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Chemistry 2810 Lecture Notes
8.
Dr. R. T. Boeré
Page 125
p-Block Elements
We have now arrived in our journey across the periodic table to that section which has the most variation in chemical
properties. This is reflected in the large range of electronegativities that is observed for this range of elements, from a low of 1.61
for Al to a high of 3.98 for F (and even higher for the noble gases, although Pauling values are not available for these systems in
that they don't all form chemical bonds.
8.1
Metals, metalloids and non-metals
B C N O
F
Al Si P S Cl
Ga Ge As Se Br
In Sn Sb Te I
Tl Pb Bi Po At
He
Ne
Ar
Kr
Xe
Rn
One example of the variation in chemical properties through the p-block is in element properties. Thus Al, Ga, In, Tl, Sn, Pb,
Sb, Bi and Po are metals, although in their highest oxidation states they have some metalloid properties. Note that all the
metalloids (with the possible exception of Be) are found in the p-block. Thus B, Si, Ge, As, Te and At have these very important
properties which are of the utmost technological importance. Si and Ge are the two most important materials for all types of
semiconductors used in the electronic industries. Thus electrical conductivity is another physical property whereby the metals,
metalloids and non-metals can be distinguished: the metals tend to be good conductors, the metalloids intrinsic
semiconductors, and the non-metals insulators. This trend is seen clearly within a single group such as 14, where carbon
(diamond) is an excellent insulator, Si and Ge are semiconductors, and Sn and Pb are conductors.
We will not say a lot about the metallic p-block elements, except to mention that In and especially Tl, Sn and especially Pb,
Sb and especially Bi can exist in two oxidation states. Thus we have Tl+ and Tl3+. By contrast, Al and Ga are only found as the
element or as the 3+ ion. Why is this? Consider the electron configurations of the thallium ions already mentioned:
Tl+
[Xe]4f145d 106s 2
Tl3+
[Xe]4f145d 10
These differ by the presence and absence of the 6s electrons.
This is also the origin of the 2+ difference in oxidation states for the other metals, i.e. Sn 2+ and Sn 4+, etc. If we remember that
the 4f electrons are especially poor at shielding, we recognize that these elements will experience an unusually high effective
nuclear charge. This has the effect of causing the 6s electrons to be particularly strongly bound to the nucleus, and therefore
harder to oxidize. The full explanation of this behavior involves relativistic behavior of the electrons in these heavy atoms, and
we will not treat it further in this course.
Chemistry 2810 Lecture Notes
8.2
Dr. R. T. Boeré
Page 126
Group 13 elements - Boron
B
Al
Ga
In
Tl
No information at the current time.
8.3
Group 14 elements; silicones; silicate minerals
C
Si
Ge
Sn
Pb
Element
C
Si
Ge
Sn
Pb
Electronic structure
[He]2s 22p 2
[Ne]3s 23p 2
[Ar]3d 104s 24p 2
[Kr]4d 105s 25p 2
[Xe]4f146s 25d 106p 2
mp, °C
3550
1410
937
232
328
1st I.E., kJ mo l-1
1086
786.3
760
708.2
715.3
Electronegativity
2.50
1.74
2.02
1.72
1.8
Covalent radius
0.77
1.17
1.22
1.40
1.44
The elemental forms of carbon are graphite (thermodynamically the most stable at room temperature) and diamond. These
two forms could not be more different! Graphite is a soft, black, conducting solid, which is used in many commercial
applications, including their use as electrodes in many important electrochemical processes (e.g. the 1.5 V dry cell, electrolysis of
bauxite/cryolite in the manufacture of aluminum.) Diamond is a clear colourless insulator, and is the hardest substances known.
These differences are due entirely to the physical structure of these important structures. We have already dealt with the
diamond structure and the graphite structure.
Diamond
Graphite
In diamond, all the carbon atoms are sp 3 hybridized, and therefore four coordinate. The whole crystal is one large molecule.
This is what gives it its immense hardness.
