions |
|
|
Other cations |
|
|
|
|
Potassium iodide is an inorganic compound with the chemical
formula KI. This white salt
is the most commercially significant iodide compound, with approximately 37,000
tons produced in 1985. It is less hygroscopic
(absorbs water less readily) than sodium
iodide, making it easier to work with. Aged and impure samples are yellow
because of aerial oxidation of the iodide to elemental
iodine.
4 KI + 2
CO2 + O2 → 2 K2CO3 + 2 I2
Potassium iodide is medicinally supplied in 130 mg
tablets (each containing 100 mg iodine as iodide) for emergency purposes
related to blockade of radioiodine uptake. Potassium iodide may also be
administered pharmaceutically for thyroid
storm or as an expectorant, as a "saturated solution of potassium
iodide" (SSKI) which in the U.S.P. generic
formulation contains 1000 mg of KI per mL of solution. This represents
333 mg KI and about 250 mg iodide (I -) in a typical adult
dose of 5 drops, assumed to be ⅓ mL. Because SSKI is a viscous liquid, it
is normally assumed to contain 15 drops/milliliter, not 20 drops/milliliter as
is often assumed for water. Thus, each drop of U.S.P. SSKI is assumed to
contain about 50 mg iodine as iodide, I -. Thus, two (2) drops
of U.S.P. SSKI solution is equivalent to one 130 mg KI tablet (100 mg
iodide).
SSKI can also be prepared by truly saturating water with
KI. This preparation can be made without a measuring scale. Since the
solubility of KI in water at room temperature is about 1.40 to 1.48 grams
per mL pure water, and the resulting solution has a density of about 1.72 g/mL,
this process also results in a final concentration of KI of about
1000 mg KI per mL of saturated KI solution, and also contains essentially
the same concentration of iodide per drop as does the U.S.P. formulation. Due
to its high potassium content, SSKI is extremely bitter, and if possible it is
administered in a sugar cube or small ball of bread. It may also be mixed into
much larger volumes of juices.
Neither SSKI or KI tablets are used as nutritional
supplements, since the nutritional requirement for iodine is only 150
micrograms or 0.15 mg of iodide per day. Thus, a drop of SSKI provides
50/0.15 = 333 times the daily iodine requirement, and a standard KI tablet
provides twice this much.
Kelp
is a natural KI source. The iodide content can range from 89 µg/g to
8165 µg/g in Asian varieties, making prepared foods content difficult to
estimate. Eating 3-5 grams of most dried, unrinsed seaweeds will provide
the 100-150 micrograms iodide recommended daily allowance for nutritional
purposes.
Potassium
iodide is ionic, K+I−.
It crystallises in the sodium chloride structure. It is produced
industrially by treating KOH with iodine.[1]
Since
the iodide ion
is a mild reducing
agent, I − is easily oxidised to I2 by
powerful oxidising
agents such as chlorine:
2
KI(aq)
+ Cl2(aq) → 2 KCl +
I2(aq)
This reaction is employed in the isolation of iodine from
natural sources. Air will oxidize iodide, as evidenced by the observation of a
purple extract when aged samples of KI are rinsed with dichloromethane.
As formed under acidic conditions, hydroiodic
acid (HI) is a stronger reducing agent.
Like other iodide salts, KI forms I3−
when combined with elemental iodine.
Unlike I2, I3−
salts can be highly water-soluble. Through this reaction, iodine is used in redox titrations.
Aqueous KI3, "Lugol's
solution," are used as disinfectants and as etchants for gold
surfaces.
Potassium iodide is the precursor to silver(I)
iodide, which is used for high speed photographic
film:
KI(aq)
+ AgNO3(aq)
→ AgI(s) + KNO3(aq)
KI serves as a source of iodide in organic
synthesis. A useful application is in the preparation of aryl iodides from arenediazonium salts. For example:
KI, acting as a source of iodide, may also act as a nucleophilic catalyst
for the alkylation of alkyl chlorides, bromides, or mesylates.
KI is a precursor to silver
iodide (AgI) an important chemical in film photography. KI is a component
in some disinfectants and hair treatment chemicals. KI is also used as a fluorescence quenching agent in biomedical
research, an application that takes advantage of collisional quenching of
fluorescent substances by the iodide ion. However, for several fluorophores
addition of KI in µM-mM concentrations results in increase of fluorescence
intensity, and iodide acts as fluorescence enhancer.
Potassium iodide is a component in the electrolyte of dye sensitized solar cells (DSSC) along
with iodine.
Potassium iodine finds its most important applications in
organic synthesis mainly in the preparation of aryl iodides in the Sandmeyer reaction, starting from aryl amines.
Aryl iodides are in turn used to attach aryl groups to other organics by
nucleophilic substitution, with iodide ion as the leaving group.
The major uses of KI include use as a nutritional
supplement in animal feeds and also the human diet. For the latter, it is the
most common additive used to "iodize" table
salt (a public health measure to prevent iodine
deficiency in populations which get little seafood). The oxidation of
iodide causes slow loss of iodine content from iodised
salts that are exposed to excess air. The alkali metal iodide salt, over
time and exposure to excess oxygen and carbon dioxide, slowly oxidizes to metal
carbonate and elemental iodine, which then evaporates.[14]
Potassium
iodate is used to add iodine to some salts so that the iodine is not lost
by oxidation.
For reasons noted above, therapeutic drops of SSKI, or
130 mg tablets of KI as used for nuclear fission accidents, are not used
as nutritional supplements, since an SSKI drop or nuclear-emergency tablet
provides 300 to 700 times more iodine than the daily adult nutritional
requirement. Dedicated nutritional iodide tablets containing 0.15 mg (150
microgram or mcg) of iodide, from KI or from various other sources (such as
kelp extract) are marketed as supplements, but they are not to be confused with
the much higher pharmaceutical dose preparations.
Pharmaceutical applications
Potassium iodide can be conveniently prepared as a
saturated solution, abbreviated SSKI. This method of delivering potassium
iodide does not require a method to weigh out the potassium iodide so it can be
used in an emergency situation. KI crystals are simply added to water until no
more KI will dissolve and instead sits at the bottom of the container. With
pure water, the concentration of KI in the solution depends only on the
temperature. Potassium iodide is highly soluble in water so SSKI is a
concentrated source of KI. At 20 degrees Celsius the solubility of KI is
140-148 grams per 100 grams of water. Because the volumes of KI and
water are approximately additive, the resulting SSKI solution will contain
about 1.40 gram (1400 mg) KI per milliliter (mL) of solution. This is
100% weight/volume (note units of mass concentration) of KI (one gram
KI per mL solution), which is possible because SSKI is significantly more dense
than pure water—about 1.72 g/mL. Because KI is about 76.4% iodide by weight,
SSKI contains about 764 mg iodide per mL. This concentration of iodide allows the
calculation of the iodide dose per drop, if one knows the number of drops per
milliliter. For SSKI, a solution more viscous than water, there are assumed to
be 15 drops per mL; the iodide dose is therefore approximately 51 mg per
drop, assuming 15 drops/mL. It is conventionally rounded to 50 mg per
drop.
The term SSKI is also used, especially by pharmacists, to
refer to a U.S.P. pre-prepared solution formula, made by adding exactly KI to
water to prepare a solution containing of 1000 mg KI per mL solution (100%
wt/volume KI solution), to closely approximate the concentration of SSKI made
by saturation. This is essentially interchangeable with SSKI made by saturation,
and also contains about 50 mg iodide per drop.
·
Saturated solutions of potassium iodide
can be an emergency treatment for hyperthyroidism
(so-called thyroid storm), as high amounts of iodide temporarily
suppress secretion of thyroxine from the thyroid gland. The dose typically
begins with a loading dose, then 1/3 mL SSKI (5 drops or 250 mg iodine as
iodide), three times per day.
·
Iodide solutions made from a few drops
of SSKI added to drinks have also been used as expectorants
to increase the water content of respiratory secretions and encourage effective
coughing.
·
SSKI has been proposed as a topical treatment
for sporotrichosis,
but no trials have been conducted to determine the efficacy or side effects of
such treatment.
·
Potassium iodide has been used for
symptomatic treatment of erythema nodosum patients for persistent lesions
whose cause remains unknown. It has been used in cases of erythema nodosum
associated with Crohn's disease.
Thyroid protection during medical treatment
Pheochromocytoma seen as dark sphere in center of
the body. Image is by MIBG
scintigraphy
with radiation from radioiodine in the MIBG. However, note unwanted uptake of
radioiodine from the pharmaceutical by the thyroid gland in the neck, in both
images (front and back) of the same patient. Radioactivity is also seen in the
bladder.
Thyroid iodine uptake blockade with potassium iodide is
used in nuclear medicine scintigraphy
and therapy with some radioiodinated compounds that are not targeted to the
thyroid, such as iobenguane (MIBG), which is used to image or treat neural tissue tumors, or
iodinated fibrinogen,
which is used in fibrinogen scans to
investigate clotting. These compounds contain iodine, but not in the iodide
form. However, since they may be ultimately metabolized or break down to
radioactive iodide, it is common to administer non-radioactive potassium iodide
to ensure that iodide from these radiopharmaceuticals is not sequestered by the
normal affinity of the thryoid for iodide.
