Category Archives: Inorganic Chemistry

Discovery of Elements: Indium


In the history of chemical elements the discovery of a new element often directly affected the discovery of another one. Thus, the discovery of thallium was a catalyst for the discovery of indium–the last of the classic group of four elements identified by spectral analysis.

The stage was set in the German town of Freiberg; and the main characters were F. Reich, professor of physics in the Mining Academy and his assistant Th. Richter. The time was the year of 1863. Interested in some properties if thallium, discovered two year earlier, F Reich decided to obtain a sufficient amount of the metal for his experiments. Searching for natural sources of thallium, he analysed samples of zinc ores mined at Himmelsfürst. In addition to zinc the ores were known to contain Sulphur, arsenic, lead, silicon, manganese, tin, and cadmium, in a word, quite a number of chemical elements. Reich believed that thallium could be added to the list. Although time-consuming chemical experiments did not produce the desired element, he obtained a straw-yellow precipitate of an unknown composition. It was told that when C. Winkler (subsequently the discoverer of germanium) entered Reich’s laboratory the latter showed him a test-tube with the precipitate and said that it contained sulphide of a new element.

It would have been surprising if F. Reich had not used spectroscopy to prove his assumption. Of course, Reich did use it but, unfortunately, he was colour-blind and, therefore, asked his assistant Richter to perform spectral analysis. Th. Richter succeeded in the very first attempt: in the spectrum of the sample he saw an extremely bright blue line which could not be confused with either cesium blue line or any other line. In a word, the observation was quite definite. Reich and Richter came to the conclusion that the ores of Himmelsfurst contained a new chemical element. They named it “indium” after “indigo”, a bright blue dye. There is an interesting fact that does credit to F. Reich. The first reports about the discovery of indium were signed by the two scientists. Reich, however, believed that this was unjust and that the honour of the discovery belonged solely to Richter.

Soon after the two scientists had proved the existence of natural indium with the help of spectroscopy, they obtained a small amount of it. Indium compounds turn the flame of a Bunsen burner blue-violet and so bright that presence of the new element could be established without a spectroscope. Subsequently Reich and Richter studied some properties of indium, with Winkler giving them considerable help.

When metallic indium, although contaminated, was prepared, Richter submitted the samples to the Paris Academy of Sciences in 1867 and estimated their value at 800 pounds sterling which was quite a lot of money at the time.

Chemical properties of indium were described soon after its discovery but its atomic mass was at first determined incorrectly (75.6). Mendeleev saw that this atomic mass would not correctly place indium in the periodic table and suggested to increase it by about 50 per cent. Mendeleev proved to be right and indium occupied its place in the third group of the periodic table.

Sodium and Potassium

Sodium and Potassium

Man had known sodium and potassium compounds for a very long time. Carbonates of these metals were used in Egypt for laundry. Common salt, one of the most widespread sodium compounds, was used in foods from time immemorial; in some countries it was very expensive and sometimes wars were waged for the right to possess salt mines. Sodium carbonate was usually obtained from salt lakes whereas potassium carbonate by leaching plant ash; for this reason, the former was named mineral alkali and the latter vegetable alkali. The word “alkali” was introduced by Geber, a medieval alchemist, although he made no distinction between the two carbonates. The differences in their nature were first mentioned in 1683. The Dutch scientist I. Bon noted that when soda and potash were used in the similar process. The shapes of the precipitated crystals were different depending on the initial product.

In 1702 G. Stahl noted the difference in crystals of some sodium and potassium compounds. This was an important step in distinguishing between soda and potash. In 1736 the French chemist A. de Monsean proved that soda was always present in common salt, Glauber’s salt, and in borax. Since an acidic constituent of soda was known, the nature of the basic constuent was of great interest. According to Monsean, soda formed Glauber’s salt with sulphuric acid cubic saltpeter (sodium nitrate) with nitric acid, and a variety of sea salt with hydrochloric acid: isn’t this reason enough to deduce that soda is the basis of sea salt?

Although chemists had suspected for a long time that alkali earths were oxides of metals, the nature of soda and potash had not been studied up to the early 19th century. Even Lavoisier had no definite idea on this subject. He did not know what the basic constituents of soda and potash were and assumed that nitrogen could be a constituent. This confusion seems to stem from the similarity between the properties of sodium, potassium, and ammonium salts. Credit for determining these constituents belongs to H. Davy. At first he was dogged by failures: he could not separate metals from soda and potash with the aid of a galvanic battery. However, soon the scientist understood his error–he used saturated aqueous solutions but the presence of water hinders decomposition. In October, 1807, Davy decide to melt anhydrous potash, and as soon as he started electrolysis of the alkali hydroxide melt, small balls resembling mercury with bright metallic lustre appeared on the negative electrode immersed into the melt. Some of the balls burnt up immediately with an explosion forming bright flame while the other did not burn, but just dimmed and became covered with a white film, Davy concluded that numerous experiments had shown that the balls were the substance which he had been looking for and this substance was highly inflammable potassium hydroxide.

Davy studied this metal thoroughly and found that when it reacted with water the resulting flame was due to burning of the hydrogen liberated from water. Having studied the metal obtained from potassium hydroxide, H. Davy began to search for sodium hydroxide using the same method and the succeeded in separating another alkali metal. The scientist noted that for its preparation much more powerful battery was required than in the experiments with potash. Never the less, the properties of both metals turned out to be similar. For a short time the scientist carefully studied the properties of potassium and sodium. Some chemists doubted the elemental nature of sodium and potassium believing that they were compounds of alkalis with hydrogen. However, Gay Lussac and Thenard proved convincingly that Davy had, indeed. Obtained simple substances.

Discovery of Elements : Lithium


The fate of the lightest metal is outwardly uneventful. It was the third alkali metal to be discovered in nature. Its abundance on Earth is much less than that of sodium and potassium, its minerals are rare and, therefore, it came relatively later to man’s attention.

At the very beginning of the 18th century the prominent Brazilian scientist and statesman J. Andrada e Silva was travelling in Scandinavia. A passionate mineralogist, he wanted to enrich his collection with new specimens. He had luck and found two new minerals which he named petalite and spodumene. J. Andrada e Silva found the minerals at the island of Uto belonging to Sweden. Soon spoudmene was found in other places but the existence of petalite was doubted until, in 1817, it was found at Uto the second time.

Therefore, spodumene was the first to become the subject of investigation. M. Klaproth studied it but discovered nothing except alumina and silica. In a word, spodumene was a typical aluminosilicate. But the total mass of the isolated components was 9.5 per cent less than the mass of the initial sample, and Klaproth could not explain the reason for this considerable loss. Meanwhile, his compatriot I. Nepomuk von Fux discovered by chance that a pinch of spodumene turned the burner flame red. The scientist did nottry to find to the reason for this phenomenon, and that was a mistake, since he could have discovered a new element in spodumene.

The second discovery of petalite attracted attention to the mineral. L. Vauquelin found alkali in it, in addition to alumina and silica, but erroneously identified it with potash. W. Hizinger obtained interesting and suggestive results but had no chance to explain them since the same data had already been published by the Swedish chemist I. Arfvedson to whom the credit for discovering lithium went. J. Berzelius in his latter to A. Berthollet, the famous French chemist, on February 9, 1818, described this event in the following way. A new alkali, he wrote, was discovered by I. Arfvedson, a skillful young chemist, who had been working in his laboratory for a year. Arfvedson found the alkali in the ore discovered earlier by Andrada at the Uto mine and named petalite. The ore consisted of 80 per cent silicon oxide, 17 per cent aluminium and 3 per cent the new alkali. The conventional method used to extract the alkali consisted in heating the ground ore with barium carbonate and separating all earths from it.

Analysing petalite, Arfvedson from the very beginning discovered that the losses of the material amounted to about 4 per cent. The Swedish chemist (like M. Klaproth in his time) tried to find the answer again and again, sweeping aside various assumptions, and at last reached the truth–it was a new alkali of unknown nature. It was clear that this alkali was formed by a new alkali metal. I. Arfvedson asked his teacher to help him choose the name for the metal and the scientists decided to name it “lithium” (from the Greek lithios for “stone”). This name is a reminder that lithium was discovered in the mineral kingdom whereas two other alkali metals (sodium and potassium) in the plant kingdom Arfvedson published the report on the discovery of lithium in petalite in 1819 but already in April, 1818, the scientist found the new alkali metal in other minerals as well. The secret of spodumene, which Klaproth had failed to solve, was finally cleared: the mineral contained about 8 per cent of lithium. And one more mineral, lepidolite, known for a long time, was also found to contain up to 4 per cent of the lightest alkali metal.

The German chemist K. Gmelin observed lithium salts to turn the burner flame a beautiful shade of red (to I. von Fux’s great irritation).

By the late 1818 H. Davy succeeded in separating pure lithium, though in very small amounts. It became possible to obtain large amounts of lithium only in the late 1850’s when the German chemists Bunsen and Matissen developed an industrial process of electrolysis of lithium chloride.



R.G. Pearson (1963) has classified the Lewis acids and Lewis bases as hard and soft acids and bases.

 A third category whose characteristics are intermediate between those of hard and soft acids/bases are called borderline acids or borderline bases.


Pearson’s hard and soft acids and bases principle (HSAB Principle)

On the basis of experimental data of various complexes obtained by the combination of Lewis acids and Lewis bases, Pearson (1963) discovered a principle known as Hard and Soft Acids and Bases Principle (HSAB Principle) (some chemists prefer the abbreviation SHAB instead of HSAB used by Pearson).

This principle states that a hard Lewis acid prefers to combine with a hard Lewis base and similarly a soft Lewis acid prefers to combine with a soft Lewis base, since this type of combination gives a more stable product.

Thus we can say that (hard acid + hard base) and (soft acid + soft base) combinations give more stable products than the (hard acid + soft base) and (soft acid + hard base) combinations.

 The combination of hard acid and hard base occurs mainly through ionic bonding as in Mg(OH)2 (Mg2+ = hard acid, OH = hard base) and that of soft acid and soft base takes place mainly by covalent bonding as in HgI2 (Hg2+ = soft acid, I = soft base.)


Table: Classification of Lewis acids and Lewis bases into hard, soft and borderline acids and bases

Lewis Acids (acceptors)

Hard acids [Ahrland and Chatt (1958) have arbitrarily called hard acids as class (a) metal ions or metal acceptors] Soft acids [These have been called class (b) metal ions or metal acceptors] Borderline (intermediate) acids
(i) They have acceptor metal atom of small size.


(ii) they have acceptor with high positive charge (oxidation state)


(iii) The valence-electrons of the acceptor atom of these acids cannot be polarised (or distorted or removed) easily (i.e., they have low polarisability), since they are held strongly and it is for this reason that these Lewis acids have been called hard acids (or hard metal ions) by Pearson (1963).

(i) They have acceptor metal atom of large size.


(ii) They have acceptor atom with low or zero positive charge


(iii) The valence-electrons of the acceptor atom of these acids can be polarised easily (i.e., they have high polarisability), since they are held weakly and for this reason these Lewis acids have been called soft acids (or soft metal ions) by Pearson.

The characteristics of borderline acids are intermediate between those of hard acids and soft acids.
Examples: H+, Li+, Na+, K+, Be2+, Mg2+, Ca2+, Sr2+, Mn2+, Ag2+, Al3+, Sc3+, Ga3+, In3+, La3+, N3+, Cl3+, Gd3+, Lu3+, Cr3+, Co3+, Fe3+, As3+, CH3Sn3+, Si4+, Ti4+, Zr4+, Th4+, Li4+, Pu4+, Ce3+, Hf4+, Wo4+, Sn4+, UO22+, MoO3+, BeMe2, BF2, B(OR)3, Al(CH3)3, AlCl3, AlH3, RPo2, SO3. Phenol, I7+, I5+, Cl7+, Cr6+, RCO+, Fe6+, Pt6+, CO2, NC+, HX (hydrogen-bonding molecules) Examples: Cu+, Ag+, Au+, Ti+, Hg+, Cs+, Pd2+, Cd2+, Pt2+, Hg2+,CH3Hg+, Co(CN)52−, Pt4+, Te4+, Ti3+, TI(CH3)3, BH3, Ga(CH3)3, GaCl3, GaI3, InCl3, RS+, RSe+, RTe+, I+, Br+, I2, Br2, ICN, trinitrobenzene, chloranil, quinones, tetracyanoethylene O, Cl, Br, I, N, RO, RO2 M (metal atoms), CH2 (carbene) Examples: Fe2+, Co2+, Ni2+, Cu2+, Zn2+, Pb2+, Sn2+, Sb3+, Bi3+, Rh3+, Ir3+, B(CH3)3, SO­2­, NO+, Ru2+, Os2+, R3C+, C6H5+, GaH3.



Lewis Bases (Donors or ligands)

Hard bases (Hard ligands) Soft bases (Soft ligands) Borderline (intermediate bases
The donor atom of a hard base: The donor atom of soft base: These bases have intermediate properties.
(i) has high electronegativity.