Elemental silicon and germanium have the diamond structure. It is a fascinating fact that silicon, with its diamond structure,
is actually a semiconductor! So is germanium. Tin and lead have close-packed metallic structures, but tin has an allotrope called
gray tin which has the diamond structure and is stable below 18 C. This form is much more brittle than the close-packed material.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 127
Direct combination of carbon with metallic elements results in the formation of metal carbides. Probably the most famous of
these is calcium carbide, CaC2. This material is an example of a saline carbide, consisting of Ca 2+ and C22- ions. It is a colourless
solid when pure, but is usually discoloured or gray. Its crystal structure is based on that of NaCl, except that at each Cl site, a C2
unit is located. All the C2 molecules are oriented in the same direction, and the expanded unit cell is tetragonal. The figures
below illustrate the CaC2 structure (left) and the BaC2 structure (right). These belong
to different crystal classes, but the basic structural units are the same, with the
c
cylindrical C22– ions aligned vertically.
A hint about the chemical structure of the carbide ion in calcium carbide is
provided by the reaction the ion undergoes with water:
CaC2(s) + H2O(l)
→
b
Ca(OH)2(aq) + H-CsC-H(g)
a
This result, the generation of acetylene, suggests that the carbide ion in saline
carbides has a triple bond. Try and write a Lewis structure for C22- using this hint. In
the early years of this century, this reaction was used in a practical application as a source of fuel for lamps. Early motorcars
invariably had carbide headlamps. It is from the manufacture of this product that the giant multinational chemical conglomerate
Union Carbide gets its name.
Consider the Frost diagram for Group 14 (here drawn in acid solution only for simplicity). Different diagrams need to be drawn
for basic solution, and different ones again for the common organic representatives of the +2, 0 and –2 oxidation states.
Frost Diagram for Group 14 in 1 M Acid
3
PbO2
2
GeH4
nE (V)
1
0
SiH4
-IV
CH4
-1
CO
C
Si
-II
0
Ge
Sn
Pb
CO2
GeO
II
Pb2+SnO
IV
GeO2
SnO2
C
Si
-2
Ge
Sn
-3
Pb
SiO2
-4
Oxidation State
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 128
They have the general valence configuration ns2np 2, and thus commonly have the oxidation state of 4. Carbon has
essentially no chemistry in the divalent state. Note that CO is unstable with respect to disproportionation to C and CO2. This is,
however, a kinetically slow reaction, and CO can be stored as a compressed gas in cylinders for years without apparent
decomposition. However, CO can be burnt to CO2, and this is in fact the common way of purifying air supplies of toxic carbon
monoxide (e.g. in close air supplies such as found in submarines or on space craft.) Certain carbenes, especially CF2, are stable
but very reactive species. These are true divalent carbon compounds, but they are extremely rare. Silicon and germanium are
also primarily tetravalent in their compounds. But for the heaviest members of the group, Sn and Pb, the oxidation state +2 is
quite common, (electron configuration s2d 10). The stability of the tetravalent state decreases down the group as well, so that
while CCl4 and SiCl4 are extremely easy to make from the elements and chlorine, PbCl4 can only be made under forcing conditions,
and decomposes readily to PbCl2.
You may be surprised to see carbon on the agenda of an inorganic chemistry course! In fact, though the vast majority of
carbon compounds are hydrocarbon derivatives (and hence the proper subject of organic chemistry), there is also and important
inorganic chemistry of carbon. Remember that these artificial sub-disciplines of our convenience make absolutely no impression
on Nature! In this laboratory we consider two important inorganic carbon compounds, carbon dioxide and calcium carbide.
O
O
O
C O
carbon monoxide
O
O
O C O
carbon dioxide
O
O
O
O C C C O
carbon suboxide
O
mellitic anhydride
Carbon has four thermodynamically stable oxides: CO, CO2, C3O2 and C12O9. The latter is the anhydride of mellitic acid,
hence is best considered an organic derivative. Carbon suboxide is and evil-smelling gas, which is stable at low temperatures,
but polymerizes when warmed to room temperature. The two important oxides are the monoxide and dioxide.
CO is formed when carbon is burned in a deficiency of oxy gen. It is a significant component of automobile exhaust fumes.
