2P32 – Principles of Inorganic Chemistry
Dr. M. Pilkington
Lecture 30 – Groups 14 and 15
1. Group 14 – Carbon oxides and Carbonates
2. Group 14 -Silicates
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Isolated ions
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Rings and chains
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Sheet
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3-D Aluminosilicates
3. Group 15 – N, P, As, Sb, Bi
1. Group 14 - Carbon Oxides and Carbonates
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Carbon forms two well-known oxides, carbon monoxide, CO, and carbon
dioxide, CO2. In addition, it also forms carbon suboxide, C3O2.
CO- Carbon Monoxide
CO2- Carbon Dioxide
C3O2- Carbon Suboxide
CO32HCO3H2CO3
C
O
O
O
C
C
O
C
C O
Carbonates
Carbon monoxide.
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Carbon monoxide is produced when graphite (one of the naturally occurring
forms of elemental carbon) is heated or burned in a limited amount of oxygen.
The reaction of steam with red-hot coke also produces carbon monoxide along
with hydrogen gas (H2). (Coke is the impure carbon residue resulting from the
burning of coal.) This mixture of CO and H2 is called water gas and is used as
an industrial fuel.
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In the laboratory, carbon monoxide is prepared by heating formic acid, HCOOH,
or oxalic acid, H2C2O4, with concentrated sulfuric acid, H2SO4. The sulfuric acid
removes the elements of water from the formic or oxalic acid and absorbs the
water produced. Because carbon monoxide burns readily in oxygen to produce
carbon dioxide,
2CO + O2
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2CO2,
It is useful as a gaseous fuel. It is also useful as a metallurgical reducing agent
because at high temperatures it reduces many metal oxides to the elemental
metal. For example, copper(II) oxide, CuO, and iron(III) oxide, Fe2O3, are both
reduced to the metal by carbon monoxide.
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Carbon monoxide is an extremely dangerous poison. Because it is an odorless and
tasteless gas, it gives no warning of its presence. It binds to the hemoglobin in
blood to form a compound that is so stable that it cannot be broken down by
body processes. When the hemoglobin is combined with carbon monoxide, it
cannot combine with oxygen; this destroys the ability of hemoglobin to carry
essential oxygen to all parts of the body. Suffocation can occur if sufficient
amounts of carbon monoxide are present to form complexes with the hemoglobin.
Carbon dioxide
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Carbon dioxide is produced when any form of carbon or almost any carbon
compound is burned in an excess of oxygen. Many metal carbonates liberate CO2
when they are heated. For example, calcium carbonate (CaCO3) produces carbon
dioxide and calcium oxide (CaO).
CaCO3 + heat
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CO2 + CaO
The fermentation of glucose (a sugar) during the preparation of ethanol, the
alcohol found in beverages such as beer and wine, produces large quantities of
CO2 as a by-product.
C6H12O6
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2C2H5OH+ 2CO2
In the laboratory CO2 can be prepared by adding a metal carbonate to an
aqueous acid; e.g.,
CaCO3 + 2H3O+
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Ca²+ 3H2O+ CO2.
Carbon dioxide is a colorless and essentially odorless gas that is 1.5 times as
dense as air. It is not toxic, although a large concentration could result in
suffocation simply by causing a lack of oxygen in the body.
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All carbonated beverages contain dissolved CO2; hence the name carbonated. One
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litre (1.06 quarts) of water at 20° C dissolves 0.9 litre of CO2 at one atmosphere,
forming carbonic acid (H2CO3), which has a mildly acidic (sour) taste.
Solid CO2 sublimes at normal atmospheric pressure. Thus, solid CO2, called dry
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ice, is a valuable refrigerant that is always free of the liquid form. Carbon
dioxide is also used as a fire extinguisher, because most substances do not burn
in it, and it is readily available and inexpensive. Air containing as little as 2.5
percent CO2 extinguishes a flame.