(ii) holds its valence-electrons strongly and hence cannot be polarised (re-moved or deformed) easily, i.e., the donor atom of hard base has low polarisability.

(iii) has filled orbitals



Examples: H2O, OH, ROH, R2O, RO, CH3COO, PO43−, SO42− RCO2, CO32−, ClO4, NO3, O2−, C2O42−, (co-ordination through O-atom), NH3, NR3, NHR2,. NH2R, N2H4, NCS (co-ordination through N-atom) F, Cl.

(i) has low electronegativity.

(ii) holds its valence-electrons weakly and hence can be polarised easily, i.e., the donor atom of a soft base has high polarisability.

(iii) has partially filled orbitals





Examples: R2S, RSH, RS, SCN (co-ordination through S-atom). S2−, R3P, R3As, I, CN, H, R, S2O32−, (RO)3P, RNC, CO, C2H4, C6H6, CH3.

These base have intermediate properties.











Examples: C6H5NH2, C5H5N, N3, Br, NO2, SO32−, N2.


Applications of HSAB principle

HSAB principle is extremely useful in explaining the following:

1. Stability of complex compounds, having the same ligands.

This application can be understood by considering the following examples :

(a)      \[AgI_{2}^{-}  is stable while \[AgI_{2}^{-}  does not exist. We know that Ag+ is a soft acid, F ion is a hard base and I ion is a soft base. Thus, since \[AgI_{2}^{-}  is obtained by the combination of a soft acid (Ag+) and soft base (I) and  \[AgI_{2}^{-} results by the interaction of a soft acid (Ag+) and a hard base (F), \[AgI_{2}^{-} ion is stable but \[AgI_{2}^{-} does not exist.


(b)       CoF63− (hard acid + hard base) is more stable than CoI63− (hard acid + soft base).

(c)        (CH3)2 \[\overset{\centerdot \,\,\centerdot }{\mathop{N}}\,-\overset{\centerdot \,\,\,\centerdot }{\mathop{P}}\,{{F}_{2}} molecule acts as a bidentate ligand, since it has two lone pairs of electrons one of which is on N-atom and the other is on P-atom. Both BH3 and BF3 molecules combine with this ligand and forms an adduct. With the help of HSAB principle we can predict the structure of this adduct. We know that BF3 is a hard acid and  \[\overset{\centerdot \,\,\centerdot }{\mathop{N}}\, (CH3)2 is a hard base. It also known that BH3 is a soft acid and −  \[\overset{\centerdot \,\,\,\centerdot }{\mathop{P}}\,{{F}_{2}}  is a soft base. On applying the principle that (hard acid + hard base) combinations and (soft acid + soft base) are preferred, the structure of the adduct should be that in which N-atom donates its lone pairs of electrons to B-atom of BF3 molecules and P-atom donates its lone pair of electrons to B-atom of BH3 molecule. Thus the structure of the adduct is:

2. To predict the nature of bonding in complex ions given by ambidentate ligands

(a)       With the help of HSAB principle we can predict which atom of an ambidentate ligand will combine with metal ion too form the complex, SCN ion is an ambidentate ligand since it can co-ordinate to the metal ion either through its S-atom or through N-atom. It has been found that Co2+ and Pd2+ both combine with four SCN ligands to form the complex ion, [M(SCN)4]2− (M = Co2+, Pd2+). With the help of HSAB principle it can be shown that in [Co(SCN)4]2− ion, Co2+ is linked with the ligand through N-atom while in [Pd(SCN)4]2− ion, Pd2+ is co-ordinated with the ligand through S-atom. Thus the complex ions given by Co2+ and Pd2+ ion should be represented as [Co(NCS)4]2− and [Pd(SCN)4]2− respectively. The reason for this is that since Co2+ ion is a hard acid, it prefers to co-ordinate with N-atom of the hard ligand, NCS. On the other, Pd2+ ion is soft acid and hence combines with the S-atom of the soft ligand, SCN.


(b)       We know that phenol is a hard acid and I2 is soft acid. It is also known that alkyl thiocyanate, RSCN (S-atom acting as a donor) is a soft ligand and alkyl iso-thiocyanate, RNCS (N-atom acting as a donor) is a hard ligand. Thus, if RSCN and RNCS are complexed with phenol and I2, RNCS will form more stable complex with phenol due to hard acid (phenol) – hard ligand (RNCS) combination than that with I2. On the other hand RSCN will give more stable complex I2 due to soft acid (I2) – soft ligand (RSCN) combination.


3. Stability of complex compounds having different ligands. Jorgensen has pointed out that in a complex compound having different ligands, if all the ligands are of the same nature, i.e., if all the ligands are soft ligands or hard ligands, the complex compound will be stable. On the other hand, of the ligands are of different nature, the complex compound would be unstable. This point may be illustrated by the following examples :

(a)       Since

  • in [Co(NH35F]2+ both the ligands , NH3 molecule and F ion are hard ligands
  • in [Co(NH35I]2+ (II) NH3 is a hard ligand and I ion is a soft ligand, therefore (I) is a stable complex ion while (II) is unstable.


  • [Co(CN)5I]3− (I) is more stable than [Co(CN)5F]3− (II) because
  • in (I) both the ligands are soft ligands
  • while in (II) CN ions are soft ligands and F ion is a hard ligand.


4. Symbiosis. Soft ligands prefers to get attached with a centre which is already linked with soft ligands. Similarly hard ligands prefers to get attached with a centre which is already linked with hard ligands. This tendency of ligands is called symbiosis and can be explained by considering the formation of (F3B ← NH3) adduct and BH4 ion. Hard ligand like NH3 co-ordinates with B-atom of BF3 molecule to form (F3B ← NH3) adduct, since F ions which are already attached with B-atom in BF3 molecule are also hard ligands. Thus :


Similarly the formation of BH4 ion by the combination of BH3 (in which H atoms are soft ligands) and H ion (soft ligands) can also be explained

The formation of F3B ← NH3 adduct can also be explained on the basis of the fact that since BF3­ and NH3 are hard acid and hard base respectively, they combine together to form a stable F3B ← NH3 adduct.

F3B (hard acid) + NH3 (Hard base) → F3B ← NH3 (Stable adduct)

Similarly since BH3 is soft acid and H ion is a soft base, their combination gives a stable BH4 ion.

BH3 (Soft acid) + H (Soft base) → BH4 (Stable ion)


5. Solubility of compounds:

This point would be more clear when we compare the relative stability of HgS and Hg(OH)2 in acidic aqueous solution. HgS (soft acid + soft base) in more stable than Hg(OH)2 (soft acid + hard base).

More stability of HgS than that of Hg(OH)2 explains why Hg(OH)2 dissolves readily in acidic aqueous solution but HgS does not.


6. Occurrence of metals in nature.

The occurrence of some metals in nature as their ores can be explained with the help of HSAB principle.

This following examples illustrate this point :

(a)       We know that since MgCO3, CaCO3 and Al2O3 are obtained by the combination of hard acids viz., Mg2+, Ca2+ and Al3+ ion with hard bases namely CO32− and O2− ions while MgS, CaS and Al2S3 are obtained by the combination of hard acids (Mg2+, Ca2+, Al3+ ions) and soft base viz., S2− ion, Mg, Ca and Al occur in nature as MgCO3, CaCO3 and Al2O3 respectively and not as their sulphides (MgS, CaS and Al2S3).

(b)       Since Cu2S, Ag2S and HgS are obtained by the combination of soft acids namely Cu+, Ag+ and Hg2+ ion and soft base viz., S2−­ ion while Cu2CO3, Ag2CO3 and HgCO3 result by the interaction of soft acids (Cu+, Ag+, Hg2+) and hard base viz., (CO32−, Cu, Ag and Hg occur in nature as their sulphides (Cu2S, Ag2S and HgS) and not as their carbonates.

(c)        Ni2+, Cu2+ and Pb2+ ions which are borderline (intermediate) acids occurs in nature both as carbonates and sulphides.


7. Jorginsen has also pointed that hard solvents tend to dissolve hard solutes and vice versa.


8. Course of reaction. The principle of (hard acid + hard base) and (soft acid + soft base) combination has also been used to predict the course of many reaction. For example :

\[\underset{(hard\,\,acid\,+\,soft\,\,base)}{\mathop{LiI}}\,+\underset{\left( soft\,\,acid\,+\,hard\,base \right)}{\mathop{CsF}}\,\xrightarrow{{}}\underset{(hard\,\,acid\,+\,hard\,\,base)}{\mathop{LiF}}\,+\underset{(soft\,\,acid\,+\,soft\,base)}{\mathop{CsI}}\,

\[\underset{soft\,\,acid\,+\,hard\,\,base)}{\mathop{Hg{{F}_{2}}}}\,+\underset{\left( hard\,\,acid\,+\,soft\,base \right)}{\mathop{Be{{J}_{2}}}}\,\xrightarrow{{}}\underset{(hard\,\,acid\,+\,hard\,\,base)}{\mathop{Be{{F}_{2}}}}\,+\underset{(soft\,\,acid\,+\,soft\,base)}{\mathop{Hg{{I}_{2}}}}\,


Limitations of HSAB principle

Although (hard + hard) and (soft + soft) combination is a useful principle, yet many reaction cannot be explained with the help of this principle. For example in the reaction:


Or    \[\underset{soft\,\,base}{\mathop{SO_{3}^{2-}}}\,+\underset{(hard\,\,acid\,+\,hard\,\,base)}{\mathop{{{H}^{+}}\,\,{{F}^{-}}}}\,\xrightarrow{{}}\,\underset{(hard\,\,acid\,+\,soft\,\,base)}{\mathop{{{[H]}^{+}}\,\,{{[S{{O}_{3}}]}^{2-}}}}\,+{{F}^{-}}

Which proceeds towards right, hard acid (H+) combines with soft or borderline base  to form [H+]  or ion which is a stable ion. (Hard acid + soft base) combination is against the HSAB principle.


Discovery of element Nickel


Nickel has very much in common with Cobalt, its neighbour in the periodic table. First of all, Nickel is also of “devilish” origin. Its name derives from the German “kupfer–nickel” (“copper devil”) and belongs to the mineral described in 1694 by the Swedish mineralogist U. Hierne, who mistook it for copper ore.

When repeated attempts to smelt copper from it failed, the metallurgists decided that it must have been Nick, the evil spirit of the mountains, at his tricks.

People came to know Nickel ages ago. Back in the 3rd century B.C the Chinese made an alloy of copper, nickel, and zinc. In the Central Asian state of Bactria coins were made from this alloy. One of them is now in the British Museum in London.

Confusion about the composition of kupfernickel remained even after the mineral had been described. In 1726 the German chemist I. Link studied the mineral  and established that its dissolution in nitric acid yields a green colour. He concluded that the mineral was most probably a cobalt ore with admixtures of copper. When Swedish miners found a reddish mineral which, being added to glass, did not produce a blue colour, they named it “cobold that had lost his soul”. It was also one of the nickel minerals.

That was how matters stood up to 1751. That year the Swedish mineralogist and chemist A. Cronstedt took an interest in the mineral found in a cobalt mine. In one of his experiments he immersed a small piece of iron into an acid solution of this ore. Had copper been present in the solution, it would have been deposited on iron in a free state. To his great surprise nothing of the kind happened. The solution did not contain copper. This contradicted the existing beliefs about this ore. Cronstedt began a thorough investigation of the green crystals dispersed in the ore. After a great number of experiments, he isolated a metal from kupfernickel which did not resemble copper at all. Cronstedt described this metal as solid and brittle, weakly attracted by a magnet, transforming into a black powder when heated, and yielding a wonderful green colour upon dissolution. Cronstedt concluded that, since the metal was contained in kupfernickel, the name could be retained and shortened to nickel. At present it is known that kupfernickel is nickel arsenide.

Many chemists in Europe recognized that a new element had been discovered. But some scientists held that nickel was mixture of cobalt, iron, arsenic and copper. All doubts were removed in 1775 by T. Bergman who showed that mixtures of these elements taken in any proportions did not possess the properties of Nickel.

Nitrogen : It’s discovery



Discovery of Nitrogen : 

The study of the atmosphere led to the discovery of nitrogen. Although it is associated with the name of a certain scientist and a certain date, this certainly is misleading. It is rather difficult to separate the history of nitrogen discovery from the mainstream of pneumatic chemistry; one can only think of a more or less logical sequence of events.

Very early in history man came across nitrogen compounds, for instance, saltpeter and nitric acid, frequently observing liberation of brown vapours of nitrogen dioxide. Obviously, it would be impossible to discover nitrogen by decomposing its inorganic compounds. Tasteless, colourless, odorless, and chemically rather inactive, nitrogen would have remained unnoticed.