Industrially it is produced on a huge scale by the reaction of steam with coke to give hydrogen and carbon monoxide (syngas):
C(s) + H2O(g)
→
CO(g) + H2(g)
These products form the feedstocks of some petrochemical processes. They can also be combined with the water gas shift
reaction to maximize hydrogen production and produce CO2:
CO(g) + H2O(g)
→
CO2 + H2
Carbon monoxide is very toxic. It replaces oxygen on the O2carrier complex hemoglobin, blocking the transfer of oxygen from the
lungs to the rest of the body. CO also reacts with many metals to form organometallic compounds in which carbon is bound to
transition metals. An example is the compound Ni(CO)4, which is an intermediate in the Mond process for the purification of
nickel.
Carbon dioxide is the main product of the combustion of fossil fuels, as well as many natural processes such as respiration.
In the laboratory it is most commonly prepared by the controlled addition of HCl(aq) to CaCO3 in the form of marble chips. CO2
is denser than air, so it can be collected directly in containers by upward displacement of the air.
CO2 does not support comb ustion, and is used as the compressed liquid in many fire extinguishers. However, although it is
one of the most versatile extinguishing materials, it cannot be used on all fires. An example is afforded in this laboratory by
demonstrating that magnesium "burns" in a CO2 environment. This is the exact reverse of obtaining the element by coke
reduction of magnesium oxide. The Ellingham diagram on p. 230 of reference 1 is of interest in this connection (see also Exercise
8.1 on p. 231.)
CO2 is present in the atmosphere in large quantities, and it is feared by many that the huge scale of human activity in the
past century has raised the total level of CO2 in the atmosphere. This may lead to global warming under the so-called
"greenhouse effect." CO2 is als o intimately connected with plant and animal life. Photosynthesis in plants is the only known
way to get usable forms of carbon compounds back from CO2.
CO2 dissolves in water to give a solution of carbonic acid:
CO2 + H2O
⇔
HCO3- + H+
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 129
This is why water in contact with the atmosphere is usually found to be slightly acidic. Carbon dioxide has a very high solubility
in water. The true pKa of carbonic acid is -3.7, a figure which takes into consideration that much of the CO2 in water is simply
solvated, and is not in the form of H2CO3. Thus solutions of CO2are less acidic than expected based on the quantity of
CO2dissolved.
8.5
Group 15 elements; fertilizers and bio-phosphates
N
P
As
Sb
Bi
The electronic configuration and the ionization energies of the element in Group 15 are shown in the following table:
Element Electronic structure
mp, °C
1st I.E., kJ mol-1
Electronegativity
2
3
N
[He]2s 2p
-210
1400
3.07
P
[Ne]3s 23p 3
44*
1012
2.06
As
[Ar]3d 104s 24p 3
817
947
2.20
Sb
[Kr]4d 105s 25p 3
631
834
1.82
Bi
[Xe]4f146s 25d 106p 3
271
703
1.67
* White phosphorus; red P melts at 990 C under pressure; black P does not melt.
Covalent radius
0.74
1.10
1.21
1.41
1.52
Observe that all the elements have only three electrons in their outermost orbitals; they all have five valence electrons.
Loss of three electrons leads to an ns 2 configuration; loss of five electrons gives a noble-gas cation, but such 5+ ions are not
stable by themselves. On the other hand, it would be possible for the atoms to fill their outer shells by gaining three electrons.
The theoretical range of the oxidation numbers of these atoms is therefore from a maximum value of +5 to a minimum value of -3.
The nitride ion N3- exists in many compounds with metals, but these are highly polarized ionic compounds because of the high
charge on the nitrogen. Much more common are compounds in which the atoms share electrons to form covalent bonds. Thus,
we would expect nitrogen atoms to share their outer electrons with atoms of more electronegative elements, in which case the
maximum oxidation state would be +5. Actually, nitrogen exhibits all the oxidation states from +5 to -3.