U.S. Food and Drug Administration-approved
dosing of potassium iodide for this purpose with iobenguane, is as follows (per
24 hours): infants less than 1 month old, 16 mg; children 1 month to 3
years, 32 mg; children 3 years to 18 years, 65 mg; adults
130 mg. However, some sources recommend alternative dosing regimens.
Not all sources are in agreement on the necessary duration
of thyroid blockade, although agreement appears to have been reached about the necessity
of blockade for both scintigraphic and therapeutic applications of iobenguane.
Commercially available iobenguane is labeled with iodine-123,
and product labeling recommends administration of potassium iodide 1 hour prior
to administration of the radiopharmaceutical for all age groups, while the
European Associated of Nuclear Medicine recommends (for iobenguane labeled with
either isotope,) that potassium iodide administration begin one day prior to radiopharmaceutical administration, and continue
until the day following the injection, with the exception of new-borns, who do
not require potassium iodide doses following radiopharmaceutical injection.
Product labeling for diagnostic iodine-131 iobenguane
recommends potassium iodide administration one day before injection and
continuing 5 to 7 days following administration, in keeping with the much longer
half-life of this isotope and its greater danger to the thyroid. Iodine-131
iobenguane used for therapeutic purposes requires a different pre-medication
duration, beginning 24–48 hours prior to iobenguane injection and continuing
10–15 days following injection.
Thyroid
protection due to nuclear accidents and emergencies
In 1982, the U.S. Food and Drug Administration approved
potassium iodide to protect thyroid glands from radioactive
iodine involving accidents or fission emergencies. In an accidental event
or attack on a nuclear power plant, or in nuclear
bomb fallout, volatile fission product radionuclides may be released. Of
these products, 131I is one of the most common and is
particularly dangerous to the thyroid gland because it may lead to thyroid
cancer. By saturating the body with a source of stable iodide prior to
exposure, inhaled or ingested 131I tends to be excreted, which
prevents radioiodine uptake by the thyroid. The protective effect of KI lasts
approximately 24 hours. For optimal prophylaxis,
KI must be dosed daily until a risk of significant exposure to radioiodine by
either inhalation or ingestion no longer exists.
Emergency 130 milligrams potassium iodide doses
provide 100 mg iodide (the other 30 mg is the potassium in the
compound), which is roughly 700 times larger than the normal nutritional need
(see recommended dietary allowance) for
iodine, which is 150 micrograms (0.15 mg) of iodine (as iodide) per
day for an adult.
Potassium iodide cannot protect against any other causes
of radiation poisoning, nor can it provide any
degree of protection against dirty bombs that produce radionuclides other than
radionuclides of iodine. See fission
products and the external links for more details concerning radionuclides.
|
WHO Recommended Dosage for Radiological
Emergencies involving radioactive iodine |
|
|
Age |
KI in mg per day |
|
Over 12 years old |
130 |
|
3 – 12 years old |
65 |
|
1 – 36 months old |
32 |
|
< 1 month old |
16 |
The potassium iodide in iodized
salt is insufficient for this use. A likely lethal dose
of salt (more than a kilogram) would be needed to equal the potassium
iodide in one tablet.
The World Health Organization does not
recommend KI prophylaxis for adults over 40 years, unless inhaled radiation
dose levels are expected to threaten thyroid function; because, the KI side
effects increases with age and may exceed the KI protective effects
"...unless doses to the thyroid from inhalation rise to levels threatening
thyroid function, that is of the order of about 5 Gy.
Such radiation doses will not occur far away from an accident site."
The U.S. Department of Health and Human Services restated
these two years later as "The downward KI (potassium iodide) dose
adjustment by age group, based on body size considerations, adheres to the
principle of minimum effective dose. The recommended standard (daily) dose of
KI for all school-age children is the same (65 mg). However, adolescents
approaching adult size (i.e., >70 kg [154 lbs]) should receive the full
adult dose (130 mg) for maximal block of thyroid radioiodine uptake. Neonates
ideally should receive the lowest dose (16 mg) of KI."
SSKI (i.e., the solution of KI rather than tablets) may
be used in radioiodine-contamination emergencies (i.e., nuclear accidents) to
"block" the thyroid's uptake of radioiodine, at a dose of two drops
of SSKI per day for an adult. This is not the same as blocking the thyroid's
release of thyroid hormone, for which the adult dose is different (and is actually
higher by a factor of 7 or 8), and for which KI anti-radiation pills (not a
common medical treatment form of KI) are not usually available in pharmacies,
or normally used in hospitals, or by physicians. Although the two forms of
potassium iodide are completely interchangeable, normally in practice the SSKI
solution, which is the historical medical form of high dose iodine, is
generally used for all medical purposes save for radioiodine prophylaxis. For
protection of the thyroid against radioiodine (iodine-131) contamination, the
convenient standard 130 mg KI pill is used if available. As noted, the
equivalent two drops of SSKI may be used for this purpose, if the pills are not
available.
Following the Chernobyl nuclear reactor disaster in April,
1986, a saturated solution of potassium iodide (SSKI) was administered to 10.5
million children and 7 million adults in Poland as a
prophylactic measure against accumulation of radioactive iodine-131
in the thyroid
gland. People in the areas immediately surrounding Chernobyl itself, however,
were not given the supplement.
Potassium iodide’s (KI) value as a radiation protective
(thyroid blocking) agent was demonstrated at the time of the Chernobyl nuclear
accident when Soviet authorities distributed it in a 30 km zone around the
plant. The purpose was to protect residents from radioactive iodine, a highly
carcinogenic material found in nuclear reactors which had been released by the
damaged reactor. Only a limited amount of KI was available, but those who
received it were protected. Later, the US Nuclear Regulatory Commission (NRC)
reported, “thousands of measurements of I-131 (radioactive iodine)
activity...suggest that the observed levels were lower than would have been
expected had this prophylactic measure not been taken. The use of KI...was
credited with permissible iodine content in 97% of the evacuees tested.”
Poland, 300 miles from Chernobyl, also distributed KI to
protect its population. Approximately 18 million doses were distributed, with
follow-up studies showing no known thyroid cancer among KI recipients. With the
passage of time, people living in irradiated areas where KI was not available
have developed thyroid cancer at epidemic levels, which is why the US Food and
Drug Administration (FDA) reported “The data clearly demonstrate the risks of
thyroid radiation... KI can be used [to] provide safe and effective protection
against thyroid cancer caused by irradiation.
Chernobyl also demonstrated that the need to protect the
thyroid from radiation was greater than expected. Within ten years of the
accident, it became clear that thyroid damage caused by released radioactive
iodine was virtually the only adverse health effect that could be measured. As
reported by the NRC, studies after the accident showed that “As of 1996, except
for thyroid cancer, there has been no confirmed increase in the rates of other
cancers, including leukemia, among the... public, that have been attributed to
releases from the accident.”
But equally important to the question of KI is the fact
that radiation releases are not “local” events. Researchers at the World Health
Organization accurately located and counted the cancer victims from Chernobyl
and were startled to find that “the increase in incidence [of thyroid cancer]
has been documented up to 500 km from the accident site... significant doses
from radioactive iodine can occur hundreds of kilometers from the site, beyond
emergency planning zones." Consequently, far more people than anticipated
were affected by the radiation, which caused the United Nations to report in
2002 that “The number of people with thyroid cancer... has exceeded
expectations. Over 11,000 cases have already been reported.”
These findings were consistent with studies of the
effects of previous radiation releases. In 1945, millions of Japanese were
exposed to radiation from nuclear weapons, and the effects can still be
measured. Today, nearly half (44.8%) the survivors of Nagasaki studied have
identifiable thyroid disease, with the American Medical Association reporting
“it is remarkable that a biological effect from a single brief environmental
exposure nearly 60 years in the past is still present and can be detected.”
This, as well as the development of thyroid cancer among residents in the North
Pacific from radioactive fallout following the United States' nuclear weapons testing in
the 1950s (on islands nearly 200 miles downwind of the tests) were instrumental
in the decision by the FDA in 1978 to issue a request for the availability of
KI for thyroid protection in the event of a release from a commercial nuclear
power plant or weapons-related nuclear incident. Noting that KI’s effectiveness
was “virtually complete” and finding that iodine in the form of potassium
iodide (KI) was substantially superior to other forms including iodate
(KIO3) in terms of safety, effectiveness, lack of side effects,
and speed of onset, the FDA invited manufacturers to submit applications to
produce and market KI.
Today, three companies (Anbex, Inc., Fleming Co, and
Recip of Sweden) have met the strict FDA requirements for manufacturing and
testing of KI, and they offer products (IOSAT, ThyroShield, and Thyro-Safe,
respectively) which are available for purchase. The Swedish manufacturing
facility for Thyrosafe, a potassium iodide tablet for thyroid protection from
radiation manufactured by Recipharm AB, was mentioned on the secret US 2008 Critical Foreign Dependencies
Initiative leaked by Wikileaks in 2010.
It was reported on March 16, 2011, that potassium iodide
tablets were given prophylactically to U.S. Naval air crew members flying
within 70 nautical miles of the Fukushima Daiichi Nuclear Power
Plant damaged in the massive Japanese earthquake (8.9/9.0 magnitude) and
ensuing tsunami on March 11, 2011. The measures were seen as precautions, and
the Pentagon said no U.S. forces have shown signs of radiation poisoning. By
March 20, the US Navy instructed personnel coming within 100 miles of the
reactor to take the pills.