Therefore, it is not easy to decide where to begin the story of the discovery of nitrogen. Although our choice may seem subjective, we start with 1767 when H. Cavendish and J. Priestley, another outstanding English physicist, chemist, and philosopher, set out to study the action of electric discharges on various gases. There were only few such gases at that time: ordinary air, fixed air, and inflammable air. Although the experiments did not produce definite results, it was shown later that electric discharge in humid air yields nitric acid. Later this fact proved to be useful for the analysis of the earth’s atmosphere.

In 1777 H. Cavendish reported in a private letter to J. Priestley that he had succeeded in preparing a new variety of air named by him asphyxiating or mephitic air. Cavendish repeatedly passed atmospheric air over red–hot coal. The resulting fixed air was absorbed with alkali. The residue was mephitic gas. Cavendish did not study it thoroughly and only reported the fact to Priestley. Cavendish returned to the study of mephitic air much later, did a large work but the credit for the discovery had already gone to another scientist.

When Priestley received the letter from Cavendish he was busy with important experiments and read it without due attention. Priestley burned various inflammable compounds in a given volume of air and calcinated metals; the fixed air formed during these processes was removed with the aid of limewater. The main thing which he noticed was that the volume of air decreased considerably. A reader will prompt that as a result of metal calcination or combustion of compounds the oxygen present in the apparatus was bonded and nitrogen remained. Priestley, however, had no idea about the existence of such a gas as oxygen (two years later, however, he became one of its discoverers) and, to explain the observed phenomenon, he turned to phlogiston. Priestley believed that the result of metal calcination was due exclusively to the action of phlogiston. The remaining air is saturated with phlogiston and, consequently, it can be named “phlogisticated” air; it sustains neither respiration nor combustion.

Thus, Priestley was in possession of a gas which subsequently became known as nitrogen. But this extremely important result was treated by him as something of secondary importance. The existence of “phlogisticated” air was for Priestley evidence of the fact that phlogiston did play a role in natural processes. This story shows once more how the erroneous phlogistic theory hampered the discovery of elemental gases.

So, neither Cavendish nor Priestley could understand the real nature of the new gas. The understanding came later when oxygen appeared on the scene of chemistry. The English physician D. Rutherford, the pupil of J. Black, who is considered to be the discoverer of nitrogen, did, in principle, nothing new compared with his famous colleagues. In September 1772, Rutherford published a magisterial thesis On the So–called fixed and Mephitic Air in which he described the properties of nitrogen. This gas, according to Rutherford, was absorbed neither by limewater nor by alkali and was unsuitable for respiration; he named it “corrupted” air.

Not properly discovered or understood as a gaseous chemical element, nitrogen in the seventies of the 18th century had three names which confused still more the fuzzy chemical concept muddled by the persisting influence of the phlogistic theory. Phlogisticated, mephitic, or corrupted air was yet receive its final name.

This name was proposed in 1787 by A. Lavoisier and other French scientists who developed the principles of as new chemical nomenclature. They derived the word “azote” from the Greek negative prefix “a” and the word “zoe” meaning “life”. Lifeless, not supporting respiration and combustion, that was how the chemists saw the main property of nitrogen. Later this view turned out to be erroneous: nitrogen is vitally important for plants. The name “azote”, however, remained. The symbol of the element, N, originates from the Latin nitrogenium which means “saltpeter–forming”.Cavendish studied the properties of nitrogen in detail. He was one of the first scientists to believe that phlogisticated air is a component of ordinary air. One day, in the course of his experiments Cavendish questioned the homogeneity of phlogisticated air. He passed an electric spark through its mixture with oxygen transforming the whole into nitrogen oxides which were removed from the reaction zone. But every time a small fraction of the phlogisticated air (nitrogen) remained unchanged and did not react with oxygen. It was a very small fraction, a slightly noticeable gas bubble–only 1/125 of all nitrogen taken for the phenomenon observed in 1785. The answer was found only over one hundred years later. You will read about it in chapter 9 devoted to inert gases.



Discovery of Hydrogen

Hydrogen is one of the most striking elements of the periodic system, its number one, and the lightest of all the existing gases. It is the element whose discovery was indispensable for the solution of many problems of chemical theory. It is an element whose atom, losing its only valence electron, becomes a “bare” proton. And, therefore, chemistry of hydrogen is, in a way, unique; it is the chemistry of an elementary particle.

Once D.I. Mendeleev called hydrogen the typical of typical elements (meaning the elements of the short periods in the system), because it begins the natural series of chemical elements.

And such a fascinating element is readily available. It can be obtained without difficulty in any school laboratory, for instance, by pouring hydrochloric acid on zinc shavings.

Even in those bygone times, when chemistry was not a science yet and when alchemists were still searching for the “philosophers’ stone”, hydrochloric, sulphuric, and nitric acids as well as iron and zinc were already known. In other words, man had in his possession all  components whose reaction could give rise to hydrogen. Only a chance was needed and chemical literature of the 16–18th centuries reported that many times chemists observed how the pouring of, for instance, sulphuric acid on iron shavings produced bubbles of a gas which they believed to be an inflammable variety of air.

  • One of those who observed this mysterious variety of air was the famous Russian Scientist M. V Lomonosov. In 1745 he wrote a thesis, On Metallic Lustre, which said, among other things: “On dissolution of some base metal, especially iron, in acidic alcohols, inflammable vapour shots out from the opening of the flask….” (According to the terminology of those times, acidic alcohols meant acids.) Thus, M.V Lomonosov observed none other than hydrogen. But the sentence went on to read: “……which is phlogiston.” since metal dissolved in the acid liberating material ignea or “inflammable vapour”, it was very convenient to assume that dissolving metal releases phlogiston: everything fits nicely into the theory of phlogiston.


  • And now is the time to meet the outstanding English scientist H. Cavendish, a man fanatically devoted to science and an excellent experimenter. He never hurried with making public his experimental results and sometimes several years had to pass before his articles appeared. Therefore, it is difficult to pinpoint the date when the scientist observed and described the liberation of “inflammable air”. What is known is that this work published in 1766 and entitled “Experiments with Artificial Air” was done as a part of pneumatic chemistry research. It is also likely that the work was performed under the influence of J. Black. H. Cavendish had become interested in fixed air and decided to see whether there existed other types of artificial air. In this manner the scientist referred to the variety of air which is contained in compounds in a bound state and which can be separated from them artificially. H. Cavendish knew that inflammable air had been observed many times. He himself obtained it by the same technique: the action of sulphuric and hydrochloric acids on Iron, Zinc, and Tin, but he was the first to obtain definite proof that the same type of air was farmed in all cases–inflammable air. And he was the first to notice the unusual properties of inflammable air. As a follower of the phlogistic theory, H. Cavendish could give only one interpretation of the substance’s nature. Like M. V. Lomonosov, he identified it as phlogiston. Studying the properties of inflammable air, he was sure that he was studying the properties of phlogiston. H. Cavendish believed that different metals contain different proportions of inflammable air. Thus, to the fixed air of J. Black, the inflammable air of H. Cavendish was added. Strictly speaking, the two scientists discovered nothing new: each of them only summarized the data of previous observations. But this summing up represented considerable progress in the history of human knowledge.

Fixed air and inflammable air differed both from ordinary air and from each other. Inflammable air was surprisingly light. H. Cavendish found that phlogiston, which he had separated, had a mass. He was the first to introduce a quantity to characterize gases, that of 

Hydrogen is a gas under standard conditions. Hydrogen is used for rocket fuel and explosives.

density. Having assumed the density of air to be unity, Cavendish obtained the density of 0.09 for inflammable air and 1.57 for fixed air. But here a contradiction arose between Cavendish the experimenter and Cavendish the adherent of the phlogistic theory. Since inflammable air had a positive mass, it could by no means be considered to be pure phlogiston. Otherwise, metals losing inflammable air would have to lose mass as well. To avoid the contradiction, Cavendish proposed an original hypothesis: inflammable air is a combination of phlogiston and water. The essence of the hypothesis was that at last hydrogen appeared in the composition of inflammable air.

The evident conclusion is that Cavendish, like his predecessors, did not understand the nature of inflammable air, although he had weighed it, described its properties, and considered it to be an independent kind of artificial air. In a word, Cavendish, unaware of the fact, studied “phlogiston” obtained by him as he would have studied a new chemical element. But Cavendish could not perceive that inflammable air was a gaseous chemical element–so strong were the chains of the phlogistic theory. And having found that the real properties of inflammable air contradicted this theory, he came up with a new hypothesis, as erroneous as the theory itself.

  • Therefore, strictly speaking, the phrase “hydrogen was discovered in 1766 by the English scientist H. Cavendish” is meaningless. Cavendish described the processes of preparation and the properties of inflammable air in greater detail than his predecessors. However, he “knew not what he was doing”. The elementary nature of inflammable air remained beyond his grasp. It was not the scientist’s fault, however; chemistry had not yet matured enough for such an insight. Many years have passed before hydrogen became, at last, Hydrogen and occupied its proper place in chemistry.
  • Its Latin name hydrogenium stems from the Greek words hydr and gennao which mean “producing water”. The name was proposed in 1779 by A. Lavoisier after the composition of water had been established. The symbol H was proposed by J. Berzelius.

Hydrogen is a unique element in the sense that its isotopes differ in their physical and chemical properties. At one time this difference prompted some scientists to consider hydrogen isotopes as independent elements and to find for them special boxes in the periodic table. Therefore, the history of the discovery of hydrogen isotopes is of special interest, as a continuation of the history of hydrogen itself.

The discovery of Isotopes of Hydrogen:

The search for hydrogen isotopes began in the twenties of this century but all attempts were unsuccessful, resulting in the belief that hydrogen had no isotopes.

In 1931 it was suggested that hydrogen, nevertheless, contains a heavy isotope with a mass number of 2. Since this isotope had to be twice as heavy as hydrogen, the scientists tried to isolate heavy hydrogen by physical methods.

  • In 1932 the American scientists Urey, Brickwedde, and Murphy evaporated liquid hydrogen and, studying the residue by spectroscopy, found a heavy isotope in it. In the atmosphere it was discovered only in 1941. The name “deuterium” originates from the Greek word deuteros which means “second, another one”.
  • The next isotope with a mass number of 3, tritium (from the Greek–the third is radioactive and was discovered in 1934 by English scientists M. Oliphant, P. Hartec, and E. Rutherford.
    The name “protium” was assigned to the main hydrogen isotope. This is the only case when isotopes of the same element have different names and symbols (H, D and T). 99.99 per cent of all hydrogen is protium; the rest is deuterium with only traces of tritium.
Bismuth : It’s discovery



Bismuth has been known to mankind for centuries but for a long time it was confused with antimony, lead and tin. Paracelsus, for instance, said that there were two varieties of antimony–a black one used for the purification of gold and very similar to lead, and a white one named bismuth and very similar to lead, and a white one named bismuth and resembling tin; a mixture of these two varieties resembles silver. Form the chemical standpoint this confusion can easily be explained. Antimony and bismuth are analogues of each other and have common features with lead and tin, the elements of the previous group.

Agricola, unlike Paracelsus, gave a rather detailed description of bismuth and of the process of its extraction from ores mined in Saxony. Miners thought that bismuth, as well as tin, was a variety of lead and that bismuth could be transformed into silver.

In Central Russia bismuth has been known since the 15th century. With the development of book-printing bismuth, along with antimony, began to be used for casting typographical types. In literature few elements have such a great number of names as bismuth. E. von Lippmann in his book History of Bismuth from 1480 to 1800 gives twenty one names of this metal used in Europe. A sufficiently clear idea of bismuth as an independent metal was formed only in the 18th century.




Antimony and its compounds have been known from times immemorial. Some scholars say that metallic antimony was used in South Babylon for making vessels about 3400 years B.C. but in antiquity antimony was mainly used for making cosmetics such as rouge and black paint for eye brows. In Egypt, however, antimony was apparently unknown or almost unknown. This is borne out by finds from Egyptian burial sites, particularly, by painted mummies.

In antiquity antimony was confused with lead. It was only in alchemical literature of the Renaissance period that antimony was given a sufficiently accurate description. For example, G. Agricola clearly pointed out that antimony is a metal different from other metals. Basilius Valentinus devoted to antimony a whole treatise, Triumphal Carraige of Antimonium, in which he described the uses of antimony and its compounds.

There are several interpretations of the Latin name of antimony antimonium. Most likely it originates from the Greek word antimonos, which means “an enemy of solitude”, and underlines simultaneous occurrence of antimony and other minerals.

Why Halogens are Coloured ?


Why Halogens are Coloured ?

General Properties :

Some chemical and physical properties of the halogens are summarized in the Table below.
It can be seen that there is a regular increase in many of the properties of the halogens proceeding down group 17 from fluorine to iodine. This includes their melting points, boiling points, intensity of their colour, the radius of the corresponding halide ion, and the density of the element. On the other hand, there is a regular decrease in the first ionization energy as we go down this group. As a result, there is a regular increase in the ability to form high oxidation states and a decrease in the oxidizing strength of the halogens from fluorine to iodine.