The Earth's atmosphere contains approximately 0.04 percent carbon dioxide by
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volume and serves as a huge reservoir of this compound. The carbon dioxide
content of the atmosphere has significantly increased in the last several years
largely because of the burning of fossil fuels. A so-called greenhouse effect can
result from increased carbon dioxide and water vapor in the atmosphere. These
gases allow visible light from the Sun to penetrate to the Earth's surface, where
it is absorbed and reradiated as infrared radiation.
Carbon suboxide.
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Carbon suboxide, C3O2, is a foul-smelling, lachrymatory gas produced by the
dehydration of malonic acid, CH2(COOH)2, with P4O10 in a vacuum at 140° to 150°
C. Carbon suboxide is a linear, symmetrical molecule whose structure can be
represented as O=C=C=C=O. At 25° C the compound is unstable and polymerizes
to highly coloured solid products, but it is a stable molecule at -78° C. Under the
influence of ultraviolet light (in the process known as photolysis), C3O2 decomposes to the very reactive molecule ketene, C2O. Since carbon suboxide is
the acid anhydride of malonic acid, it reacts slowly with water to produce that
acid.
Silcon oxides
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Compared to Carbon there is only one.
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SiO2 – where every Silicon is bonded to 4 Oxygen’s, every Oxygen is bonded to
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Si(OH)4
two Silicon’s. It is a network structure in sharp contrast to the CO2 molecule.
SiO2 (crystalline structure = Quartz)
-H2O
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The silicates are the largest, the most interesting and the most complicated
class of minerals than any other minerals. Approximately 30% of all minerals are
silicates and some geologists estimate that 90% of the Earth's crust is made up
of silicates, SiO44- based material. Thus, oxygen and silicon are the two most
abundant elements in the earth's crust.
Silicates is based on the basic chemical unit SiO44-, tetrahedron shaped anionic
group. The central silicon ion has a charge of positive four while each oxygen has
a charge of negative two (-2) and thus each silicon-oxygen bond is equal to one
half (½ ) the total bond energy of oxygen. This condition leaves the oxygens with
the option of bonding to another silicon ion and therefore linking one SiO44tetrahedron to another.
In the extreme case, the tetrahedra are arranged in a
regular, orderly fashion forming a three-dimensional
network. Quartz is such a structure (see the diagram), and
its formula is SiO2. If silica in the molten state is cooled
very slowly it crystallizes at the freezing point. But if
molten silica is cooled more rapidly, the resulting solid is a
disorderly arrangement which is called glass, often also
called quartz.
2.
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Silicates and Silica
The chemistry of silicon is dominated by silicon-oxygen bonds.
The basic unit of all of the silicates is the SiO44- tetrahedron.
The various types of silicates are characterized by the sharing of from zero
to all four of the oxygen’s in this unit.
The charge on each repeating unit is determined by recalling that the
oxidation state of silicon is +4 while oxygen is as expected, -2.
How to Classify Silicates
1.
2.
3.
The orthosilicates feature a discrete, self-contained SiO44- unit (a). The
cations are some other metals. For example, in olivine, (Fe, Mg)2SiO4, the
cations are either Fe2+ or Mg2 .
The rare pyrosilicates or disilicates are characterized by the sharing of one
oxygen of each SiO44- unit, to produce the Si2O76- ion (b). For example,
thortveitite, Sc2Si2O7 is a pyrosilicate.
Cyclic silicates find two oxygen’s of the SiO44- being shared to produce sixmembered Si3O96- (c) and twelve membered Si4O1812- (d). The precious stone
beryl Be3Al2(SiO3)6 contain six-silica rings.
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The mineral Beryl Be3Al2[Si16O18] contains the cyclic ion (d)
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When Cr3+ replaces a little Al3+ , we have Emerald (green).
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When Fe3+ replaces the Al3+ then we have Aquamarine.
4.
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5.
Single Chain Silicates - Find two oxygen’s of the SiO44- unit being shared but
not closing to form a ring. The repeating unit of these infinite chains is SiO32shown below (e). Spodumene, LiAl(SiO3)2 is one of the few important lithium
ores.