Adverse reactions
There is reason for caution with prescribing the
ingestion of high doses of potassium iodide and iodate, as their unnecessary
use can cause conditions such as the Jod-Basedow phenomena, and the Wolff-Chaikoff effect, trigger and/or worsen hyperthyroidism
and hypothyroidism,
and ultimately cause temporary or even permanent thyroid conditions. It can
also cause sialadenitis (an inflammation of the salivary gland),
gastrointestinal disturbances, allergic reactions and rashes. Potassium iodide
is also not recommended for those who have had an allergic reaction to iodine,
and people with dermatitis herpetiformis and hypocomplementemic vasculitis,
conditions that are linked to a risk of iodine sensitivity.
There have been some reports of potassium iodide
treatment causing swelling of the parotid
gland (one of the three glands which secrete saliva), due to its
stimulatory effects on saliva production.
A saturated solution of KI (SSKI) is typically given orally
in adult doses of about 250 mg iodide several times a day (5 drops of SSKI
assumed to be ⅓ ml) for thyroid blockade (to prevent the thyroid from
excreting thyroid hormone) and occasionally this dose is also used, when iodide
is used as an expectorant (the total dose is about one gram KI per day for an
adult). The anti-radioiodine doses used for I-131 uptake blockade
are lower, and range downward from 100 mg a day for an adult, to less than
this for children (see table). All of these doses should be compared with the
far lower dose of iodine needed in normal nutrition, which is only
150 μg per day (150 micrograms, not milligrams).
At maximal doses, and sometimes at much lower doses, side
effects of iodide used for medical reasons, in doses of 1000 times the normal
nutrional need, may include: acne, loss of appetite, or upset stomach
(especially during the first several days, as the body adjusts to the
medication). More severe side effects which require notification of a physician
are: fever, weakness, unusual tiredness, swelling in the neck or throat, mouth
sores, skin rash, nausea, vomiting, stomach pains, irregular heartbeat,
numbness or tingling of the hands or feet, or a metallic taste in the mouth.
The administration of known goitrogen
substances can also be used as a prophylaxis
in reducing the bio-uptake of iodine, (whether it be the nutritional
non-radioactive iodine-127 or radioactive iodine, radioiodine - most
commonly iodine-131,
as the body cannot discern between different iodine isotopes). perchlorate
ions, a common water contaminant in the USA due to the aerospace industry, has been shown to reduce
iodine uptake and thus is classified as a goitrogen.
Perchlorate ions are a competitive inhibitor of the process by which iodide, is
actively deposited into thyroid follicular cells. Studies involving healthy
adult volunteers determined that at levels above 0.007 milligrams per kilogram
per day (mg/(kg·d)), perchlorate begins to temporarily inhibit the thyroid
gland’s ability to absorb iodine from the bloodstream ("iodide uptake
inhibition", thus perchlorate is a known goitrogen). The reduction of the
iodide pool by perchlorate has dual effects – reduction of excess hormone
synthesis and hyperthyroidism, on the one hand, and reduction of thyroid
inhibitor synthesis and hypothyroidism on the other. Perchlorate remains very
useful as a single dose application in tests measuring the discharge of
radioiodide accumulated in the thyroid as a result of many different
disruptions in the further metabolism of iodide in the thyroid gland.
Treatment of thyrotoxicosis (including Graves' disease)
with 600-2,000 mg potassium perchlorate (430-1,400 mg perchlorate)
daily for periods of several months or longer was once common practice,
particularly in Europe, and perchlorate use at lower doses to treat thryoid
problems continues to this day. Although 400 mg of potassium perchlorate
divided into four or five daily doses was used initially and found effective,
higher doses were introduced when 400 mg/day was discovered not to control
thyrotoxicosis in all subjects.
Current regimens for treatment of thyrotoxicosis
(including Graves' disease), when a patient is exposed to additional sources of
Iodine, commonly include 500 mg potassium perchlorate twice per day for
18–40 days.
Prophylaxis with perchlorate containing water at
concentrations of 17 ppm, which corresponds to 0.5 mg/kg-day
personal intake, if one is 70 kg and consumes 2 litres of water per day,
was found to reduce baseline radioiodine uptake by 67% This is equivalent to
ingesting a total of just 35 mg of Perchlorate ions per day. In another
related study were subjects drank just 1 litre of perchlorate containing water
per day at a concentration of 10 ppm, i.e. daily 10 mg of Perchlorate ions
were ingested, an average 38% reduction in the uptake of Iodine was observed.
However when the average perchlorate absorption in
perchlorate plant workers subjected to the highest exposure has been estimated
as approximately 0.5 mg/kg-day, as in the above paragraph, a 67% reduction
of iodine uptake would be expected. Studies of chronically exposed workers
though have thus far failed to detect any abnormalities of thyroid function,
including the uptake of iodine. this may well be attributable to sufficient
daily exposure or intake of healthy Iodine-127 among the workers and the short
8 hr Biological half life of Perchlorate in the
body.
To completely block the uptake of Iodine-131 by the
purposeful addition of perchlorate ions to a populaces water supply, aiming at
dosages of 0.5 mg/kg-day, or a water concentration of 17 ppm, would
therefore be grossly inadequate at truly reducing radioiodine uptake.
Perchlorate ion concentrations in a regions water supply, would need to be much
higher, at least 7.15 mg/kg of body weight per day or a water
concentration of 250 ppm,assuming people drink 2 liters of water per
day, to be truly beneficial to the population at preventing bioaccumulation
when exposed to a radioiodine environment, independent of the availability of Iodate or Iodide drugs.
The continual distribution of perchlorate tablets or the
addition of perchlorate to the water supply would need to continue for no less
than 80–90 days, beginning immediately after the initial release of radioiodine
was detected, after 80–90 days had passed released radioactive iodine-131 would
have decayed to less than 0.1% of its initial quantity at which time the danger
from biouptake of iodine-131 is essentially over.
In the event of a radioiodine release the ingestion of
prophylaxis potassium iodide, if available, or even iodate, would rightly take
precedence over perchlorate administration, and would be the first line of
defense in protecting the population from a radioiodine release. However in the
event of a radioiodine release too massive and widespread to be controlled by
the limited stock of iodide & iodate prophylaxis drugs, then the addition
of perchlorate ions to the water supply, or distribution of perchlorate tablets
would serve as a cheap, efficacious, second line of defense against carcinogenic
radioiodine bioaccumulation.
The ingestion of goitrogen drugs is, much like potassium
iodide also not without its dangers, such as hypothyroidism.
In all these cases however, despite the risks, the prophylaxis benefits of
intervention with iodide, iodate or perchlorate outweigh the serious cancer
risk from radioiodine bioaccumulation in regions were radioiodine has
sufficiently contaminatated the environment.
Sulfide
|
Sulfide |
|
|
|
|
|
Sulfanediide (substitutive) |
|
|
Identifiers |
|
|
Jmol-3D images |
|
|
Properties |
|
|
S2− |
|
|
32.065 g mol-1 |
|
|
Related compounds |
|
|
Other anions |
|
|
Except where noted
otherwise, data are given for materials in their standard
state (at 25 °C, 100 kPa) |
|
A sulfide (IUPAC-recommended
spelling) or sulphide (UK) is an anion of sulfur in its lowest oxidation
state of 2-. Sulfide is also a slightly archaic term for thioethers, a
common type of organosulfur compound that are well known for
their bad odors.
The dianion S2− exists only in strongly alkaline aqueous
solutions. Such solutions can form by dissolution of H2S or alkali
metal salts such as Li2S, Na2S, and K2S in the
presence of extra hydroxide. The ion S2− is exceptionally
basic with a pKa > 14. It does not exist in appreciable
concentrations even in highly alkaline water, being undetectable at pH < ~15
(8 M NaOH).
Hydrogen sulfide ion
Hydrogen sulfide
Instead, sulfide combines with protons to form HS−,
which is variously called hydrogen sulfide ion, hydrosulfide ion,
sulfhydryl ion, or bisulfide ion. At still lower pH (<7), HS−
converts to H2S, hydrogen
sulfide.
Sulfides are moderately strong reducing agents. They
react with oxygen in the air in elevated temperatures to form higher-valence
sulfur salts, such as sulfates and sulfur
dioxide.
Aqueous solutions of transition
metals cations
react with sulfide sources (H2S, NaHS, Na2S) to
precipitate solid sulfides. Such inorganic
sulfides typically have very low solubility in water, and many are related to
minerals with the same composition (see below). One famous example is the
bright yellow species CdS or "cadmium
yellow". The black tarnish formed on sterling silver is Ag2S.
Such species are sometimes referred to as salts. In fact, the bonding in
transition metal sulfides is highly covalent, which gives rise to their semiconductor
properties, which in turn is related to the deep colors. Several have practical
applications as pigments, in solar cells, and as catalysts.