Properties of Group 17 (The Halogens) F Cl Br I
Atomic number, Z 9 17 35 53
Ground state electronic configuration [He]2s2 2p5 [Ne]3s2 3p5 [Ar]3d10 4s2 4p5 [Kr]4d10 5s2 5p5
Colour pale yellow gas yellow-green gas red-brown liquid blue-black solid
Density of liquids at various temperatures, /kg m-3 1.51(85°K) 1.66(203°K) 3.19 (273°K) 3.96 (393°K)
Melting point, /K 53.53 171.6 265.8 386.85
Boiling point, /K 85.01 239.18 331.93 457.5
Enthalpy of atomisation, ΔaH° (298K) / kJ mol-1 79.08 121.8 111.7 106.7
Standard enthalpy of fusion of X2, ΔfusH°(mp) / kJ mol-1 0.51 6.4 10.57 15.52
Standard enthalpy of vaporization of X2, ΔvapH°(bp) / kJ mol-1 6.62 20.41 29.96 41.57
First ionization energy, IE1 / kJ mol-1 1681 1251.1 1139.9 1008.4
ΔEAH1°(298K) / kJ mol-1 -333 -348 -324 -295
ΔhydH°(X,g) / kJ mol-1 -504 -361 -330 -285
ΔhydS°(X,g) / JK-1 mol-1 -150 -90 -70 -50
ΔhydG°(X,g) / kJ mol-1 -459 -334 -309 -270
Standard redox potential, E°(X2 /2X) /V 2.87 1.36 1.09 0.54
Covalent radius, rcov = ½ X-X bond length /pm 72 100 114.2 133.3
Ionic radius, rion for X /pm 133 181 196 220
van der Waals radius, rv /pm 135 180 195 215
X-X(g)bond energy /kJ mol-1 159 243 193 151
H-X(g)bond energy /kJ mol-1 562 431 366 299
C-X(g)bond energy /kJ mol-1 484 338 276 238
Pauling electronegativity, χP 3.98 3.16 2.96 2.66



The origin of the colour of the halogens stems from the excitation between the highest occupied π* MO and the lowest unoccupied σ* MO. The energy gap between the HOMO and LUMO decreases according to F2 > Cl2 > Br2 > I2.

The amount of energy required for excitation depends upon the size of the atom. Fluorine is the smallest element in the group and the force of attraction between the nucleus and the outer electrons is very large. As a result, it requires a large excitation energy and absorbs violet light (high energy) and so appears pale yellow. On the other hand, iodine needs significantly less excitation energy and absorbs yellow light of low energy. Thus it appears dark violet. Using similar arguments, it is possible to explain the greenish yellow color of chlorine and the reddish brown color of bromine.

The halogens show a variety of colours when dissolved in different solvents. Solutions of iodine can be bright violet in CCl4, pink or reddish brown in aromatic hydrocarbons and deep brown in alcohols for example. This can be explained by weak donor-acceptor interaction and complex formation. The presence of charge-transfer bands further supports this since they are thought to be derived from interaction with the HOMO σu* orbital.

The X-ray structure of some of these have been obtained and often the intense colour can be used for characterisation and determination such as the bright blue colour of iodine in the presence of starch. In the case of the solid formed between dibromine and benzene, the structure is shown below and a new charge transfer band occurs at 292 nm. The Br-Br bond length is essentially unchanged from that of dibromine (228 pm).

In a study of the reaction of dibromine with substituted phosphines in diethyl ether, all but one showed a tetrahedral arrangement where one bromine was linked to the phosphorus.[3]

R3P + Br2 (Et2O, N2/r.t.) → R3PBr2
The X-ray study of the triethylphosphine was interpreted as [Et3PBr]Br where the Br-Br separation was 330 pm. This is considerably longer than the 228 pm found above and was taken to mean that the compound was ionic In the case of the tri (perfluorophenyl) phosphine however the structure showed both bromines linked to give a trigonal bipyramid arrangement with D3 symmetry. Why (C6F5)3PBr2 was the only R3PBr2 compound that adopted trigonal bipyramidal geometry was reasoned to be due to the very low basicity of the parent tertiary phosphine.


MP’s and BP’s

Intermolecular forces are the attractive forces between molecules without which all substances would be gases. The various types of these interactions span large differences in energy and for the halogens and interhalogens are generally quite small. The dispersion forces involved in these cases are called London forces (after Fritz Wolfgang London, 1900-1954). They are derived from momentary oscillations of electron charge in atoms and hence are present between all particles (atoms, ions and molecules).

The ease with which the electron cloud of an atom can be distorted to become asymmetric is termed the molecule’s polarizability. The greater the number of electrons an atom has, the farther the outer electrons will be from the nucleus, and the greater the chance for them to shift positions within the molecule. This means that larger nonpolar molecules tend to have stronger London dispersion forces. This is evident when considering the diatomic elements in group 17, the Halogens. All of these diatomic elements are nonpolar, covalently bonded molecules. Descending the group, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. For nonpolar molecules, the farther you go down the group, the stronger the London dispersion forces.

To picture how this occurs, compare the situation 1) where the electrons are evenly distributed and then consider 2) an instantaneous dipole that would arise from an uneven distribution of electrons on one side of the nucleus. When two molecules are close together, the instantaneous dipole of one molecule can induce a dipole in the second molecule. This results in synchronised motion of the electrons and an attraction between them. 3) Multiply this effect over numerous molecules and the overall result is that the attraction keeps these molecules together, and for diiodine is sufficient to make this a solid.

On average the electron cloud for molecules can be considered to be spherical in shape. When two non-polar molecules approach, attractions or repulsions between the electrons and nuclei can lead to distortions in their electron clouds (i.e. dipoles are induced). When more molecules interact these induced dipoles lead to intermolecular attraction.


The changes seen in the variation of MP and BP for the dihalogens and binary interhalogens van be attributed to the increase in the London dispersion forces of attraction between the molecules. In general they increase with increasing atomic number.


Redox properties
The most characteristic chemical feature of the halogens is their oxidising strength. Fluorine has the strongest oxidising ability, so that a simple chemical preparation is almost impossible and it must be prepared by electrolysis. Note that since fluorine reacts explosively with water oxidising it to dioxygen, finding reaction conditions for any reaction can be difficult. When fluorine is combined with other elements they generally exhibit high oxidation states. Chlorine is the next strongest oxidising agent, but it can be prepared by chemical oxidation. Most elements react directly with chlorine, bromine and iodine, with decreasing reactivity going down the Group, but often the reaction must be activated by heat or UV light. [2]

1. Enthalpy of Atomisation,              2. ΔEAH1,              3. ΔhydH°(X,g)

The redox potential, E°, X2/2X, measures a free-energy change, usually dominated by the ΔH term. The values in the Table show that there is a decrease in oxidising strength proceeding down the group (2.87, 1.36, 1.09, 0.54 V). This can be explained by comparing the steps shown above.

1) Atomisation of the dihalide is the energy required to break the molecule into atoms ½ X2(g) → X (g) note that only F2 and Cl2 are gases in their natural state so the energies associated with atomisation of Br2 and I2 requires converting the liquid or solid to gas first.

2) ΔEAH1 is the energy liberated when the atom is converted into a negative ion and is related to the Electron Affinity
X(g) + e → X(g) Addition of an electron to the small F atom is accompanied by larger e/e repulsion than is found for the larger Cl, Br or I atoms. This would suggest that the process for F should be less exothermic than for Cl and not fit the trend that shows a general decrease going down the group.

3) ΔhydH°(X, g) is the energy liberated on the hydration of the ion, the Hydration energy.
X(g) + H2O → X(aq)

The overall reaction is then:
½ X2(g) → X(aq)

Halogen atomisation energy
(kJ mol-1)
(kJ mol-1)
hydration enthalpy
(kJ mol-1)
(kJ mol-1)
F +79.08 -333 -504 -758
Cl +121.8 -348 -361 -587
Br +111.7 -324 -330 -542
I +106.7 -295 -285 -473


This shows a very negative energy change for the fluoride compared to the others in the group. This comes about because of two main factors: the high hydration energy and the low atomisation energy. For F2 2) is less than for Cl2 but since the energy needed to break the F-F bond is also less and the hydration more, the total energy drop is much greater. In spite of their lower atomisation energies, Br2 and I2 are weaker oxidising agents than Cl2 and this is due to their smaller ΔEAH1 and smaller ΔhydH°.

It can be seen that the ΔEAH1 value for fluorine is in between those for chlorine and bromine and so this value alone does not provide a good explanation for the observed variation.

Each of the halogens is able to oxidise any of the heavier halogens situated below it in the group. They can oxidise hydrogen and nonmetals such as:

X2 + H2(g) → 2HX(g)
In water, the halogens disproportionate according to:

X2 + H2O(l) → HX(aq) + HXO(aq), (where X=Cl, Br, I)

When base is added then the reaction goes to completion forming hypohalites, or at higher temperatures, halates, for example heating dichlorine:
3Cl2(g) + 6OH(aq) → ClO3(aq) + 5Cl−(aq) + 3H2O(l)


First Ionisation Energies

The trend seen for the complete removal of an electron from the gaseous halogen atoms is that fluorine has the highest IE1 and iodine the lowest. To overcome the attractive force of the nucleus means that energy is required so that the Ionisation Energies are all positive. The variation with size can be explained since as the size increases it take less energy to remove an electron. This inverse relationship is seen for all the groups, not just group 17. As the distance from the nucleus to the outermost electrons increases, the attraction decreases so that those electrons are easier to remove. The high value of IE1 for Fluorine is such that it does not exhibit any positive oxidation states, whereas Cl, Br and I can exist as high as 7.


Oxidation states

Fluorine is the most electronegative element in the periodic table and exists in all its compounds in either the -1 or 0 oxidation state. Chlorine, bromine, and iodine however can be found in a range of oxidation states including: +1, +3, +5, and +7, as shown below.


Common Oxidation States for the Halogens

Oxidation States Examples
-1 CaF2, HCl, NaBr, AgI
0 F2, Cl2, Br2, I2
1 HClO, ClF
3 HClO2ClF3
5 HClO3BrF5, [BrF6]IF5
7 HClO4BrF6+IF7, [IF8]

In general, odd numbered groups (like group 17) form odd-numbered oxidation states and this can be explained since all stable molecules contain paired electrons. (Free radicals are obviously much more reactive). When covalent bonds are formed or broken two electrons are involved so the oxidation state changes by 2.

When difluorine reacts with diiodine initially iodine monofluoride is formed. I2 + F2 → 2IF

Adding a second difluorine uses two more iodine valence electrons to form two more bonds:

2IF + F2 → IF3

more on this in the next lecture

Arsenic, a poison for mystery killing : it’s discovery



Arsenic compounds, namely its sulphides As2S3(orpiment) and As4S4(realgar or sandarac), were well known to Greeks and Romans. Orpiment was also known under the name of “arsenic”. Pliny the Elder and Dioscorides mentioned the toxicity of these compounds; Dioscorides noted calcination of “arsenic” to obtain white arsenic (oxide).

Arsenic is sometimes found in nature in native state and is fairly easily extracted from its compounds. It is not known who was the first to produce elemental arsenic. Usually its discovery is ascribed to the alchemist. Albert the great. Paracelsus described the process of preparing metallic arsenic by the calcination of “arsenic” with egg-shells. According to some reports, metallic arsenic was known much earlier but it was considered to be a variety of native mercury. This is due to the fact that arsenic sulphide resembles one of mercury minerals and the extraction of arsenic from its ores is rather simple.

In the Middle Ages arsenic was known not only in Europe but in Asia as well. Chinese alchemists could extract arsenic from its ores. Medieval Europians had no way of knowing whether death of a person was caused by arsenic poisoning but Chinese alchemists had a method of making sure. Unfortunately, their method of analysis is unknown. In Europe the test for estimating arsenic content in human body and the food eaten before death was developed by D. Marsh. This test is very sensitive and is still used.

Since arsenic sometimes accompanies tin, there are reported cases (for instance, in Chinese literature) when people were poisoned by water or wine kept for some time in new tin vessels.

For a long time people confused white arsenic, or its oxide, with arsenic itself believing the two to be the same substance. The confusion was eliminated at first by H. Brand and then by A. Lavoisier who proved that arsenic is an independent chemical element.

Arsenic oxide has for a long time been used to kill rodents and insects. The symbol As originates from the Latin word arsenicum whose etymology is obscure.

Predict the shape of the molecule using Steric Number

Steric number is the number of atoms bonded to a central atom of a molecule plus the number of lone pairs attached to the central atom.

The steric number of a molecule is used in VSEPR (valence shell electron pair repulsion) theory to determine the molecular geometry of a molecule.

Steric Number helps to predict the geometry and shape of the molecule easily.


Use the Lewis structure to determine the steric number. The steric number gives the electron-pair arrangement for the geometry that maximizes distance between valence electron pairs.

To minimize the repulsion between Valence shell electrons , the distance between them should be maximum.

When the distance between valence electrons is maximized, the energy of the molecule is at its lowest state.