Double chain silicates, (f) find some SiO44- units sharing two oxygen’s while
others share three. The repeating unit is Si4O116-
Sheet silicates find the bulk of the tetrahedra sharing three oxygens to form
a Si2O52- repeating unit (g). Talc or soapstone is a sheet silicate as is the
mineral petalite from which Arfwedson first isolated lithium.
The arrangement of sheets in brucite,
Mg(OH)2, in which the sheets consist of corner
sharing octahedrons of Mg(OH)6. In chlorite,
there are two types of sheets. Half of the
sheets are the same as those of brucite, but
half of the brucite-sheets are sandwiched
between sheets of silicates. The talc consists
of only the sandwiched sheets.
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3. Group 15 (5A)
N, P, As (Arsenic), Sb (Antimony), Bi(Bismuth)
Metals
Non Metals
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Known as the pnicogens a name that means “choking producers” from the Greek.
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There are some uniform properties within the group, but there is still a great
deal of diversity among the elements of the group.
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P, As and Sb have strong connections to Alchemy (magicians-chemists in the
middle ages).
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Antimony comes from the Latin “stilbium”, formally used for the element. It may
also have Arabic origin. Early application of Sb was as a vomiting inducer. Sb2S3
was used by the ancients as a cosmetic to darken and beautify eyebrows.
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Arsenic comes from the Greek word “arsenikon” which is an adaptation of the
Persion word for “yellow ointment”, a common sulfide ore of the element. Albert
the Great a 13th Century German scholar is credited with its discovery because
of the clear descriptions he gave it in his writings dated to 1250.
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The poisonous nature of arsenic compounds has been long known. Both accidental
and intentional poisonings have been well documented.
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Phosphorus was first discovered by burning human urine, then it was discovered
in human bones and in the 1800 the demand for phosphorus in England was so
great that battlefields were being combed for human remains to make the
matches.
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Phosphorus burns with a bright white flame hence its name comes from the
Greek word meaning “light bearing”.
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The first isolation of bismuth is unclear. It is most likely derived from the
Latin word for “white metal”.
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It is a white crystalline metal with a pinkish tinge.
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Early uses for bismuth included addition to tin to increase its hardness and
brilliancy.
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Alloys are low melting and are used today for safety devices in fire-detection
and sprinkler systems, as well as fuses and relief valves.
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Nitrogen compounds were well known before the free element was isolated.
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Ammonium salts have been characterized since the fifth century.
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The element was given the name “nitrogen producer” after it was found to be
the component of nitric acid and nitrates.
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N is still manufactures by liquefying and fractionally distilling common air.
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It is the most abundant uncombined element known; the atmosphere contains 4
trillion tons of nitrogen.
The differences between N and P are due to:
Their abilities to π-bond
N-pπ bonding (no availability of 3d orbitals)
P-dπ bonding i.e the availability of d-orbitals for P means it can have sp3d and sp3d2
hybrids.
Allotropes
a.
N2
pπ-pπ bonding, the N are sp hybridized and there is only one allotrope.
b.
P2 – several allotropes. P2 does not exist at room temperature
:N
N:
:P
P:
pπ-pπ bonding good overlap - 1 allotrope
requires pπ-pπ bonding - this is weak in elements
beyond the second row, the p orbitals are further
appart because the atoms are larger so this is weak.
P4 - stable allotrope at room temperature - White phosphorous
burns spontaneously in air-stored in water
P
P
P
P
sp3 hybridized - the angles for sp3 are 1090
although the angles are much less in the triangle
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Allotropes of P
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The discovery of Phosphorous
Phosphorus is the 15th element on the
periodic table. Located between silicon
and sulfur, the element is widely known
for its many different allotropes and
association with fire. You see unlike
iron or sodium, which only really exist
in one form, phosphorus exists in three
different forms. These different
forms, or more properly speaking
“allotropes,” are all the same element,
but they differ in their structure and
atomic arrangement.
These slightly
different arrangements cause the
chemical and physical properties of the
allotropes to differ ever so slightly.