Dissolved free sulfides (H2S, HS−
and S2−) are very aggressive species for the corrosion of many
metals such as steel, stainless steel, and copper. Sulfides present in aqueous
solution are responsible for stress corrosion cracking (SCC) of steel,
and is also known as sulfide stress cracking. Corrosion is a
major concern in many industrial installations processing sulfides: sulfide ore
mills, deep oil
wells, pipeline transporting soured oil, Kraft paper
factories. Microbially-induced corrosion (MIC) or biogenic sulfide corrosion are also
caused by sulfate reducing bacteria producing
sulfide.
Oxidation of sulfide can also form thiosulfate
(S2O32−) an intermediate species
responsible for severe problems of pitting
corrosion of steel and stainless steel while the medium is also acidified
by the production of sulfuric acid when oxidation is more advanced.
In organic
chemistry, "sulfide" usually refers to the linkage C-S-C,
although the term thioether is less ambiguous. For example, the thioether dimethyl
sulfide is CH3-S-CH3. Polyphenylene sulfide (see below) has the
empirical formula C6H4S. Occasionally, the term sulfide
refers to molecules containing the -SH functional
group. For example, methyl sulfide can mean CH3-SH. The
preferred descriptor for such SH-containing compounds is thiol or mercaptan,
i.e. methanethiol, or methyl mercaptan.
Confusion arises from the different meanings of the term
"disulfide".
Molybdenum disulfide (MoS2)
consists of separated sulfide centers, in association with molybdenum in the
formal 4+ oxidation state (Mo4+). Iron disulfide (pyrite, FeS2)
on the other hand consists of S22−, or −S–S−
dianion, in association with divalent iron in the formal 2+ oxidation state
(ferrous ion: Fe2+). Dimethyldisulfide
has the chemical binding CH3-S–S-CH3, whereas carbon
disulfide has no S–S bond, being S=C=S (linear molecule analog to CO2).
Most often in sulfur chemistry and in biochemistry, the disulfide term is
commonly ascribed to the sulfur analogue of the peroxide −O–O−
bond. The disulfide
bond (−S–S−) plays a major role in the conformation of proteins and in
the catalytic activity of enzymes.
Examples
|
Formula |
|
Melting
point (°C) |
Boiling
point (°C) |
CAS
number |
|
H2S |
Hydrogen sulfide is a very toxic and corrosive gas
characterised by a typical odour of "rotten egg". |
-85,7 |
-60,20 |
|
|
CdS |
Cadmium sulfide can be used in photocells. |
1750 |
|
|
|
|
Calcium polysulfide ("lime sulfur") is a traditional fungicide in gardening. |
|
|
|
|
CS2 |
Carbon disulfide is sometimes used as a solvent in industrial chemistry. |
-111.6 |
46 |
|
|
PbS |
Lead sulfide is used in infra-red sensors. |
1114 |
|
|
|
MoS2 |
Molybdenum
disulfide,
the mineral molybdenite, is used as a catalyst to remove
sulfur from fossil fuels; also as lubricant for high-temperature and
high-pressure applications. |
|
|
|
|
Cl-CH2CH2-S-CH2CH2-Cl |
Sulfur mustard (mustard gas) is an organosulfide
(thioether) that has been used as a chemical weapon in the First World War,
the chloride on the molecule acts as a leaving group when in the presence of
water and forms a thioether-alcohol and HCl. |
13 - 14 |
217 |
|
|
Ag2S |
Silver sulfide is formed on silver electrical
contacts
operating in an atmosphere rich in hydrogen sulfide. |
|
|
|
|
Na2S |
Sodium sulfide is an important industrial
chemical, used in manufacture of kraft paper, dyes, leather tanning, crude petroleum processing, treatment of heavy
metal pollution, and others. |
920 |
1180 |
|
|
ZnS |
Zinc sulfide is used for lenses and other optical devices in the infrared part of the spectrum. Zinc sulfide doped with silver is used in alpha detectors while zinc
sulfide with traces of copper has applications in photoluminescent strips for emergency lighting and
luminous watch dials. |
|
1185 |
|
|
MeS |
Several metal sulfides are used as pigments in art, although their use has declined somewhat due to their toxicity. Sulfide pigments include cadmium, mercury, and arsenic. |
|
|
|
|
C6H4S |
Polyphenylene
sulfide is
a polymer commonly called "Sulfar". Its repeating units are bonded
together by sulfide (thioether) linkages. |
|
|
|
|
SeS2 |
Selenium sulfide is an antifungal used in anti-dandruff
preparations, such as Selsun Blue. The presence of the highly toxic
selenium in healthcare and cosmetics products represents a general health and
environmental concern. |
<100 |
|
|
|
FeS2 |
The crystal lattice of pyrite is made of iron disulfide, in which iron is divalent and
present as ferrous ion (Fe2+). |
600 |
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Characteristic
reactions S2- ions
Silver nitrate forms black precipitate of Silver sulphide:
2Ag+ + S2- =
Ag2S¯.
Precipitate Ag2S is not dissolved in
diluted HNO3 on a cold, but well dissolved by heating.
Reaction
performance. To
2-3 drops of an investigated solution add some drops of AgNO3. If
there are ions S2- the black precipitate forms.
Sodium nitroprusside Na2[Fe(CN)5(NO)] with ions S2- by pH>7 forms complex Na4[Fe(CN)5NOS]
red-violet colour.
Reaction
performance. To
2-3 drops of an investigated solution add some drops of NaOH and 1-2 drops of Sodium
nitroprussid solution. If there are ions S2- red-violet solution
forms.
Acids, for
example, diluted H2SO4 and HCl, with sulphide ions form
gas H2S:
Na2S + H2SO4 =
Na2SO4 + H2S.
H2S is detected on smell or by reaction with paper moistened by
Pb(CH3COO)2 or Na4[Pb(OH)6] solution (it
will be black):
H2S + Pb(CH3COO)2
= PbS¯ + 2CH3COOH;
H2S + Na4[Pb(OH)6]
= PbS¯ + 4NaOH + 2H2O.
Reaction
performance. In
crucible place some drops of an investigated solution and add 1 mol/L H2SO4
or HCl solution. A crucible is covered by thin glass with a paper moistened
with a solution of Lead (II) salt. If there are ions S2- the paper
blackens.
Salts of Cadmium Cd2+ with S2- ions form pale
yellow precipitate CdS:
Cd2+ + HS- =
CdS¯ + H+.
If to this precipitate add 1-2 drops of CuSO4
solution the precipitate will be black CuS:
CdS + Cu2+ = Cd2+
+ CuS¯.
Formation CdS is used for separation of S2-
ions from others anions which contain S. For sedimentation S2- ions use
Cadmium carbonate.
Reaction
performance. To
4-5 drops of an investigated solution (neutral or basic) add some crystals of
CdCO3 and mix, shake. If there are sulphide ions, near white precipitate
of CdCO3 will be yellow precipitate CdS or a solution will be
yellow.
Thiocyanate
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Thiocyanate |
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cyanosulfanide |
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Other names[hide] sulphocyanate, thiocyanide |
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Identifiers |
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Jmol-3D images |
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Properties |
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SCN- |
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58.0824 |
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Thiocyanate (also known as rhodanide)
is the anion [SCN]−. It is the conjugate
base of thiocyanic acid. Common derivatives include the
colourless salts potassium thiocyanate and sodium thiocyanate. Organic
compounds containing the functional
group SCN are also called thiocyanates. Mercury(II) thiocyanate was formerly used
in pyrotechnics.
Thiocyanate is analogous to the cyanate ion,
[OCN]−, wherein oxygen is replaced by sulfur. [SCN]−
is one of the pseudohalides, due to the similarity of its reactions
to that of halide
ions. Thiocyanate used to be known as rhodanide (from a Greek
word for rose)
because of the red colour of its complexes with iron. Thiocyanate is
produced by the reaction of elemental sulfur or thiosulfate
with cyanide:
8 CN− + S8 → 8 SCN−
CN− + S2O32−
→ SCN− + SO32−
The
second reaction is catalyzed by the enzyme sulfotransferase
known as rhodanase
and may be relevant to detoxification of cyanide in the body.
Structure,
bonding and coordination chemistry
Resonance structures of the thiocyanate ion
Thiocyanate shares its negative charge approximately equally
between sulfur and nitrogen. As a consequence, thiocyanate can act as a nucleophile
at either sulfur or nitrogen — it is an ambidentate
ligand. [SCN]− can also bridge two (M−SCN−M)
or even three metals (>SCN− or −SCN<). Experimental evidence
leads to the general conclusion that class
A metals (hard acids) tend to form N-bonded thiocyanate
complexes, whereas class B metals (soft acids)
tend to form S-bonded thiocyanate complexes. Other factors, e.g.
kinetics and solubility, are sometimes involved, and linkage isomerism can
occur, for example [Co(NH3)5(NCS)]Cl2 and
[Co(NH3)5(SCN)]Cl2.
Organic and transition metal derivatives of the
thiocyanate ion can exist as "linkage
isomers." In thiocyanates, the organic group (or metal ion) is
attached to sulfur: R−S−C≡N has a S-C single bond and a C-N
triple bond. In isothiocyanates, the substituent
is attached to nitrogen: R−N=C=S has a S-C double bond and a C-N double
bond:
Phenylthiocyanate and phenylisothiocyanate are linkage
isomers and are bonded differently.
Organic thiocyanates are hydrolyzed to thiocarbamates
in the Riemschneider thiocarbamate
synthesis.