When the Energy (potential energy) of the molecule is lowest , molecule is in its most stable arrangement.

The steric number is calculated using the following formula:

Steric Number = (number of lone electron pairs on the central atom) + (number of atoms bonded to the central atom)


Steric Number = number of lone pairs of electrons + number of σ (sigma) bonds formed by central atom

Steric Number Hybridization* Geometry Examples 
2 sp linear BeCl2
3 sp2 trigonal planar BF3
4 sp3 Tetrahedral CH4
5 sp3d Trigonal bipyramidal PCl5
6 sp3d2 Octahedral

or square bipyramidal

7 sp3d3 pentagonal bipyramidal IF7
1 pure s spherical H atom

* Steric number can be used to predict the hybridization. 

Examples for Calculation of Steric Number 

(i) Methane (CH4) – Methane consists of carbon bonded to 4 hydrogen atoms and 0 lone pairs. Steric numer is 4.

Water (H2O) – Water has two hydrogen atoms bonded to oxygen and also 2 lone pairs, so its steric number is 4.

Ammonia (NH3) – Ammonia also has a steric number of 4 because it has 3 hydrogen atoms bonded to nitrogen and 1 lone electron pair.


Element Mercury : it’s discovery


Mercury : liquid metal 

There is a science-fiction story by a Russian scientist I.A.EfremovThe Lake of the Mountain Spirits. Anybody who visited the lake in a sunny weather died. People living in the area were sure that the lake was inhabited with evil spirits who hated all visitors. When geologists reached the lake high in the mountains, they were amazed to learn that the lake contained not only water, but also native mercury element. And the “evil spirits” were nothing but element mercury vapour; in hot weather they rose above the surface of small and large mercury pools surrounding the lake.

Indeed, mercury is often found in native state, sometimes in most unexpected places. For instance, in some mountain regions of Spain, mercury was found at bottoms of wells. In antiquity mercury was known China and India. Mercury was also found in excavations of Egyptian tombs dating from about the Middle of the cinnabar was the only mercury containing mineral known in antiquity. Theophrastos (300 B.C.) described the process of extracting mercury from cinnabar by treating it with copper and vinegar. Man discovered mercury in ancient times owing to the fact that it is comparatively easily liberated from cinnabar at a sufficiently high temperature.

Mercury occurs in deposits throughout the world mostly as cinnabar (mercuric sulfide). The red pigment vermilion is obtained by grinding natural cinnabar or synthetic mercuric sulfide.

The world’s biggest mercury deposit is at Almaden (Spain). Exploitation of this deposit began at the time of the Roman Empire, and Romans extracted 4.5 tons of mercury annually. In antiquity mercury had many uses. Mirrors were made with amalgamated mercury; mercury and its compound were used as medicines. Cinnabar was mainly used as a pigment; and not for producing pure mercury. Before the invention of the galvanization process, mercury had been used in gilding and silvering processes. Amalgam of the metal was applied to a metal plate and heated to a high temperature. When mercury evaporated a thin coat of gold or silver remained on the plate. But this process was very unhealthy. Mercury played an important role in studies of gases; it was used in gas pumps and gas vessels.

Aristotle named mercury “liquid silver” and Dioskorides named in “silver water”. From this comes the Latin name of mercury –hydrargium.


Lead : It’s discovery


Lead discovery

Lead is very rarely encountered in a native state but is smelted fairly easily from ores. Lead become known to Egyptians simultaneously with iron and silver and was produced as early as the second millenium B.C. in India and china, In Europe production of lead began somewhat later although in the 6th century. B.C records we find mention of lead which was brought

Lead is a naturally occurring metal but its natural status doesn’t mean it’s healthy. In fact, lead is extremely toxic to humans and affects the liver, kidneys, reproductive system, and nervous system

to the Tyre trade fair. Lead was produced in great amounts during the reign of Hammurabi in Babylon. For a long time lead was confused with tin. Tin was named “plum bum album” and lead–“plumbum nigrum”. Only in the Middle Ages were they recognized as different metals.

Greeks and Phoenicians started many lead mines in Spain which later were taken over by Romans. In ancient Rome lead was widely used: for making crockery, styluses, and pipes for the famous Roman water-main. Lead was also used for manufacturing white lead. The island of Rhodes was the biggest exporter of white lead.

Lead Preparation :

The process of its preparation is still used as follows: lead pieces are immersed into vinegar and the salt thus obtained is boiled with water for a long time.

But red lead was first obtained unexpectedly. When a fire broke out in the Greek port of Piraeus barrels with lead were enveloped in flames. After the fire had been extinguished, red substance was found in the charred barrels–it was red lead.

Although in Russia lead has been known for a long time, up to the 18th century the process of lead production was very primitive. After the invention of firearms lead was used for making bullets and the military importance of lead is still great. But in addition to its “military” uses” lead has many peaceful ones; for instance, typographical types are made of its alloy with antimony. Lead is also used for protection against radiation in experiments.

Greeks named leadmolibdos; its chemical symbol Pb originates from Latin plumbum.



Silver Discovery

Silver is a more active metal than gold .

Although its abundance in the earth’s crust is about fifteen times that of gold, but it occurs much less frequently in a native state. It is not surprising that in antiquity silver was valued higher than gold. In ancient Egypt, for instance, the ratio between the costs of these metals was 2.5:1. Gold was used mainly for coins and jewelry; silver had other uses: for example, for making water vessels.

In the 4th century B.C the army of Alexander the Great conquered Persia and Phoenicia and invaded India. Here the Greek army was struck by an out break of a mysterious gastrointestinal disease and the men demanded to be sent home. Interestingly, the Greek military commanders fell victim to the disease far less frequently than their men, although they shared all the burdens of camp life with the soldiers. More than two thousand years had passed before scientists found an explanation of it. The soldiers drank from tin cups and their superiors from silver ones. It was proved that silver dissolves in water forming a colloid solution that kills pathogenic bacteria. And although the solubility of silver in in water in low, it is quite enough for disinfection.

Silver mines have known from time immemorial. The largest deposits of silver were in Greece, Spain, and Germany. After the discovery of America silver deposits were also found in Peru and Mexico. Lead minerals are often observed as constituents in silver ores.

Old process of Silver Extraction :

An old process of extracting silver from such ores is described as follows:-

Silver ore was ground, washed with water, and dried. Then it was fused together with flux and the alloy thus obtained was heated with charcoal. The resulting alloy of silver and lead was calcinated. On heating in air silver is practically unoxidized whereas lead transforms into oxide almost completely. The melting point of lead oxide is 896oC and that of silver, 960oC. Thus, practically pure silver was obtained.

At present more perfect processes of purifying silver are used.

Silver like gold was used in coins but the cost of silver compared to that of gold was gradually decreasing. In 1874 the cost of one pound of gold was equal to that of 15.5 pounds of silver but after the discovery of silver deposits in Australia this ratio fell to 1:46. In England bimetallism, i.e. the use of gold and silver jointly as a monetary standard, was discontinued in 1816. Later other countries followed this example.

In the last 40 years, silver has had two big bull runs, surrounded by even bigger bear markets. In the late 1970s, silver skyrocketed in price.

Russian words “rubl”’ (rouble) and “kopeika” (kopeck) owe their origin to silver. Rouble came into being in Kievan Russia in the 13th century –a silver bar weighing about 200 grams. It is believed that in the process of manufacturing roubles a long silver bar was cast and then hacked into parts (“rubit” is the Russian for “to hack”). The word “kopeika” appeared somewhat later (in 1534) when coins with an image of a horseman holding a speak (“kop’ e” in Russian) were first minted.

The name “silver” seems to stem from the Assyrian “serpu” or Gothic “silbur” the Latin argentum originates most likely from the Sanscrit arganta, which means “light, white”.

Polarizing Power and Fajan’s rule


It explains the generation of ionic character in a covalent bond.

A bond can be classified into Ionic bond and covalent bond using the  concepts of Effective Nuclear Charge and Electronegativity. 

Fajans’ Rules

Rules formulated by Kazimierz Fajans in 1923, can be used to predict whether a chemical bond is expected to be predominantly ionic or covalent, and depend on the relative charges and sizes of the cation and anion. If two oppositely charged ions are brought together, the nature of the bond between them depends upon the effect of one ion on the other.

non-polar covalent      polar covalent       ionic 


Fajan’s rules for predicting whether a bond is predominantly Covalent or Ionic

Covalent Ionic
Small cation (< ~100 pm) Large cation (> ~100 pm)
Large anion Small anion
High charges Low charges


Although the bond in a compound like X+Y may be considered to be 100% ionic, it will always have some degree of covalent character. When two oppositely charged ions (X+ and Y) approach each other, the cation attracts electrons in the outermost shell of the anion but repels the positively charged nucleus. This results in a distortion, deformation or polarization of the anion. If the degree of polarization is quite small, an ionic bond is formed, while if the degree of polarization is large, a covalent bond results.
The ability of a cation to distort an anion is known as its polarization power and the tendency of the anion to become polarized by the cation is known as its polarizability.

The polarizing power and polarizability that enhances the formation of covalent bonds is favoured by the following factors:

Small cation: the high polarizing power stems from the greater concentration of positive charge on a small area. This explains why LiBr is more covalent than KBr (Li+ 90 pm cf. K+ 152 pm).

Large anion: the high polarizability stems from the larger size where the outer electrons are more loosely held and can be more easily distorted by the cation. This explains why for the common halides, iodides, are the most covalent in nature (I 206 pm).

Large charges: as the charge on an ion increases, the electrostatic attractions of the cation for the outer electrons of the anion increases, resulting in the degree of covalent bond formation increasing.

Reminder. Large cations are to be found on the bottom left of the periodic table and small anions on the top right. The greater the positive charge, the smaller the cation becomes and the ionic potential is a measure of the charge to radius ratio.

On the left, the cation charge increases (size decreases) and on the right, the anion size increases, 
both variations leading to an increase in the covalency.
Thus covalency increases in the order:
[Na+ Cl, NaCl] < [Mg2+ 2(Cl), MgCl2] < [Al3+ 3(Cl), AlCl3] and
[Al3+ 3(F), AlF3] < [Al3+ 3(Cl), AlCl3] < [Al3+ 3(Br), AlBr3]

Electronic configuration of the cation: for two cations of the same size and charge, the one with a pseudo noble-gas configuration (with 18 electrons in the outer-most shell) will be more polarizing than that with a noble gas configuration (with 8 electrons in the outermost shell). Thus zinc (II) chloride ( Zn(II) 1s2 2s2 2p6 3s2 3p6 3d10 and Cl 1s2 2s2 2p6 3s2 3p6 ) is more covalent than magnesium chloride ( Mg(II) 1s2 2s2 2p6) despite the Zn2+ ion (74 pm) and Mg2+ ion (72 pm) having similar sizes and charges.

From an MO perspective, the orbital overlap disperses the charge on each ion and so weakens the electrovalent forces throughout the solid, this can be used to explain the trend seen for the melting points of lithium halides.

LiF = 870 °C, LiCl = 613 °C, LiBr = 547 °C, LiI = 446 °C

It is found that the greater the possibility of polarization, the lower is the melting point and heat of sublimation and the greater is the solubility in non-polar solvents.

Example: The melting point of KCl is higher than that of AgCl though the crystal radii of Ag+ and K+ ions are almost the same.

Solution: When the melting points of two compounds are compared, the one having the lower melting point is assumed to have the smaller degree of ionic character. In this case, both are chlorides, so the anion remains the same. The deciding factor must be the cation. (If the anions were different, then the answer could be affected by the variation of the anion.) Here the significant difference between the cations is in their electronic configurations. K+= [Ar] and Ag+ =[Kr] 4d10. This means a comparison needs to be made between a noble gas core and pseudo noble gas core, which as noted above holds that the pseudo noble gas would be the more polarizing.

Percentage of ionic character and charge distribution

Based on Fajan’s rules, it is expected that every ionic compound will have at least some amount of covalent character. The percentage of ionic character in a compound can be estimated from dipole moments.

The bond dipole moment uses the idea of electric dipole moment to measure the polarity of a chemical bond within a molecule. It occurs whenever there is a separation of positive and negative charges. The bond dipole μ is given by:

μ = δ d 

A bond dipole is modeled as +δ – δ- with a distance d between the partial charges. It is a vector, parallel to the bond axis and by convention points from minus to plus (note that many texts appear to ignore the convention and point from plus to minus). The SI unit for an electric dipole moment is the coulomb-meter, (C m). This is thought to produce values too large to be practical on the molecular scale so bond dipole moments are commonly measured in Debye, represented by the symbol, D. 

Historically the Debye was defined in terms of the dipole moment resulting from two equal charges of opposite sign and separated by 1 Ångstrom (10-10 m) as 4.801 D. This value arises from (1.602 x 10-19 * 1 x 10-10) / 3.336 x 10-30
where D = 3.336 x 10-30 C m (or 1 C m = 2.9979 x 1029 D).