As we go down the group, the atoms get larger in size and become more metallic in nature. Size of atom and metallic (or
nonmetallic) nature are closely associated. The nitrogen atoms are the smallest ones in this family, and it is difficult to dislodge
their outermost electrons. The high ionization energy of nitrogen confirms the fact that the outermost electrons are very tightly
held and we therefore conclude that nitrogen behaves as a nonmetal. Bismuth atoms are the largest in the group, and therefore
their outer electrons are easiest to dislodge. The ionization energy of bismuth is only 703 kJ mol-1, compared to 496 for Na and
731 for Ag. It therefore follows that bismuth is the most metallic element of the group. At the centre of the group, elements are
on the borderline between metallic and nonmetallic properties. The hydroxide of antimony, for example, can be acidic or basic
and, as such, is call amphoteric. Arsenic and antimo ny are usually classed as metalloids.
Nitrogen
The elemental form of nitrogen is relatively inert. Thus N2 is commonly used as an inert atmosphere gas. There are
limitations to this, and for example nitrogen reacts with lithium to form a coating of Li3N on the exposed surface of the metal, and
some transition metal complexes take up N2. Under stringent conditions, argon is to be preferred over nitrogen as an inert gas.
Liquefied nitrogen (bp -196 C) is an extremely useful coolant, in part because it evaporates to harmless nitrogen gas (except for
the danger of asphyxiation in enclosed spaces.)
A nitrogen molecule is diatomic, N2. It can be represented by the Lewis structure :NsN: which implies a nitrogen-nitrogen
triple bond. A good deal of energy is needed to break these bonds, and they must be broken before nitrogen atoms can react
with atoms of other elements. This is one of the reasons why nitrogen molecules are so inactive. Conversely, the large bond
energy of N2 is one of the thermodynamic driving forces behind the explosive nature of many nitrogen compounds. The
formation of N2 from a compound simultaneously releases a lot of heat and gas volume.
Nitrogen constitutes about 80% of the atmosphere. This high proportion of nitrogen might lead us to think of nitrogen as a
plentiful element. But such is not the case. Indeed, nitrogen is a relatively scarce element, only about one-third as abundant as
carbon. Then how can we explain the high proportion of nitrogen in air? This is due to its chemical inactivity. In nature it forms
relatively few compounds. Practically the only simple nitrogen compound in nature is Chile saltpetre (NaNO3). Of course
considerable quantities of nitrogen are involved in living systems and their remains (e.g. mineral oil.)
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 130
The inactivity of nitrogen has posed a crucial problem for the chemist. All plant and animal life depends on nitrogen
compounds. Therefore, nitrogen must be made to combine, for otherwise human life would be in jeopardy. Plant and animal cells
consist of proteins which are highly complex compounds of carbon, hydrogen, oxygen and nitrogen (about 17% nitrogen). A
growing plant takes nitrogen from the minerals in soil to make proteins. The soil, steadily depleted of its nitrogen minerals, must
be replenished, otherwise plants would starve and so would the animals which depend on plants.
The challenge to the chemist was to learn how to make suitable nitrogen compounds which could serve as plant food. The
way the chemist has met this challenge is one of the major scientific achievements of the present century.
The preparation of nitrogen
Nitrogen is needed to make various nitrogen compounds on an industrial scale, and for this purpose it is obtained by the
fractional distillation of liquid air. Nitrogen is more volatile than oxygen (the boiling points are -196 C for nitrogen and -183 C for
oxygen) and therefore evaporates in the first fraction.
In the laboratory, an easier way to prepare nitrogen from air is simply to remove the more active oxygen by making it
combine with some substance. For example, if air is passed over hot copper, the oxygen combines with it to form copper (II)
oxide:
Cu + O2 (+N2) → 2 CuO + (+N2)
If all the oxygen is used up, the residual gas is about 99% nitrogen; it contains small proportions of the inactive gases carbon
dioxide and argon, and water vapour.
Pure nitrogen (or chemical nitrogen) can be prepared by the decomposition of ammonium nitrite (NH4NO2). This compound
readily decomposes at a low temp erature to give nitrogen as the only gaseous product:
NH4NO2
→
N2 + 2 H2O
This is an example of a comproportionation reaction.
Nitric acid
Nitric acid is one of the most important acids in the chemical industry; it is used in the manufacture of fertilizers, drugs,
dyes, and plastics. In industry, it is made by the catalytic oxidation of ammonia, the so-called Ostwald process, described by the
following 3 equations. The Ostwald process is a cycle, in which the NO produced in step 3 is automatically sent back into the
catalytic converter used in step 2.