Test for
iron(III)
If [SCN]− is
added to a solution containing iron (III) ions (Fe3+), a blood red solution is
formed due to the formation of [Fe(NCS)(H2O)5]2+.
-3D-balls.png)
The
blood-red coloured complex pentaaqua(thiocyanato-N)iron(III), [Fe(NCS)(H2O)5]2+,
indicates the presence of Fe3+ in solution
iron(III)_chloride.jpg)
Iron(III) test
Biological
chemistry of thiocyanate in medicine
Thiocyanate is known to be an important
part in the biosynthesis of hypothiocyanite
by a lactoperoxidase. Thus the complete absence of
thiocyanate or reducted thiocyanate in the human body, (e.g., cystic
fibrosis) is damaging to the human host defense system. Thiocyanate is a
potent competitive inhibitor of the thyroid sodium-iodide symporter.
Thiocyanate is a metabolite
of sodium
nitroprusside, after rhodanese
catalyses its reaction with thiosulfate.
Characteristic reactions SCN– ions.
Silver nitrate AgNO3 with tiotsianat ions forms
white precipitate AgSCN:
Ag+ + SCN- =
AgSCN¯.
The precipitate is not dissolved in diluted nitric
acid, slightly dissolved in an aqueous ammonia solution.
Reaction
performance. To
2-3 drops of an investigated solution add 1-2 drops of solution AgNO3.
If there are SCN- ions, white precipitate AgSCN forms.
Iron (˛˛˛) salts with SCN- ions form red or pink (at small
concentration SCN–) complex:
Fe3+ + 3SCN- = Fe(SCN)3.
Reaction
performance. To 2-3 drops of an investigated
solution add 1-2 drops of 1 mol/L HCl solution and 1 drop of FeCl3
solution. If there are SCN- ions the solution will be pink or
bright red.
Bromate
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Bromate |
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Identifiers |
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Jmol-3D images |
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Properties |
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Br O3 |
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The bromate anion, BrO3−,
is a bromine-based
oxoanion. A bromate
is a chemical compound that contains this ion.
Examples of bromates include sodium
bromate, (NaBrO3), and potassium
bromate, (KBrO3).
Bromates are formed many different ways in municipal
drinking water. The most common is the reaction of ozone and bromide:
Br− + O3
→ BrO3−
Electrochemical processes, such as electrolysis of brine without a membrane
operating to form hypochlorite, will also produce bromate when bromide ion
is present in the brine solution.
Photoactivation (sunlight
exposure) will encourage liquid or gaseous chlorine to generate bromate in
bromide-containing water.
In laboratories bromates can be synthesized by dissolving
Br2 in a concentrated solution of potassium hydroxide (KOH). The following
reactions will take place (via the intermediate creation of hypobromite):
Br2 + 2 OH−
→ Br− + BrO− + H2O
3 BrO− → BrO3−
+ 2 Br−
Bromate in drinking water is undesirable because it is a suspected human carcinogen.
Its presence in Coca Cola's Dasani bottled water forced a recall of that product in the UK.
Although few by-products are formed by ozonation, ozone
reacts with bromide ions in water to produce bromate. Bromide can be found in
sufficient concentrations in fresh water to produce (after ozonation) more than
10 ppb of bromate--the maximum contaminant level established by the USEPA.
Proposals to reduce bromate formation include: lowering the water pH below 6.0,
limiting the doses of ozone, using an alternate water source with a lower
bromide concentration, pretreatment with ammonia and addition of small
concentrations of chloramines prior to ozonation.
On December 14, 2007, the Los Angeles Department of
Water and Power (LADWP) announced that it would drain Silver Lake Reservoir and Elysian Reservoir
due to bromate contamination. At the Silver Lake and Elysian reservoirs a
combination of bromide from well water, chlorine, and sunlight had formed
bromate. The decontamination took 4 months, discharging over 600 million US
gallons (2.3×106 m3)
of contaminated water.
On June 9, 2008 the LADWP began covering the surface of
the 10-acre (4 ha), 58-million-US-gallon (0.22×106 m3)
open Ivanhoe Reservoir with black, plastic balls to block the sunlight which
causes the naturally present bromide to react with the chlorine used in
treatment. It will require 30 million of the 40 cent balls ($12 million) to
cover the Ivanhoe and Elysian reservoirs.
Characteristic reactions BrÎ3- ions.
Silver nitrate AgNO3 with concentrated solutions
of bromate ions froms
pale yellow precipitate AgBrÎ3, which is soluble in diluted nitric and
sulphatic acids:
Ag+ + BrÎ3-= AgBrÎ3¯.
Silver bromate is dissolved also in a aqueous ammonia solution
and Potassium cyanide.
Reaction
performance. To 5-6 drops of an investigated
solution add 2-3 drops of Silver nitrate solution. If there are bromate ions,
pale yellow precipitate forms.
Bromide and iodide ions in the acidic medium are oxidised bromate ions to free bromine and
iodine:
BrÎ3-+ 5Br- +
6H+ = 3Br2 + 3H2O;
BrÎ3-+ 6I- +
6H+ = 3I2 + Br- +3H2O.
Reaction
performance. To 2-3 drops of an investigated
solution add 1-2 drops of solution HCl and 1-2 drops of KI or KBr solution and
0,5 mL of chloroform. If there are bromate ions the chloroformic layer will be
painted in violet or violet-red colour (by addition of KI), and red or brown
colour (by addition of KBr).
Iodate
The
iodate anion, IO3−
Space-filling model of the iodate anion
An iodate is a conjugate
base of iodic
acid. In the iodate anion, iodine is bonded to three oxygen atoms and
the molecular formula is IO3−.
The molecular geometry of iodate is trigonal pyramidal.
Iodate can be obtained by reducing periodate
with a thioether.
The byproduct of the reaction is a sulfoxide.
Iodates are a class of chemical
compounds containing this group. Examples are sodium
iodate (NaIO3), silver
iodate (AgIO3), and calcium
iodate (Ca(IO3)2). iodates resemble chlorates with
iodine instead of chlorine.
In acid conditions, iodic acid
is formed. Potassium hydrogen iodate (KH(IO3)2) is a double salt
of potassium iodate and iodic acid and an acid as well. Iodates
are used in the iodine clock reaction.
Characteristic reactions IO3- ions.
Silver nitrate AgNO3 with iodate ions forms white curdled
precipitate AgIO3 which is dissolved in all reagents which give
complex of Silver:
Ag+ + IO3-
= AgIO3¯;
AgIO3 + 2NH3×H2O = [Ag(NH3)2]+
+ 2H2O + IO3-;
AgIO3 + 2CN- = [Ag(CN)2]-
+ IO3-.
Reaction
performance. To 2-3 drops of an investigated
solution add 1-2 drops of solution AgNO3. If there are iodate ions,
white curdled precipitate forms.
Iodide in
the acidic medium (even acetic acid) are oxidised iodate ions to iodine:
IO3- + 5I- +
6H+ = 3I2 + 3H2O.
Reaction
performance. To 2-3 drops of an investigated
solution add 1-2 drops of HCl solution and 1-2 drops of KI solution. If there
are iodate ions, the solution is painted in brown colour (if it isn’t enough iodate
ions, the solution will have orange colouring).
Iron (II) Salts
in the acidic medium reduce HIO3 to iodine:
10Fe2+ + 2IO3-
+ 12H+ = 10Fe3+ + I2 + 6H2O.
Reaction performance. To 2-3 drops of an investigated solution add
1-2 drops of 2 mol/L HCl solution and 1-2 drops of FeCl2 solution.
If there are iodate ions, the solution will be brown-red colourin Complex compounds
A complex (or coordination compound)
is a compound, which consist either of complex ions with other ions of opposite
charge or a neutral complex species.
Complex ions
are ions formed from a metal atom or ion with Lewis bases attached to it by
coordinate covalent bonds.
Ligands are the
Lewis bases attached to the metal atom in a complex. They are electron-pair
donors, so ligands may be neutral molecules (such as H2O or NH3)
or anions (such as CN– or Cl–) that have at least one atom with alone pair of
electrons.
Cations
only rarely function as ligands. We might expect this, because an electron pair
on a cation is held securely by the positive charge, so it would not be
involved in coordinate bonding. A cation in which the positive charge is far
removed from an electron pair that could be donated can function as a ligand.
An example is the pyrazinium ion.
A polydentate ligand ("having many teeth") is a ligand
that can bond with two or more atoms to a metal atom. A complex formed by
polydentate ligands is frequently quite stable and is called a chelate.
Because of the stability of chelates, polydentate ligands (also called
chelating agents) are often used to remove metal ions from a chemical system.
Complexation Reactions
A more general definition of acids and bases was
proposed by G. N. Lewis (1875–1946) in
1923. The Brønsted–Lowry definition of acids and bases focuses on an acid’s
proton-donating ability and a base’s proton-accepting ability. Lewis theory, on the other hand,
uses the breaking and forming of covalent bonds to describe acid–base
characteristics. In this treatment, an acid is an electron pair acceptor, and a base is an
electron pair donor. Although Lewis theory can be applied to the treatment of acid–base
reactions, it is more useful for treating complexation reactions between metal ions
and ligands.