Typical dipole moments for simple diatomic molecules are in the range of 0 to 11 D (see Table below). 
The % ionic character = μobserved / μcalculated (assuming 100% ionic bond) * 100 %

Example: From the Table below the observed dipole moment of KBr is given as 10.41 D, (3.473 x 10-29 coulomb metre), which being close to the upper level of 11 indicates that it is a highly polar molecule. The interatomic distance between K+ and Br is 282 pm. From this it is possible to calculate a theoretical dipole moment for the KBr molecule, assuming opposite charges of one fundamental unit located at each nucleus, and hence the percentage ionic character of KBr.

Solution: Dipole moment μ = q * e * d coulomb metre
q = 1 for complete separation of unit charge
e = 1.602 x 10-19 C
d = 2.82 x 10-10 m for KBr (282 pm)

Hence calculated μKBr = 1 * 1.602 x 10-19 * 2.82 x 10-10 = 4.518 x 10-29 Cm (13.54 D)

The observed μKBr = 3.473 x 10-29 Cm (10.41 D)

the % ionic character of KBr = 3.473 x 10-29/ 4.518 x 10-29 or 10.41 / 13.54 = 76.87% and the % covalent character is therefore about 23% (100 – 77).

Given the observed dipole moment is 10.41 D (3.473 x 10-29) it is possible to estimate the charge distribution from the same equation by now solving for q:

Dipole moment μ = q * e * d Coulomb metre 
but since q is no longer 1 we can substitute in values for μ and d to obtain an estimate for it.

q = μ /(e * d) = 3.473 x 10-29 / (1.602 x 10-19 * 2.82 x 10-10

thus q = 3.473 x 10-29 / (4.518 x 10-29) = 0.77 and the δ- and δ+ are -0.8 and +0.8 respectively.
Example. For HI, calculate the % of ionic character given a bond length = 161 pm and an observed dipole moment 0.44 D.

Solution: To calculate μ considering it as a 100% ionic bond

μ = 1 * 1.602 x 10-19 * 1.61 x 10-10 / (3.336 x 10-30) = 7.73 D

the % ionic character = 0.44/7.73 * 100 = 5.7%

The calculated % ionic character is only 5.7% and the % covalent character is (100 – 5.7) = 94.3%. The ionic character arises from the polarizability and polarizing effects of H and I. Similarly, knowing the bond length and observed dipole moment of HCl, the % ionic character can be found to be 18%. Thus it can be seen that while HI is essentially covalent, HCl has significant ionic character.

Note that by this simplistic definition, to achieve 100 % covalent character a compound must have an observed dipole moment of zero. Whilst not strictly true for heteronuclear molecules it does provide a simple qualitative method for predicting the bond character.

Bond character based on electronegativity differences

It is possible to predict whether a given bond will be non-polar, polar covalent, or ionic based on the electronegativity difference, since the greater the difference, the more polar the bond.


Electronegativity difference, ΔχP Bond
Δχ < 0.4 covalent
0.4 < Δχ < 1.7 polar covalent
Δχ > 1.7 ionic


Linus Pauling proposed an empirical relationship which relates the percent ionic character in a bond to the electronegativity difference.

percent ionic character= (1-e-(Δχ/2)^2 )* 100

This is shown as the curve in red below and is compared to the values for some diatomic molecules calculated from observed and calculated dipole moments.


diatomic Δ χ %ionic bond dist
Obs μ
Calc μ
Cl2 0.0 0.0 200 0.00 9.60
IBr 0.3 5.9 247 0.70 11.86
HI 0.4 5.7 161 0.44 7.73
ICl 0.5 5.4 232 0.60 11.14
HBr 0.7 12.1 141 0.82 6.77
HCl 0.9 17.7 127 1.08 6.10
ClF 1.0 11.2 163 0.88 7.83
BrF 1.2 15.1 178 1.29 8.55
LiI 1.5 65.0 238 7.43 11.43
HF 1.9 41.2 92 1.82 4.42
LiBr 1.8 69.8 217 7.27 10.42
KI 1.7 73.7 305 10.80 14.65
LiCl 2.0 73.5 202 7.13 9.70
KBr 2.0 76.9 282 10.41 13.54
NaCl 2.1 79.4 236 9.00 11.33
KCl 2.2 80.1 267 10.27 12.82
CsCl 2.3 74.6 291 10.42 13.97
LiF 3.0 86.7 152 6.33 7.30
KF 3.2 82.5 217 8.60 10.42
CsF 3.3 64.4 255 7.88 12.25



London dispersion forces
Intermolecular forces are the attractive forces between molecules without which all substances would be gases. The various types of these interactions span large differences in energy and for the halogens and interhalogens are generally quite small. The forces involved in these cases are called London dispersion forces (after Fritz Wolfgang London, 1900-1954). They are derived from momentary oscillations of electron charge in atoms and hence are present between all particles (atoms, ions and molecules).

The ease with which the electron cloud of an atom can be distorted to become asymmetric is termed the molecule’s polarizability. The greater the number of electrons an atom has, the farther the outer electrons will be from the nucleus, and the greater the chance for them to shift positions within the molecule. This means that larger nonpolar molecules tend to have stronger London dispersion forces. This is evident when considering the diatomic elements in Group 17, the Halogens. All of these homo-nuclear diatomic elements are nonpolar, covalently bonded molecules. Descending the group, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. For nonpolar molecules, the farther you go down the group, the stronger the London dispersion forces.

To picture how this occurs, compare the situation 1) where the electrons are evenly distributed and then consider 2) an instantaneous dipole that would arise from an uneven distribution of electrons on one side of the nucleus. When two molecules are close together, the instantaneous dipole of one molecule can induce a dipole in the second molecule. This results in synchronised motion of the electrons and an attraction between them. 3) when this effect is multiplied over numerous molecules the overall result is that the attraction keeps these molecules together, and for diiodine is sufficient to make this a solid.


On average the electron cloud for molecules can be considered to be spherical in shape. When two non-polar molecules approach, attractions or repulsions between the electrons and nuclei can lead to distortions in their electron clouds (i.e. dipoles are induced). When more molecules interact these induced dipoles lead to intermolecular attraction.

The changes seen in the variation of MP and BP for the dihalogens and binary interhalogens can be attributed to the increase in the London dispersion forces of attraction between the molecules. In general they increase with increasing atomic number.



Anomalous behaviour of the 2nd row elements: Li, Be, B, C, N, O, F

For the elements in the 2nd row, as the atomic number increases, the atomic radius of the elements decreases, the electronegativity increases, and the ionization energy increases.

The 2nd row has two metals (lithium and beryllium), making it the least metallic period and it has the most nonmetals, with four. The elements in the 2nd row often have the most extreme properties in their respective groups; for example, fluorine is the most reactive halogen, neon is the most inert noble gas, and lithium is the least reactive alkali metal.

These differences in properties with the subsequent rows are a result of:

  • the smaller size of the atoms
  • an outer shell with a maximum of 8 electrons (2s and 2p) and an underlying shell with just 2 electrons
  • no acessible d-orbitals – energy too high for use in bonding

Apart from the 2nd row (ignoring H/He 1st row) the later rows all end with inert gases but these do not have completed quantum levels. The 2nd row elements in general can only use the 2s and 2p electrons for bonding restricting the total number of bonds to 4.

So N is not expected to have more than 4 bonds and 3 is common, while for P 5 and 6 bonded species are quite common.
Reactivity of metals and metalloids
For Lithium, compared to other alkali metals

Reaction with water:
Li reacts slowly with water at 25 °C 
Na reacts violently and K in flames

2M(s) + 2H2O(l) → 2M+(aq) + 2OH + H2(g)

In general Li, Be, B, C, N, O, F are less reactive towards water than their heavier congeners.

Reaction with oxygen:
In conditions of excess oxygen, only Li forms a simple oxide, Li2O. Other metals form peroxides and superoxides 

Reaction with nitrogen:
Li reacts directly with N2 to form Li3N

6Li(s) + N2 (g) → 2Li3N(s). 

No other alkali metal reacts with N2


LiF, LiOH and Li2CO3 are less soluble than the corresponding Na and K compounds

For Beryllium compared to the other alkaline earth metals:

With water:

All Group 2 metals except Be, react with water

M(s) + 2H2O(l) → M2+(aq)+ 2OH(aq) + H2 (g)

With oxygen (air):
Be only reacts with air above 600 °C if it is finely powdered. The BeO that is formed is amphoteric (other Group 2 oxides are basic).

Of the Group 2 elements only Be reacts with NaOH or KOH to liberate H2 and form [Be(OH)4]2-.

Li and Be are metals but are less conducting than the higher members of Group 1 and 2 elements due to their high IEs (electrons are close to nucleus).

Ionization of Boron to B3+ requires a large input of energy and B adopts a covalent polymeric structure with semi-metallic properties.

The other elements of Group 14 become increasingly metallic as the group is descended due to the decrease in ionization energies.

Crystalline Boron is chemically inert – unaffected by boiling HCl and only slowly oxidized by hot concentrated HNO3 when finely powdered.

Covalent character
Li+ and Be2+ are small and have strong polarizing abilities. Their compounds are more covalent than those of the heavier elements in their groups.

BeCl2 is covalent while MCl2 (M = Mg-Ba) are ionic. The conductivity of fused beryllium chloride is only 1/1000 that of sodium chloride under similar conditions.

Catenation is the linkage of atoms of the same element into longer chains. Catenation occurs most readily in carbon, which forms covalent bonds with other carbon atoms to form longer chains and structures. This is the reason for the presence of the vast number of organic compounds in nature.

The ability of an element to catenate is primarily based on the bond energy of the element to itself, which decreases with more diffuse orbitals (those with higher azimuthal quantum number) overlapping to form the bond. Hence, carbon, with the least diffuse valence shell 2p orbital is capable of forming longer p-p sigma bonded chains of atoms than heavier elements which bond via higher valence shell orbitals.

Hetero-catenation is quite common in Inorganic Chemistry. Phosphates and silicates with P-O-P-O and Si-O-Si-O linkages are examples of this.


Multiple Bonds
C, N and O are able to form multiple bonds (double and/or triple). In Group 14, C=C double bonds are stable (134 pm) but Si=Si double bonds (227 pm) are uncommon. The diagram below shows how multiple bonds are formed involving π overlap of 2p orbitals. By comparison the 3p orbitals of the corresponding third row elements Si, P, and S are more diffuse and the longer bond distances expected for these larger atoms would result in poor π overlap.


C=C bond length = 134 pm and Si=Si bond length = 227 pm

Oxidizing ability of oxygen and fluorine
Due to the high electron affinities and electronegativities of oxygen and fluorine, they tend to form strong ionic bonds with other elements. They even react with noble gases to form compounds such as XeO3, XeO4, XeF4 and XeF6.

In 1962 Neil Bartlett at the University of British Columbia reacted platinum hexafluoride and xenon, in an experiment that demonstrated the chemical reactivity of the noble gases. He discovered the mustard yellow compound, xenon hexafluoroplatinate, which is perhaps now best formulated as a mixture of species, [XeF+][PtF5], [XeF+][Pt2F11], and [Xe2F3]+[PtF6]

A few hundred compounds of other noble gases have subsequently been discovered: in 1962 for radon, radon difluoride (RnF2), and in 1963 for krypton, krypton difluoride (KrF2). The first stable compound of argon was reported in 2000 when argon fluorohydride (HArF) was formed at a temperature of 40 K (-233.2 °C). Neutral compounds in which helium and neon are involved in chemical bonds have still not been formed. 

Noble gas compounds have already made an impact on our daily lives. XeF2 is a strong fluorinating agent and has been used to convert uracil to 5-fluorouracil, one of the first anti-tumor agents.




Discovery of Hydrogen : 

In 1671, Robert Boyle discovered and described the reaction between iron filings and dilute acids, which resulted in the production of hydrogen gas. In 1766-81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, and that it produced water when burned. He named it “flammable air”. In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek υδρο- hydro meaning “water” and -γενης genes meaning “creator”) when he and Pierre-Simon Laplace reproduced Cavendish’s finding that water was produced when hydrogen was burned. 
Hydrogen was liquefied for the first time by James Dewar in 1898 by using regenerative cooling and his invention, the vacuum flask. He produced solid hydrogen the next year. Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck. Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey’s group in 1932.

The nickel hydrogen battery was used for the first time in 1977 aboard the U.S. Navy’s Navigation technology satellite-2 (NTS-2). It had two caesium atomic clocks on board and helped to show that satellite navigation based on precise timing was possible. In the dark part of its orbit, the Hubble Space Telescope is powered by nickel-hydrogen batteries, which were finally replaced in May 2009, more than 19 years after launch, and 13 years passed their design life.

from NASA (accessed 2 Feb 2015)


Isotopes of hydrogen

Hydrogen has three naturally occurring isotopes, denoted 1H, 2H and 3H. Other, highly unstable nuclei (4H to 7H) have been synthesized in the laboratory but are not observed in nature.