If we cut white phosphorous the cut line turns red:
P
P
P
P
P
P
P
P
P
P
P
P
Red P - alot of the bond strain is released from P4 as there are 3-memebered rings
This compound does not burn in air spontaneously (unless flamed).
Black P is the most stable allotrope – 6 membered chair rings joined together
to form sheets.
ƒ Insoluble in organic solvents and water.
ƒ Non-Volatile.
The three main allotropes of phosphorus are
white phosphorus, red phosphorus, and black
phosphorus. Yellow and violet phosphorus
are sometimes mentioned, but those really
aren’t true allotropes. They are just
combinations of two different allotropes
into a homogenous mixture. (Yellow = White
+ Red Phosphorus, and Violet = Black + Red
Phosphorus).
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Nitrogen
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It has the most number of oxidation states of any of the main group elements
(-3, -2, -1, 0, 1, 2, 3, 4, 5) i.e 9 oxidation states, don’t forget Mn – transition
metal with 11 oxidation states (-3 to +7).
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Examples include the following:
1.
NH3 – (-3) ammonia
2.
N2H4 – (-2) hydrazine
3.
NH2OH – (-1) hydroxylamine
4.
N2- (O) nitrogen
5.
N2O – (+1)
6.
NO – (+2)
7.
N2O3 – (+3)
8.
NO2 – (+4)
9.
N2O5 – (+5)
Note that NO and NO2 have an odd number of electros so no Lewis
structures.
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The chemistry is characterized by a variety of single and double bonds
involving both localized and delocalized pπ-pπ interactions.
Phosphorous
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The phosphorous oxides and corresponding acids are mildly acidic.
Common oxidation states for P compounds are -3 (PH3, phosphine) +3 (P4O6) and
+5 (P4O10).
For the oxides and acids both the +3 and the +5 phosphorous oxides are known.
Normally we would expect these oxides to have formulas of P2O3 and P2O5 and
indeed these compounds are referred to as phosphorous trioxide and
phosphorous pentoxide.
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The molecular formulas however are P4O6 and P4O10 respectively.
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White phosphorous P4 and the two oxides are structurally related.
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P4 is a tetrahedron of P atoms and P4O6 has six oxygen atoms bridging the edges
of the tetrahedron. In P4O10 four terminal P=O bonds (of the pπ-dπ type) have
been added. (Remember empty d orbitals accept electron density from filled
oxygen p orbitals).
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Acids
N HNO2 nitrous acid
HNO3 nitric acid
oxo group
O
N
O H
nitrous acid (less O)
O
N
O H
nitric acid
O
two oxo groups - stronger acid than HNO2
P
H3PO3 phosphorous acid
H3PO4 phosphoric acid
one oxo group
O
H
P
HO
OH
O
P
one oxo group
OH
OH
OH
H3PO4
H3PO3
similar Ka's
H3PO2 –hypophosphorous acid
pπ-dπ bonding
O
H
P
one oxo group
OH
H
simlar Ka to the previous acids
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Size differences between N and P is responsible for the different numbers of
atoms bonded to N and P
N – 3 bonded atoms
P – 4 bonded atoms (uses d orbitals to accept the electrons donated).
Remember that P does not obey the Octet rule when you draw your Lewis
structures and that the chemistry of phosphorous is dominated by element to
element single bonds but also by the availability of 3d orbitals to form pπ-dπ
double bonds with a variety of other atoms such as oxygen, nitrogen and even
sulfur. d-Orbital participation results in exzpanded octets found in compound
such as PF5, SbCl5, X3P=O (where X = F, Cl, Br, I).
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Phosphoric acid (a), two phosphoric acid
moleules split out a molecule of water to
form pyrophosphoric acid H4P2O2.
Pyrophosphate (c)
Tricyclophosphate (d)
Tetracyclophosphate (e)
Polyphosphate (f).
Phosphorous acid H3PO3, two phosphorous
Acid molecules split out a water to form
Pyrophosphorous acid H4P2O5.
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