The following
reaction between the metal ion Cd2+ and the ligand NH3 is
typical of a complexation reaction.
Cd2+ +
4(:NH3) =
Cd(:NH3)42+
The product of
this reaction is called a metal–ligand complex. In writing the equation for this reaction,
we have shown ammonia as :NH3 to emphasize the pair of electrons it
donates to Cd2+. In subsequent reactions we will omit this notation.
The formation of a
metal–ligand complex is described by a formation constant, Kf. The complexation
reaction between Cd2+ and NH3, for example, has the following
equilibrium constant
The reverse of reaction is called a dissociation
reaction and is characterized by a dissociation
constant, Kd, which is the reciprocal of Kf.
Many complexation reactions occur in a stepwise
fashion. For example, the reaction
between Cd2+ and NH3 involves
four successive reactions
Cd2+ + NH3 = Cd(NH3)2+
Cd(NH3)2+ + NH3
= Cd(NH3)22+
Cd(NH3)22+
+ NH3 = Cd(NH3)32+
Cd(NH3)32+
+ NH3 = Cd(NH3)42+
This creates a problem since it no longer is clear
what reaction is described by a formation constant. To avoid
ambiguity, formation constants are divided into two categories.
Stepwise formation
constants, which are designated as Ki for the ith step, describe the successive
addition of a ligand to the metal–ligand complex formed in the previous step.
Thus, the equilibrium constants for these reactions are, respectively, K1, K2, K3, and K4. Overall, or cumulative formation
constants, which are designated as bi, describe the addition of i ligands to the
free metal ion. The equilibrium constant expression given
in equation 6.16, therefore, is correctly identified as b4, where
b4 = K1 ´ K2 ´ K3 ´ K4
In general
bi = K1 ´ K2 ´ . . . ´ Ki
The
formation constant, or stability constant, Kf, of a
complex ion is the equilibrium constant for the formation of the complex ion
from the aqueous metal ion and the ligands:
Ag+
+ 2NH3 « Ag(NH3)2+
Kf = 
The
dissociation constant, Kd, for a complex ion is the reciprocal, or
inverse, value of Kf:
Ag(NH3)2+
« Ag+
+ 2NH3 Kd
= 
Ladder
Diagrams for Complexation Equilibria
The same
principles used in constructing and interpreting ladder diagrams for acid–base
equilibria can be applied to equilibria involving metal–ligand complexes. For
complexation reactions the ladder diagram’s scale is defined by the concentration
of uncomplexed, or free ligand, pL. Using the formation of Cd(NH3)2+
as an example
Cd2+ +
NH3 = Cd(NH3)2+
we can easily
show that the dividing line between the predominance regions for Cd2+
and Cd(NH3)2+ is log(K1).
Since K1 for Cd(NH3)2+
is 3.55·102, log(K1) is 2.55. Thus, for a pNH3
greater than 2.55 concentrations of NH3
less than 2.8·10–3 M), Cd2+ is the predominate species.
A complete ladder diagram for the metal–ligand complexes of Cd2+ and NH3
is shown in Figure.
Influence various factors on complex compound stability
1. Stability
of complex compounds is more in complexes with high coordination number.
2. Concentration
of complex compounds in solution direct depends to ligand concentration and is
inversely proportional to metal ion concentration.
3. Equilibrium
in solution of complex compounds depend to pH (concentration of hydrogen ions)
and dissociation constant. Increasing the pH value is a cause of complex
compounds destroying (hydrolysis).
4. The
most complicated is temperature influence on complex compound stability.
Reaction of complex formation may be endothermic or exothermic. Heating can
induces such chemical processes:
–
changing acidic-basic equilibrium,
–
destroying some ligands,
–
oxidation some ligands or metal ions,
–
hydrolysis complex ions.
The most important complex compounds with inorganic ligands, used in
analysis
1.
Ammonia:
–
selection (colourless complex):
[Ag(NH3)2]+, [Zn(NH3)4]+2,
[Cd(NH3)4]+2;
–
detection (coloured complex): [Cu (NH3)4]+2,
[Co(NH3)6]+3, [Ni(NH3)4]+2.
2.
Halogen and rhodanide:
–
selection with extraction in
inorganic solvents;
–
detection (coloured complex):
[Fe(SCN)3]–3, [BiJ4]–,
[CoCl4]–2.
3.
Fluor – separation and masking (colourless
complex): [FeF6]–3.
4.
Cyanide – determination (coloured
complex): [Fe(CN)6]–3, [Fe(CN)6]–2.
Using complex ions in analysis
1. On
application and investigation of complex compounds in analysis may arise next
problems:
1)
determination of nature and quantity
of complex particles in solution;
2)
determination of structure of complex
compounds in solution;
3)
calculation of dissociation constant;
4)
determination of molar particles of
metal ions and ligands in complex compounds.
1.
Determination of cations with
coloured complex compounds.
2.
Masking of preventing cations in
stabile colourless complex compounds.
3.
Selection of cations with hydroxo- or
ammonia- complex compounds on systematic analysis.
4.
Dissolving of insoluble sediments:
AgCl + NH4OH, HgO + KCN.
5.
Changing of acidic-basic properties
of weak electrolytes: boric acid + glycerine.
Chemistry and
Properties of EDTA
Ethylenediaminetetraacetic acid, or EDTA, is an
aminocarboxylic acid. The structure of EDTA is shown in
Figure:
EDTA, which is a Lewis acid, has six binding
sites (the four carboxylate groups and
the two amino groups), providing six pairs of
electrons. The resulting metal–ligand complex, in which EDTA forms a cage-like structure
around the metal ion (Figure 9.25b), is very stable. The actual number of coordination
sites depends on the size of the metal ion; however, all metal–EDTA complexes
have a 1:1 stoichiometry.
MetalÐEDTA Formation Constants To
illustrate the formation of a metal–EDTA complex
consider the reaction between Cd2+ and EDTA
where Y4– is a shorthand notation for the
chemical form of EDTA shown in Figure. The formation constant for this reaction
is quite large, suggesting that the reaction’s
equilibrium position lies far to the right. Formation
constants for other metal–EDTA complexes are found in Appendix 3C.
EDTA Is a Weak Acid Besides its
properties as a ligand, EDTA is also a weak acid. The fully
protonated form of EDTA, H6Y2+, is a hexaprotic weak acid
with successive pKa values of pKa1 = 0.0 pKa2 = 1.5 pKa3 = 2.0 pKa4 = 2.68 pKa5 = 6.11 pKa6 = 10.17.
The first four values are for the carboxyl protons,
and the remaining two values are for the ammonium
protons. A
ladder diagram for EDTA is shown in Figure 9.26.
The species Y4– becomes the predominate form of EDTA
at pH levels greater than 10.17. It is only
for pH levels greater than 12 that Y4– becomes the only significant form of EDTA.
Conditional MetalÐLigand Formation Constants Recognizing
EDTA’s acid–base properties is important. The
formation constant for CdY2– in equation assumes that
EDTA is present as Y4–. If we restrict the pH to levels greater than
12, then equation 9.11 provides an adequate description of
the formation of CdY2–. For pH levels
less than 12, however, Kf overestimates the stability of the
CdY2– complex. At any pH a
mass balance requires that the total concentration of unbound EDTA
equal the combined concentrations of each of its forms.
CEDTA =
[H6Y2+] + [H5Y+] + [H4Y]
+ [H3Y–] + [H2Y2–] + [HY3–]
+ [Y4–]
To correct the formation constant for EDTA’s acid–base
properties, we must account for the
fraction, aY4–, of EDTA present as Y4–.
Values of a(Y4–) are shown in Table
9.12. Solving equation 9.12 for [Y4–] and substituting into the equation
for the formation constant gives
If we fix the pH using a buffer, then a(Y4–) is a constant. Combining a(Y4–) with Kf
gives
where Kf´ is a conditional formation
constant whose value depends on the pH. As
shown in Table 9.13 for CdY2–, the
conditional formation constant becomes smaller, and the
complex becomes less stable at lower pH levels.
EDTA Must Compete with Other Ligands To maintain a
constant pH, we must add a buffering agent.
If one of the buffer’s components forms a metal–ligand complex with Cd2+,
then EDTA must compete with the ligand for Cd2+. For example, an NH4+/NH3
buffer includes the ligand NH3, which forms several stable Cd2+–NH3 complexes. EDTA
forms a stronger complex with Cd2+ and will displace NH3.
The presence of NH3, however, decreases the stability of
the Cd2+–EDTA complex. We can account for
the effect of an auxiliary complexing agent, such as NH3, in the same way we
accounted for the effect of pH. Before adding EDTA, a mass balance on Cd2+
requires that the total concentration of Cd2+, CCd, be
CCd = [Cd2+]
+ [Cd(NH3)2+] + [Cd(NH3)22+] +
[Cd(NH3)32+] + [Cd(NH3)42+]
The fraction, α(Cd2+),
present as uncomplexed Cd2+ is
Solving equation 9.14 for [Cd2+] and
substituting into equation 9.13 gives
If the concentration of NH3 is held
constant, as it usually is when using a buffer, then
we can rewrite this equation as
where Kf˝ is a new conditional formation
constant accounting for both pH and the presence
of an auxiliary complexing agent. Values of α(Mn+) for
several metal ions are provided in
Table 9.14.