  • 1H is the most common hydrogen isotope with an abundance of more than 99.98%. Because the nucleus of this isotope consists of only a single proton, it is given the descriptive, but rarely used formal name of protium.
  • 2H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. Essentially all deuterium in the universe is thought to have been produced at the time of the Big Bang, and has endured since that time. Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for 1H-NMR spectroscopy. Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.
  • 3H is known as tritium and contains one proton and two neutrons in its nucleus. It is radioactive, decaying into helium-3 through beta decay with a half-life of 12.32 years. It is sufficiently radioactive that it can be used in luminous paint, making it useful in such things as watches where the glass moderates the amount of radiation getting out. Small amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests. It is used in nuclear fusion reactions, as a tracer in isotope geochemistry, and specialized in self-powered lighting devices. Tritium has been used in chemical and biological labeling experiments as a radiolabel.

Hydrogen is the only element that has different names for its isotopes in common use today. During the early study of radioactivity, various heavy radioactive isotopes were given their own names, but these names are no longer used, except for deuterium and tritium.

nuclide symbol Z(p) N(n) isotopic mass (u) half-life decay mode Daughter Isotope representative isotopic composition
1H 1 0 1.00782503207(10) Stable 0.999885(70)
2H – D 1 1 2.0141017778(4) Stable 0.000115(70)
3H – T 1 2 3.0160492777(25) 12.32(2) y β 3He <1 in 1017 atoms

Properties of hydrogen

The difference of mass between isotopes of most elements is only a small fraction of the total mass and so this has very little effect on their properties, this is not the case for hydrogen. Consider chlorine with Z=17, there are 2 stable isotopes 35Cl (75.77%) and 37Cl (24.23%). The increase is therefore 2 in 35 or less than 6%. Deuterium and tritium are about double and triple the mass of protium and show significant physical and chemical differences particularly with regard to those properties related to mass, e.g. rate of diffusion, density, etc.

Some physical properties of the hydrogen isotopes.
isotope MP /K BP /K ΔHdiss /kJmol-1 Interatomic Distance /pm
H2 13.99 20.27 435.99 74.14
D2 18.73 23.67 443.4 74.14
T2 20.62 25.04 446.9 74.14

Differences between H2O and D2O


Property H2O D2O
Melting point /K 273.15 276.97
Boiling point /K 373.15 374.5
Temperature of maximum density /K 277 284.2
Maximum density /g cm3 0.99995 1.1053
Relative permittivity (at 298 K) 78.39 78.06
Kw (at 298 K) 1 *1014 2 * 1015
Symmetric stretch, ν1 /cm-1
(gaseous molecule)
3657 2671

Given that the boiling point of D2O is 101.4 °C (compared to 100.0 °C for H2O), evaporation or fractional distillation can be used to increase the concentration of deuterium in a sample of water by the selective removal of the more volatile light water, H2O. Thus bodies of water that have no outlet, such as the Great Salt Lake in Utah, USA and the Dead Sea in the Jordan Rift Valley, which maintain their level solely by evaporation, have significantly higher concentrations of deuterated water than do lakes or seawater with at least one outlet.

 The isotopic ratio for H and D is not fixed and so a range is given
for the standard atomic weight in the IUPAC Periodic Table of isotopes.

Heavy water is 10.6% denser than ordinary water, a difference not immediately obvious since they are otherwise physically and chemically similar. The difference can be observed by freezing a sample and dropping it into normal water, where it sinks.

With respect to taste and smell, rats given a choice between distilled normal water and heavy water avoided the heavy water, based on smell, and it may be that they detected a different taste as well.

The difference in weight increases the strength of water’s hydrogen-oxygen bonds, and this in turn is sufficient to cause differences that are important to some biochemical reactions. The human body naturally contains deuterium equivalent to about five grams of heavy water, which is harmless. When a large fraction of water (> 50%) in higher organisms is replaced by heavy water, the result is cell dysfunction and death.

In normal water, about 1 molecule in 3,200 is HDO (one hydrogen in 6,400 is in the form of D), and heavy water molecules (D2O) only occur in a proportion of about 1 molecule in 41 million (i.e. one in 6,4002). Thus semiheavy water molecules are far more common than “pure” (homoisotopic) heavy water molecules.

Deuterium oxide was initially obtained by the electrolysis of ordinary water over a considerable period of time. This method of production requires a large cascade of stills or electrolysis chambers and consumes large amounts of power, so that chemical methods are generally now preferred. The most important chemical method is the Girdler sulfide process.

In this process, demineralised and deaerated water is trickled through a series of perforated (seive) plates in a tower, while hydrogen sulfide gas (BP -60 °C) flows upward through the perforations. Deuterium migration preferentially takes place from the gas to the liquid water. This “enriched” water from the cold tower (maintained at 32 °C) is fed to the hot tower (at 130 °C) where deuterium transfer takes place from the water to the hydrogen sulfide gas. An appropriate “cascade” setup accomplishes enrichment via the reversible reaction: 

H2O +HDS ⇄ HDO + H2

The equilibrium constant, K for, this reaction in terms of the concentrations, can be written as:

K = ([HDO][H2S]) / ([H2O][HDS]) or alternatively:

K = ([HDO]/[H2O]) / ([HDS]/[H2S])

If H and D were the same chemically, the equilibrium constant for the reaction would be equal to unity. However, what is found is that K is not equal to unity, and furthermore it is temperature dependent: 

at 25 °C,  K = 2.37
at 128 °C, K = 1.84

From the above information, at 32 °C, the equilibrium favours the concentration of deuterium in water. However, at around 130 °C, the equilibrium is now relatively more favorable to the concentration of deuterium in the hydrogen sulfide. In other words, the concentration of HDO in H2O is greater than the concentration of HDS in H2S but the relative concentration of HDS in H2S increases with increasing temperature, making it possible to separate D from H.

In the first stage, the gas is enriched from 0.015% deuterium to 0.07%. The second column enriches this to 0.35%, and the third column achieves an enrichment between 10% and 30% deuterium oxide, D2O. Further enrichment to “reactor-grade” heavy water (> 99% D2O) still requires distillation or electrolysis. The production of a single litre of heavy water requires ~340,000 litre of feed water.

In 1934, Norway built the first commercial heavy water plant with a capacity of 12 tonnes per year. From 1940 and throughout World War II, the plant was under German control and the Allies decided to destroy the plant and its heavy water to inhibit German development of nuclear weapons. In late 1942, a planned raid by British airborne troops failed, both gliders crashing. The raiders were killed in the crash or subsequently executed by the Germans. On the night of 27 February 1943 Operation Gunnerside succeeded. Norwegian commandos and local resistance managed to demolish small, but key parts of the electrolytic cells, dumping the accumulated heavy water down the factory drains. Had the German nuclear program followed similar lines of research as the United States Manhattan Project, the heavy water would not have been crucial to obtaining plutonium from a nuclear reactor, but the Germans did not discover the graphite reactor design used by the allies for this purpose.

ortho- and para-dihydrogen

In dihydrogen, the two electrons in the molecule will be spin paired but there is no similar requirement for the two nuclei; they may be parallel or opposed. There are therefore two nuclear spin isomers possible which are called ortho and para.

In the parahydrogen form the nuclear spins of the two protons are antiparallel and form a singlet (2I+1= 1) with a molecular spin quantum number, I, of 0 (½ – ½). In the orthohydrogen form, the spins are parallel and form a triplet state (2I+1= 3) with a molecular spin quantum number, I, of 1 (½ + ½). At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the “normal form”.

The para form has slightly lower energy:-

o-H2 ⇄ p-H2; ΔH = -1.5 kJmol-1

but due to the small difference this has little effect at room temperature.

The amount of ortho and para hydrogen varies with temperature:

  • At 20 K, hydrogen contains mainly para (singlet) hydrogen (99.8%) which is the more stable form.
  • At the temperature of liquefaction of air, ~80 K, the ratio of ortho and para hydrogen is 1 : 1.
  • At room temperatures, the ratio of ortho to para hydrogen is 3 : 1.
  • Even at very high temperatures, the ratio of ortho to para hydrogen never exceeds 3 : 1.

It is possible then to get pure para hydrogen by cooling ordinary hydrogen gas to very low temperatures (close to 20 K) but it is not possible to get a sample of hydrogen containing more than 75% of ortho (triplet) hydrogen. The first synthesis of pure parahydrogen was achieved in 1929.
This conversion of ortho- to para-hydrogen liberates some heat which can cause evaporation of hydrogen within storage vessels. Since orthohydrogen molecules make up 75% of “normal” hydrogen at room temperature, this can considerably complicate the performance of storing liquid hydrogen. Without an ortho-para conversion catalyst, (such as hydrous ferric oxide) extra refrigeration equipment is required to remove the heat generated by the natural conversion to para hydrogen.

Production of hydrogen

Laboratory preparations
In the laboratory, H2 can be prepared by the action of a dilute non-oxidizing acid on a reactive metal such as zinc, with a Kipp’s apparatus. 

Zn + 2Haq+ ⇄ Znaq2+ + H2

Aluminium can produce H2 upon treatment with bases:

2Al + 6 H2O + 2 OH ⇄ 2 Al(OH)4 + 3 H2

The electrolysis of water is another simple method of producing hydrogen. A low voltage current is passed through the water, and gaseous dioxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or other inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used versus energetic value of hydrogen produced) is in the range 80-94%.

2 H2O(l) ⇄ 2 H2(g) + O2(g) 

In 2007, it was discovered that an alloy of aluminium and gallium in pellet form added to water could be used to generate hydrogen. The process creates alumina, but the expensive gallium, which prevents the formation of an oxide skin on the pellets, can be re-used. This has important potential implications for a hydrogen economy, as hydrogen could be produced on-site without the need of being transported.

Industrial preparation of hydrogen
Steam reforming is a method for producing hydrogen, carbon monoxide or other useful products from hydrocarbon fuels such as natural gas. This is achieved in a processing device called a reformer which reacts steam at high temperature with the fossil fuel.

At high temperatures (700 – 1100 °C) and in the presence of a metal-based catalyst (nickel), steam reacts with methane to yield carbon monoxide and hydrogen.

CH4 + H2O → CO + 3 H2

In order to produce more hydrogen from this mixture, more steam is added and the water gas shift reaction is carried out: 

CO + H2O → CO2 + H2

The mixture of CO and H2 is called “synthesis gas or syngas”. Syngas is used as an intermediate in producing synthetic petroleum for use as a fuel or lubricant via the Fischer-Tropsch process and previously the Mobil methanol to gasoline process.

Enzymatic route from xylose 
In 2013 a low-temperature, 50 °C, atmospheric-pressure, enzyme-driven process to convert xylose into hydrogen with nearly 100% of the theoretical yield was announced. The process employed 13 enzymes, including a novel polyphosphate xylulokinase (XK).

It was noted that: “Approximately 50 million metric tons of dihydrogen are produced annually from nonrenewable natural gas, petroleum, and coal. H2 production from water remains costly. Technologies for generating H2 from less costly biomass, such as microbial fermentation, enzymatic decomposition, gasification, steam reforming, and aqueous phase reforming, all suffer from low product yields.

Compounds of Hydrogen

The chemistry of hydrogen depends mainly on four processes:

  1. donation of the valency electron to form the hydrogen ion, H+
  2. accepting an electron to form the hydride ion H
  3. sharing the electron with a partner atom to form a pair bond (covalent bond) H-H
  4. sharing the electron with an ensemble of atoms to form a metallic bond H.

While H2 is not very reactive under standard conditions, it does form compounds with most elements. Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I), or oxygen; in these compounds hydrogen takes on a partial positive charge. When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of medium-strength noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules. Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides.

The term “hydride” suggests that the H atom has acquired a negative or anionic character, denoted H-, and is used when hydrogen forms a compound with a more electropositive element. The existence of the hydride anion, suggested by Gilbert N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 by the electrolysis of molten lithium hydride (LiH), producing a stoichiometry quantity of hydrogen at the anode.

Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminium hydride. A simple binary indium hydride has not yet been identified, although larger complexes exist.

The position of H in the Periodic Table

In some respects, H does not seem to have a perfect position in the Periodic Table and so many designers have it in more than one position, e.g. in Group 1 or Group 17 and even in Group 14.

Ionization energy of hydrogen

Hydrogen has a single outer electron, like the alkali metals, but they all form positive ions quite readily whereas hydrogen has little tendency to do so. Hydrogen often tends to share its electron with nonmetals rather than losing it to them.

The first ionization energies for H, Li, Na and K are 1312, 520.2, 495.8 and 418.8 kJmol-1. The high IE for H (even bigger than for Xe) can be attributed to the very small size of the atom and the strong attractive force between the proton and electron.

H(g) → H+(g) + e        ΔH = 1312 kJmol-1

Xe(g) → Xe+(g) + e      ΔH = 1170 kJmol-1

The free proton can only be obtained under extreme conditions such as by an electric arc or in a discharge tube and even then only exists for about half a second. H+ can be found in solvated form where the solvation energy provides the energy needed to overcome the very high ionization energy. Examples are in ammonia, alcohol or water with species like NH4+, ROH2+ and H3O4+ being formed.