4 NH3 + 5 O2 → 4 NO + 6 H2O
2 NO + O2 → 2 NO2
3 NO2 + H2O → 2 HNO3 + NO
The Lewis structure of nitric acid is:
..
O..
..
H-O-N
..
..
..O .
.
..
O..
..
H-O-N
..
..
..O .
.
which shows that the nitrogen atom is in its highest oxidation state of +5.
Pure nitric acid is a colourless liquid that boils at 86°C. It decomposes in sunlight, or when heated, forming nitrogen dioxide
(NO2) which turns the solution brown.
4 HNO3 → 2 H2O + O2 + 4 NO2
Water is usually added to retard this decomposition and, as a result, ordinary concentrated nitric acid contains about 68% acid and
is about 15 molar.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 131
Properties of Nitric Acid
Nitric acid is a strong acid and can be thought of as completely ionized in dilute solution:
HNO3 → H+ + NO3Thus, with metallic hydroxides, carbonates, or oxides the reactions are the same as those of dilute hydrochloric acid or dilute sulfuric
acid. For example,
Na + + OH- + H+ + NO3- → Na + + NO3- + H2O
With metals, however, nitric acid does not yield hydrogen as is the case with the other acids, because nitric acid is a
powerful oxidizing agent. You will recall that dilute hydrochloric acid is an oxidizer because it contains hydrogen ion. Or,
expressed another way, the following is a reducing reaction:
2 H+ + 2 e- → H2
In nitric acid, however, both hydrogen ion and nitrate ion (NO3-) are good oxidizers. We would expect nitrate ion to be an
oxidizer because nitrogen is in its highest oxidation state of +5. If in any reaction the oxidation state of nitrogen is reduced, then
of necessity N+5 (or the nitrate ion) is an oxidizer. That is why free hydrogen is not usually a product when dilute nitric acid
reacts with a metal. Indeed, the products of such a reaction depend upon the concentration of hydrogen ions and nitrate ions,
the temperature of the reaction, and the activity of the reducing agent. In short, these reactions can be highly complicated,
particularly with active metals.
Since dilute hydrochloric and sulfuric acids owe their oxidizing power solely to hydrogen ions, they will react only with
metals above hydrogen in the emf series. But this is not the case with dilute nitric acid, which will react with all metals above
hydrogen and most of the metals below hydrogen. Let us consider the reactions of nitric acid with copper.
(a) Concentrated nitric acid and copper
As stated earlier, concentrated nitric acid decomposes rapidly at its boiling point to give nitrogen dioxide (NO2) as one of
the products. This same decomposition product is formed when copper reacts with the concentrated acid, the reaction being
Cu + 4 HNO3 à Cu(NO3)2 + 2 H2O + 2 NO2
(b) Dilute nitric acid and copper
If dilute nitric acid is used, a colourless gas, nitric oxide (NO), is formed rather than the dark brown nitrogen dioxide.
3 Cu + 8 HNO3 à 3 Cu(NO3)2 + 4 H2O + 2 NO
To summarize: when concentrated nitric acid is the oxidizer, nitrogen dioxide is a product; when dilute nitric acid is the oxidizer,
nitric oxide is a product. Notice also that in the more dilute acid, the nitrogen is reduced to a lower oxidation state! This complex
behaviour can be summarized using the Latimer diagrams found in Ref.1, appendix 4.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 132
Nitrates, the salts of nitric acid
An unusual but characteristic feature of nitrates is that their salts are all soluble in water. It is well to bear this in mind if a
solution of a metallic ion is needed. Nitrates, like nitric acid, decompose when heated. Nitrates of the heavy metals decompose
in an analogous manner to nitric acid; instead of O2, they yield the metallic oxide:
2 Pb(NO3)2 à 2 PbO + 4 NO2 + O2
Nitrates of the alkali metals, however, are more strongly bound in their crystal lattices and, in consequence, are more difficult to
decompose. When heated, they lose only oxygen, not nitrogen dioxide, to form the corresponding nitrite. For example:
2 NaNO3 à 2 NaNO2 + O2
It should be mentioned that ammonium nitrate is exceptional in the way it decomposes.