Coordination
complex
Cisplatin,
PtCl2(NH3)2
A platinum atom with four ligands
In chemistry, a coordination
complex or metal complex, consists of an atom or ion (usually
metallic), and a surrounding array of bound
molecules or anions, that are in turn known as ligands or
complexing agents. Many metal-containing compounds consist of coordination
complexes.
Coordination complexes are so pervasive that the
structure and reactions are described in many ways, sometimes confusingly. The
atom within a ligand that is bonded to the central atom or ion is called the donor
atom. A typical complex is bound to several donor atoms, which can be the
same or different. Polydentate (multiple bonded) ligands consist of several donor
atoms, several of which are bound to the central atom or ion. These complexes
are called chelate complexes, the formation of such complexes
is called chelation, complexation, and coordination.
The central atom or ion, together with all ligands
comprise the coordination sphere. The central atoms or ion
and the donor atoms comprise the first coordination sphere.
Coordination refers to the
"coordinate covalent bonds" (dipolar
bonds) between the ligands and the central atom. Originally, a complex
implied a reversible association of molecules, atoms, or ions through such weak chemical
bonds. As applied to coordination chemistry, this meaning has evolved. Some
metal complexes are formed virtually irreversibly and many are bound together
by bonds that are quite strong.
Structure of hexol
Coordination complexes were known – although not
understood in any sense – since the beginning of chemistry, e.g. Prussian
blue and copper vitriol. The key breakthrough occurred
when Alfred
Werner proposed in 1893 that Co(III) bears six ligands in an octahedral geometry. His theory
allows one to understand the difference between coordinated and ionic in a
compound, for example chloride in the cobalt ammine chlorides
and to explain many of the previously inexplicable isomers.
In 1914, Werner resolved the first coordination complex,
called hexol, into
optical isomers, overthrowing the theory that only carbon compounds could
possess chirality.
The ions or molecules surrounding the central atom are
called ligands.
Ligands are generally bound to the central atom by a coordinate covalent bond (donating
electrons from a lone electron pair into an empty metal orbital), and are
said to be coordinated to the atom. There are also organic ligands such
as alkenes whose
pi bonds
can coordinate to empty metal orbitals. An example is ethene in the
complex known as Zeise's salt, K+[PtCl3(C2H4)]−.
In coordination chemistry, a structure is first described
by its coordination number, the number of ligands
attached to the metal (more specifically, the number of donor atoms). Usually
one can count the ligands attached, but sometimes even the counting can become
ambiguous. Coordination numbers are normally between two and nine, but large
numbers of ligands are not uncommon for the lanthanides and actinides. The
number of bonds depends on the size, charge, and electron configuration of the metal ion and
the ligands. Metal ions may have more than one coordination number.
Typically the chemistry of complexes is dominated by
interactions between s and p molecular
orbitals of the ligands and the d orbitals of the metal ions. The s, p, and
d orbitals of the metal can accommodate 18 electrons (see 18-Electron
rule). The maximum coordination number for a certain metal is thus related
to the electronic configuration of the metal ion (to be more specific, the
number of empty orbitals) and to the ratio of the size of the ligands and the
metal ion. Large metals and small ligands lead to high coordination numbers,
e.g. [Mo(CN)8]4−. Small metals with large ligands
lead to low coordination numbers, e.g. Pt[P(CMe3)]2. Due
to their large size, lanthanides, actinides, and
early transition metals tend to have high coordination numbers.
Different ligand structural arrangements result from the
coordination number. Most structures follow the points-on-a-sphere pattern (or,
as if the central atom were in the middle of a polyhedron
where the corners of that shape are the locations of the ligands), where
orbital overlap (between ligand and metal orbitals) and ligand-ligand
repulsions tend to lead to certain regular geometries. The most observed
geometries are listed below, but there are many cases that deviate from a
regular geometry, e.g. due to the use of ligands of different types (which
results in irregular bond lengths; the coordination atoms do not follow a
points-on-a-sphere pattern), due to the size of ligands, or due to electronic
effects (see, e.g., Jahn–Teller distortion):
·
Linear for two-coordination
·
Trigonal
planar for three-coordination
·
Tetrahedral or square
planar for four-coordination
·
Trigonal bipyramidal or square pyramidal for
five-coordination
·
Octahedral (orthogonal) or trigonal
prismatic for six-coordination
·
Pentagonal bipyramidal for seven-coordination
·
Square
antiprismatic for eight-coordination
·
Tri-capped trigonal prismatic
(Triaugmented triangular prism) for nine-coordination.
Some exceptions and provisions should be noted:
·
The idealized descriptions of 5-, 7-,
8-, and 9- coordination are often indistinct geometrically from alternative
structures with slightly different L–M–L (ligand–metal–ligand) angles. The
classic example of this is the difference between square pyramidal and trigonal
bipyramidal structures.
·
Due to special electronic effects such
as (second-order) Jahn–Teller stabilization, certain geometries
are stabilized relative to the other possibilities, e.g. for some compounds the
trigonal prismatic geometry is stabilized relative to octahedral structures for
six-coordination.
The arrangement of the ligands is fixed for a given
complex, but in some cases it is mutable by a reaction that forms another
stable isomer.
There exist many kinds of isomerism in
coordination complexes, just as in many other compounds.
Stereoisomerism
Stereoisomerism occurs with the same bonds in
different orientations relative to one another. Stereoisomerism can be further
classified into:
Cis–trans
isomerism and facial–meridional isomerism
Cis–trans isomerism occurs in octahedral and square
planar complexes (but not tetrahedral). When two ligands are mutually
adjacent they are said to be cis, when opposite each other, trans.
When three identical ligands occupy one face of an octahedron, the isomer is
said to be facial, or fac. In a fac isomer, any two identical
ligands are adjacent or cis to each other. If these three ligands and
the metal ion are in one plane, the isomer is said to be meridional, or mer.
A mer isomer can be considered as a combination of a trans and a cis,
since it contains both trans and cis pairs of identical ligands.
.png)
cis-[CoCl2(NH3)4]+ .png){kind=spacer}
trans-[CoCl2(NH3)4]+
.png)
fac-[CoCl3(NH3)3] .png)
mer-[CoCl3(NH3)3]
Optical
isomerism
Optical isomerism occurs when a molecule is not
superposable with its mirror image. It is so called because the two isomers are
each optically active, that is, they rotate the plane
of polarized light in opposite directions. The symbol
Λ (lambda)
is used as a prefix to describe the left-handed propeller twist formed by three
bidentate ligands, as shown. Likewise, the symbol Δ (delta)
is used as a prefix for the right-handed propeller twist.
ferrate(III)-3D-balls.png)
Λ-[Fe(ox)3]3− ferrate(III)-3D-balls.png)
Δ-[Fe(ox)3]3− cobalt(III).png)
Λ-cis-[CoCl2(en)2]+
cobalt(III).png)
Δ-cis-[CoCl2(en)2]+
Structural isomerism occurs when the bonds are
themselves different. There are four types of structural isomerism: ionisation
isomerism, solvate or hydrate isomerism, linkage isomerism and coordination
isomerism.
1.
Ionisation isomerism – the
isomers give different ions in solution although they have the same
composition. This type of isomerism occurs when the counter ion of the complex
is also a potential ligand. For example pentaaminebromidocobalt(III)sulfate
[Co(NH3)5Br]SO4 is red violet and in solution
gives a precipitate with barium chloride, confirming the presence of sulfate
ion, while pentaaminesulfatecobalt(III)bromide [Co(NH3)5SO4]Br
is red and tests negative for sulfate ion in solution, but instead gives a
precipitate of AgBr with silver nitrate.
2.
Solvate or hydrate isomerism – the
isomers have the same composition but differ with respect to the number of
solvent ligand molecules as well as the counter ion in the crystal lattice. For
example [Cr(H2O)6]Cl3 is violet colored, [Cr(H2O)5Cl]Cl2·H2O
is blue-green, and [Cr(H2O)4Cl2]Cl·2H2O
is dark green
3.
Linkage
isomerism occurs with ambidentate ligands that can bind in more
than one place. For example, NO2 is an ambidentate ligand: It can
bind to a metal at either the N atom or an O atom.
4.
Coordination isomerism – this
occurs when both positive and negative ions of a salt are complex ions and the
two isomers differ in the distribution of ligands between the cation and the
anion. For example [Co(NH3)6][Cr(CN)6] and
[Cr(NH3)6][Co(CN)6].
Electronic
properties
Many of the properties of metal complexes are dictated by
their electronic structures. The electronic structure can be described by a
relatively ionic model that ascribes formal charges to the metals and ligands.
This approach is the essence of crystal field theory (CFT). Crystal field
theory, introduced by Hans Bethe in 1929, gives a quantum
mechanically based attempt at understanding complexes. But crystal field
theory treats all interactions in a complex as ionic and assumes that the
ligands can be approximated by negative point charges.
More sophisticated models embrace covalency, and this
approach is described by ligand field theory (LFT) and Molecular orbital theory (MO). Ligand
field theory, introduced in 1935 and built from molecular orbital theory, can
handle a broader range of complexes and can explain complexes in which the
interactions are covalent. The chemical applications of group
theory can aid in the understanding of crystal or ligand field theory, by
allowing simple, symmetry based solutions to the formal equations.