NH4NO3 à N2O + 2 H2O
This is because the ammonium ion is itself capable of being oxidized, and takes part in the decomposition to yield nitrous oxide, N2O.
This gas is used as an anaesthetic, and is sometimes referred to as laughing gas.
Phosphorus
Phosphorus is much mo re abundant than nitrogen. It occurs mainly as phosphate rock which contains a high percentage of calcium
phosphate Ca 3(PO4)3. Phosphorus, like sulfur, occurs in allotropic forms, the two common ones being white (or yellow) phosphorus
and red phosphorus.
Allotropes of Phosphorus
White phosphorus is a soft, waxlike solid, exceedingly poisonous and very reactive chemically. It ignites spontaneously in
air. White phosphorus is soluble in carbon disulfide but insoluble in water, and so it can be stored under water. Red
phosphorus does not oxidize rapidly at room temperature, although it burns very readily if it is heated sufficiently. Red
phosphorus is much less poisonous than yellow phosphorus, and it is insoluble in carbon disulfide. These marked differences in
properties between the allotropes suggest a difference in molecular structure.
A phosphorus atom has five valence electrons and its electronic symbol is usually written as P The molecular weight of
phosphorus, as determined experimentally from its vapour density, is 124. Therefore, phosphorus vapour consists of P4
molecules. These molecules have an unusual tetrahedral shape, in which the four atoms are located at the corners of a regular
tetrahedron. (see below) In contrast to nitrogen which forms N2 in the gas phase with a triple bond, phosphorus forms P4 which
has four strained single bonds. (At least this is true below 800 C; above this temperature the vapour consists of P2 molecules.)
This is consistent with the known weaker double bond versus single bond strength on going down the periodic table. Indeed,
compounds possessing stable P=P bonds were unknown until 1981.
Phosphorus vapour condenses to a liquid at 280 C and then to a solid at 44°C. The attractive forces between the molecules
are van der Waals forces. If the solid or liquid is heated, the P4 molecules separate and the vapour is formed. From this we
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 133
conclude that the covalent bonds which bind atoms into molecules are generally stronger than the van der Waals forces which
bind mo lecules into a liquid or solid.
Black
White
Red – (unsure, amorphous)
Black phosphorus is a crystalline network solid. White phosphorus is a nonpolar substance and, like rhombic sulfur, it will
dissolve in nonpolar liquids such as carbon dis ulfide and carbon tetrachloride, but not in water.
Red phosphorus is amorphous rather than crystalline. It is therefore much harder to obtain the detailed structure of this
solid, since crystallographic methods cannot be used. The structure of red phosphorus is thought to be chains of ring-opened
P4 cages, as shown above. It is prepared from white phosphorus by heating in the absence of oxygen at atmospheric pressure.
This thermal activation allows the ring-opening polymerization reaction to occur. As one might expect, red phosphorus is the
more stable allotrope; it does not oxidize in air at ordinary temperatures nor does it dissolve in nonpolar liquids such as carbon
disulfide. If red phosphorus is heated to a sufficiently high temperature, the bonds between the atoms are broken and, on
cooling, P4 molecules are formed which condense to the liquid and solid forms of white phosphorus.
Frost diagram in 1 M acid comparing all the Group 15 Elements
P
O
O
O
P
P
O
P
O
O
P4O6
O
P
O
O
O
O
P
P
O
P
O
O
O
P 4O10
O
We note that unlike N(V), P(V) is highly stable, in fact the most stable form of phosphorus. Thus H3PO4 is one of the best
examples of a non-oxidizing strong acid, even though it is not particularly strong on the Brønsted scale. All the truly strong
acids turn out to have oxidizing properties. Note that phosphorous acid and its salts, including the anhydride P4O6 are potent
reducing agents, as is elemental phosphorus and phosphane, PH3. Note that arsenic and antimony are most stable in the
elemental form, and that Bi(V) is also a potent oxidizing agent (e.g. compare Pb(IV) in the same period!)