Chemists tend to employ the simplest model required to
predict the properties of interest; for this reason, CFT has been a favorite
for the discussions when possible. MO and LF theories are more complicated, but
provide a more realistic perspective.
The electronic configuration of the complexes gives them
some important properties:
Color
Synthesis
of copper(II)-tetraphenylporphyrin, a metal complex, from tetraphenylporphyrin
and copper(II) acetate monohydrate.
Metal complexes often have spectacular colors caused by
electronic transitions by the absorption of light. For this reason they are
often applied as pigments. Most transitions that are related to colored metal
complexes are either d–d transitions or charge transfer bands. In
a d–d transition, an electron in a d orbital on the metal is excited by a
photon to another d orbital of higher energy. A charge transfer band entails
promotion of an electron from a metal-based orbital into an empty ligand-based
orbital (Metal-to-Ligand Charge Transfer or MLCT).
The converse also occurs: excitation of an electron in a ligand-based orbital
into an empty metal-based orbital (Ligand to Metal Charge Transfer or LMCT).
These phenomena can be observed with the aid of electronic spectroscopy; also
known as UV-Vis.
For simple compounds with high symmetry, the d–d transitions can be assigned
using Tanabe–Sugano diagrams. These assignments are
gaining increased support with computational chemistry.
Metal complexes that have unpaired electrons are magnetic.
Considering only monometallic complexes, unpaired electrons arise because the
complex has an odd number of electrons or because electron pairing is
destabilized. Thus, monomeric Ti(III) species have one "d-electron"
and must be (para)magnetic, regardless of the geometry or the
nature of the ligands. Ti(II), with two d-electrons, forms some complexes that
have two unpaired electrons and others with none. This effect is illustrated by
the compounds TiX2[(CH3)2PCH2CH2P(CH3)2]2:
when X = Cl,
the complex is paramagnetic (high-spin configuration), whereas when X = CH3,
it is diamagnetic (low-spin configuration). It is important to realize that
ligands provide an important means of adjusting the ground
state properties.
In bi- and polymetallic complexes, in which the
individual centers have an odd number of electrons or that are high-spin, the
situation is more complicated. If there is interaction (either direct or
through ligand) between the two (or more) metal centers, the electrons may
couple (antiferromagnetic coupling, resulting in a
diamagnetic compound), or they may enhance each other (ferromagnetic
coupling). When there is no interaction, the two (or more) individual metal
centers behave as if in two separate molecules.
Complexes show a
variety of possible reactivities:
·
Electron transfers
A common reaction between coordination complexes involving ligands are inner and outer sphere electron transfers.
They are two different mechanisms of electron
transfer redox
reactions, largely defined by the late Henry Taube.
In an inner sphere reaction, a ligand with two lone electron pairs acts as a bridging
ligand, a ligand to which both coordination centres can bond. Through
this, electrons are transferred from one centre to another.
·
(Degenerate) ligand
exchange
One important indicator of reactivity is the rate of degenerate exchange of
ligands. For example, the rate of interchange of coordinate water in [M(H2O)6]n+
complexes varies over 20 orders of magnitude. Complexes where the ligands are
released and rebound rapidly are classified as labile. Such labile complexes
can be quite stable thermodynamically. Typical labile metal complexes either
have low-charge (Na+), electrons in d-orbitals that are antibonding
with respect to the ligands (Zn2+), or lack covalency (Ln3+,
where Ln is any lanthanide). The lability of a metal complex also depends on
the high-spin vs. low-spin configurations when such is possible. Thus,
high-spin Fe(II) and Co(III) form labile complexes, whereas low-spin analogues
are inert. Cr(III) can exist only in the low-spin state (quartet), which is
inert because of its high formal oxidation state, absence of electrons in
orbitals that are M–L antibonding, plus some "ligand field
stabilization" associated with the d3 configuration.
·
Associative processes
Complexes that have unfilled or half-filled orbitals often show the
capability to react with substrates. Most substrates have a singlet
ground-state; that is, they have lone electron pairs (e.g., water, amines,
ethers), so these substrates need an empty orbital to be able to react with a
metal centre. Some substrates (e.g., molecular oxygen) have
a triplet ground state, which results that metals with half-filled orbitals
have a tendency to react with such substrates (it must be said that the dioxygen
molecule also has lone pairs, so it is also capable to react as a 'normal'
Lewis base).
If the ligands around the metal are carefully chosen, the
metal can aid in (stoichiometric or catalytic)
transformations of molecules or be used as a sensor.
Classification
Metal complexes, also known as
coordination compounds, include all metal compounds, aside from metal vapors, plasmas,
and alloys. The
study of "coordination chemistry" is the study of "inorganic
chemistry" of all alkali and alkaline earth metals, transition
metals, lanthanides,
actinides,
and metalloids.
Thus, coordination chemistry is the chemistry of the majority of the periodic
table. Metals and metal ions exist, in the condensed phases at least, only
surrounded by ligands.
The areas of coordination chemistry can
be classified according to the nature of the ligands, in broad terms:
·
Classical (or "Werner
Complexes"): Ligands in classical coordination chemistry bind to metals,
almost exclusively, via their "lone pairs"
of electrons residing on the main group atoms of the ligand. Typical ligands
are H2O, NH3, Cl−,
CN−,
en
Examples: [Co(EDTA)]−,
[Co(NH3)6]Cl3,
[Fe(C2O4)3]K3
·
Organometallic Chemistry: Ligands are
organic (alkenes, alkynes, alkyls) as well as "organic-like" ligands
such as phosphines, hydride, and CO.
Example: (C5H5)Fe(CO)2CH3
·
Bioinorganic Chemistry: Ligands are
those provided by nature, especially including the side chains of amino acids,
and many cofactors such as porphyrins.
Example: hemoglobin
Many natural ligands are "classical"
especially including water.
·
Cluster Chemistry: Ligands are all of
the above also include other metals as ligands.
Example Ru3(CO)12
·
In some cases there are combinations of
different fields:
Example: [Fe4S4(Scysteinyl)4]2−,
in which a cluster is embedded in a biologically active species.
Mineralogy,
materials science, and solid state chemistry – as they apply to
metal ions – are subsets of coordination chemistry in the sense that the metals
are surrounded by ligands. In many cases these ligands are oxides or sulfides,
but the metals are coordinated nonetheless, and the principles and guidelines
discussed below apply. In hydrates, at least some of the ligands are
water molecules. It is true that the focus of mineralogy, materials science,
and solid state chemistry differs from the usual focus of coordination or
inorganic chemistry. The former are concerned primarily with polymeric
structures, properties arising from a collective effects of many highly
interconnected metals. In contrast, coordination chemistry focuses on
reactivity and properties of complexes containing individual metal atoms or
small ensembles of metal atoms.
Traditional classifications of the kinds of isomer have
become archaic with the advent of modern structural chemistry. In the older
literature, one encounters:
·
Ionisation isomerism
describes the possible isomers arising from the exchange between the outer
sphere and inner sphere. This classification relies on an archaic
classification of the inner and outer sphere. In this classification, the
"outer sphere ligands," when ions in solution, may be switched with
"inner sphere ligands" to produce an isomer.
·
Solvation isomerism
occurs when an inner sphere ligand is replaced by a solvent molecule.
This classification is obsolete because it considers solvents as being distinct
from other ligands. Some of the problems are discussed under water of crystallization.
Naming
complexes
The basic procedure for naming a complex:
1.
When naming a complex ion, the ligands
are named before the metal ion.
2.
Write the names of the ligands in the
order,-neutral,negative,positive. If there are multiple ligands of the same
charge type, they are named in alphabetical order. (Numerical prefixes do not
affect the order.)
o
Multiple occurring monodentate ligands receive
a prefix according to the number of occurrences: di-, tri-, tetra-,
penta-, or hexa. Polydentate ligands (e.g.,
ethylenediamine, oxalate) receive bis-, tris-, tetrakis-,
etc.
o
Anions end in ido. This replaces
the final 'e' when the anion ends with '-ate', e.g. sulfate becomes sulfato.
It replaces 'ide': cyanide becomes cyanido.
o
Neutral ligands are given their usual
name, with some exceptions: NH3 becomes ammine; H2O
becomes aqua or aquo; CO becomes carbonyl; NO becomes nitrosyl.
3.
Write the name of the central atom/ion.
If the complex is an anion, the central atom's name will end in -ate,
and its Latin name will be used if available (except for mercury).
4.
If the central atom's oxidation state
needs to be specified (when it is one of several possible, or zero), write it
as a Roman numeral (or 0) in parentheses.
5.
Name cation then anion as separate words
(if applicable, as in last example)
Examples:
[NiCl4]2− → tetrachloridonickelate(II)
ion
[CuNH3Cl5]3− →
amminepentachloridocuprate(II) ion
[Cd(en)2(CN)2] →
dicyanidobis(ethylenediamine)cadmium(II)
[Co(NH3)5Cl]SO4 →
pentaamminechloridocobalt(III) sulfate
The
coordination number of ligands attached to more than one metal (bridging
ligands) is indicated by a subscript to the Greek symbol μ placed before
the ligand name. Thus the dimer of aluminium
trichloride is described by Al2Cl4(μ2-Cl)2.