Electron affinity of hydrogen

Hydrogen, like the halogens, exists as diatomic molecules and H atoms have electron configurations with one electron short of a filled outer shell hence the idea of placing H in Group 17. However unlike the halogens with large EA values, the EA for hydrogen is quite small. The formation of H is much less favourable than the formation of a chloride ion, as seen from the thermodynamic profiles below and it is rare whereas halide ions are common and stable. In addition H has a lower electronegativity value than any of the halogens.

Much more energy is required as well to break the H-H bond compared to the Cl-Cl bond where the steps for comparison are:

½H2 (g) → H. (g)                   ΔH = 218 kJmol-1
H. (g) + e → H (g)                ΔH = -72.8 kJmol-1
so overall for hydrogen
½H2 (g) + e → H (g)            ΔH = +145.2 kJmol-1
½Cl2 (g) → Cl. (g)                 ΔH = 121 kJmol-1
Cl. (g) + e → Cl (g)              ΔH = -348.6 kJmol-1
overall for chlorine
½Cl2 (g) + e → Cl (g)          ΔH = -227.6 kJmol-1

As a result, only the most active elements, whose Ionization Energies are low, can form ionic hydrides, e.g. NaH.

The covalent radius for H is 37 pm and the estimated radius for H is ~140 pm indicating a substantial increase. This comes about as a result of the interelectronic repulsion when a second electron is added to the 1s atomic orbital. All the alkali metal hydrides crystallize with the NaCl-type structure and are all considered ionic. They are sometimes called “saline” hydrides.

Saline hydrides

The instability of the hydride ion compared to the halide ions can be seen by comparison of the ΔHf for alkali metal hydrides and chlorides.

Cation ΔHf MH/ kJmol-1 ΔHf MCl/ kJmol-1
Li -90.5 -409
Na -56.3 -411
K -57.7 -436
Rb -52.3 -430
Cs -54.2 -433

Saline hydrides are formed by the group 1 and 2 metals when heated with dihydrogen (H2). They are white, high melting point solids that react immediately with protic solvents, for example:

NaH + H2O → NaOH + H2

(Their moisture sensitivity means that reaction conditions must be water-free.)

Evidence for the ionic nature of these hydrides is:
1) molten salts show ionic conductivity.
2) X-ray crystal data gives reasonable radius ratios expected for ionic compounds.
3) Observed and calculated Lattice Energies (from Born-Haber cycles etc.) are in good agreement (i.e. show little covalency).
NaH is capable of deprotonating a range of even weak Brønsted acids to give the corresponding sodium derivatives.

NaH + Ph2PH → Na[PPh2] + H2

Sodium hydride is sold by many chemical suppliers as a mixture of 60% sodium hydride (w/w) in mineral oil. Such a dispersion is safer to handle and weigh than pure NaH. The compound can be used in this form but the pure grey solid can be prepared by rinsing the oil with pentane or tetrahydrofuran, THF, care being taken because the washings will contain traces of NaH that can ignite in air. Reactions involving NaH require an inert atmosphere, such as nitrogen or argon gas. Typically NaH is used as a suspension in THF, a solvent that resists deprotonation but solvates many organosodium compounds.

Hydride reducing agents

LiH and Al2Cl6 gives lithium aluminium hydride (lithal LiAlH4), NaH reacts with B(OCH3)3 to give sodium borohydride (NaBH4). These find wide scope and utility in organic chemistry as reducing agents.

LiAlH4 is commonly used for the reduction of esters and carboxylic acids to primary alcohols; previously this was a difficult conversion that used sodium metal in boiling ethanol (the Bouveault-Blanc reduction). The solid is dangerously reactive toward water, releasing gaseous hydrogen (H2). Some related derivatives have been discussed for hydrogen storage.

NaBH4 is used in large amounts for the production of sodium dithionite from sulfur dioxide: Sodium dithionite is used as a bleaching agent for wood pulp and in the dyeing industry. NaBH4 consists of the tetrahedral BH4 anion in the crystalline form and is found to exist as three polymorphs: α, β and γ. The stable phase at room temperature and pressure is α-NaBH4, which is cubic and adopts an NaCl-type structure. Millions of kilograms are produced annually, far exceeding the production levels of any other hydride reducing agent.

NaBH4 will reduce many organic carbonyls, depending on the precise conditions. Most typically, it is used in the laboratory for converting ketones and aldehydes to alcohols. For example, reduction of acetone (propanone) to give propan-2-ol.

Molecular hydrides – covalent hydrides and organic compounds

Hydrogen forms a vast number of compounds with carbon, (the hydrocarbons), and an even larger array with heteroatoms that, because of their general association with living things, are called organic compounds. The study of their properties is covered in organic chemistry and their study in the context of living organisms is covered in biochemistry. By some definitions, “organic” compounds are only required to contain carbon. However, most of them also contain hydrogen, and because it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word “organic” in chemistry. Millions of hydrocarbons are known, and they are usually formed by complicated synthetic pathways, which seldom involve direct reaction with elementary hydrogen.

Most molecular hydrides are volatile and many have simple structures that can be predicted by the VSEPR model. There are a large number of B hydrides known (boranes) and although the simplest BH3 has been found in the gas phase it readily dimerises to give B2H6

In inorganic chemistry, hydrides can serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminium complexes, as well as in clustered carboranes, (composed of boron, carbon and hydrogen atoms). The bonding of the bridging hydrogens in many of the boranes is explained in terms of 3 centre – 2 electron bonds.

Diborane is a colourless and highly unstable gas at room temperature with a repulsively sweet odour. Diborane mixes well with air, easily forming explosive mixtures. Diborane will ignite spontaneously in moist air at room temperature.


Metallic (interstitial) hydrides

Many transition metal elements form metallic (interstitial) hydrides, in which H2 molecules (and H atoms) can occupy the holes in the metal’s crystal structure. They are traditionally termed ‘compounds’, even though they do not strictly conform to the definition of a compound; more closely resembling common alloys such as steel. These systems are usually non-stoichiometric, with variable amounts of hydrogen atoms in the lattice.

Palladium is unique in its ability to reversibly absorb large amounts of H2 or D2 (up to 900 times its own volume of hydrogen, but no other gases, at room temperature) to form palladium hydride. Structural studies show that the absorbed H fits into octahedral holes in the cubic close packed Pd lattice with a non-stoichiometric formula approximating to PdH0.6 for the β-form. This material has been considered as a means to carry hydrogen for vehicular fuel cells. Interstitial hydrides show some promise as a way for safe hydrogen storage. During the last 25 years many interstitial hydrides have been developed that readily absorb and discharge hydrogen at room temperature and atmospheric pressure. At this stage their application is still limited, as they are capable of storing only about 2 weight percent of hydrogen, insufficient for automotive applications.

Hydrogen bonds

A hydrogen bond is the name given to the electrostatic attraction between polar molecules that occurs when a hydrogen (H) atom bound to a highly electronegative atom such as nitrogen (N), oxygen (O) or fluorine (F) experiences attraction to some other nearby highly electronegative atom. The name is something of a misnomer, as it represents a particularly strong dipole-dipole attraction, rather than a typical covalent bond.

The 2011 IUPAC definition specifies that “The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X-H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation.

These hydrogen-bond attractions can occur between molecules (intermolecular) or within different parts of a single molecule (intramolecular). The hydrogen bond (5 to 30 kJ/mole) is stronger than a van der Waals interaction, but weaker than covalent or ionic bonds. This type of bond can occur in inorganic molecules such as water and in organic molecules like DNA and proteins.

Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides that have no hydrogen bonds. Intramolecular hydrogen bonding is partly responsible for the secondary and tertiary structures of proteins and nucleic acids. It plays an important role in the structure of polymers, both synthetic and natural.

BP’s of MG hydrides with Noble gases for comparison /K

Hydrogen bonding in biological systems.

Base pairs, which form between specific nucleobases (also termed nitrogenous bases), are the building blocks of the DNA double helix and contribute to the folded structure of both DNA and RNA. Dictated by specific hydrogen bonding patterns, Watson-Crick base pairs (guanine-cytosine and adenine-thymine) allow the DNA helix to maintain a regular helical structure that is subtly dependent on its nucleotide sequence. The complementary nature of this based-paired structure provides a backup copy of all genetic information encoded within double-stranded DNA. The regular structure and data redundancy provided by the DNA double helix make DNA well suited to the storage of genetic information, while base-pairing between DNA and incoming nucleotides provides the mechanism through which DNA polymerase replicates DNA, and RNA polymerase transcribes DNA into RNA. Many DNA-binding proteins can recognize specific base pairing patterns that identify particular regulatory regions of genes.

Applications of hydrogen

Large quantities of H2 are used by the petroleum and chemical industries. The largest application of H2 is for the processing (“upgrading”) of fossil fuels, and in the production of ammonia. The key consumers of H2 in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H2 has several other important uses. H2 is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H2 is used as a reducing agent of metallic ores.

Nitrogen is a strong limiting nutrient in plant growth. Carbon and oxygen are also critical, but are more easily obtained by plants from soil and air. Even though air is 78% nitrogen, atmospheric nitrogen is nutritionally unavailable because nitrogen molecules are held together by strong triple bonds. Nitrogen must be ‘fixed’, i.e. converted into some bioavailable form, through natural or man-made processes. It was not until the early 20th century that Fritz Haber developed the first practical process to convert atmospheric nitrogen to ammonia, which is nutritionally available.

Fertilizer generated from ammonia produced by the Haber process is estimated to be responsible for sustaining one-third of the Earth’s population. It is estimated that half of the protein within human beings is made of nitrogen that was originally fixed by this process; the remainder was produced by nitrogen fixing bacteria and archaea.

Dozens of chemical plants worldwide produce ammonia, consuming more than 1% of all man-made power. Ammonia production is thus a significant component of the world energy budget. Modern ammonia-producing plants depend on industrial hydrogen production to react with atmospheric nitrogen using a magnetite catalyst or over a promoted Fe catalyst under high pressure (100 standard atmospheres (10,000 kPa)) and temperature (450 °C) to form anhydrous liquid ammonia. This step is known as the ammonia synthesis loop (also referred to as the Haber-Bosch process):

3 H2 + N2 ⇄ 2 NH3 (ΔH = -92.4 kJmol-1)

Nitrogen (N2) is very unreactive because the molecules are held together by strong triple bonds. The Haber process relies on catalysts that accelerate the cleavage of this triple bond.

At room temperature, the equilibrium is strongly in favor of ammonia, but the reaction doesn’t proceed at a detectable rate. Thus two opposing considerations are relevant to this synthesis. One possible solution is to raise the temperature, but because the reaction is exothermic, the equilibrium quickly becomes quite unfavourable at atmospheric pressure. Low temperatures are not an option since the catalyst requires a temperature of at least 400 °C to be efficient. By increasing the pressure to around 200 atm the equilibrium concentrations are altered to give a profitable yield.

The reaction scheme, involving the heterogeneous catalyst, is believed to involve the following steps:

1.   N2 (g) → N2 (adsorbed)

2.   N2 (adsorbed) → 2 N (adsorbed)

3.   H2(g) → H2 (adsorbed)

4.   H2 (adsorbed) → 2 H (adsorbed)

5.   N (adsorbed) + 3 H (adsorbed) → NH3 (adsorbed)

6.   NH3 (adsorbed) → NH3 (g)

Reaction 5 actually consists of three steps, forming NH, NH2, and then NH3. Experimental evidence suggests that reaction 2 is the slow, rate-determining step. This is not unexpected given that the bond broken, the nitrogen triple bond, is the strongest of the bonds that must be broken.



Periodic Properties


Trends across a period follow from the increasing number of protons in the nucleus and the decrease in radius. Both contributions can be explained by the change in effective nuclear charge.

Trends down a group follow from the increasing number of electron shells and the increased distance of the outer electrons from the nucleus. The major factor is the increasing size.

The properties of an element are largely determined by their electronic configurations, giving rise to recurring patterns or periodic behaviour. Examples are shown in the diagrams below including ionization energy, electron affinity, electronegativity and atomic radii. It is this periodicity of properties, manifestations of which were noticed well before the underlying theory was known, that led to the establishment of the periodic law (the properties of the elements recur at varying intervals) and the development of the first periodic tables. The modern periodic table is a tabular arrangement of the chemical elements, organized on the basis of their atomic number (number of protons in the nucleus), electronic configurations, and recurring chemical properties.

In the RSC Tutorial Chemistry Text on Main Group Chemistry, it notes that “When an element forms a chemical compound, electrons can be considered to be either lost, gained or shared with other atoms. These tendencies can be assessed by the parameters of ionization energy (IE), electron affinity (EA) and electronegativity (E). Prediction of bond types as either ionic or covalent then allows prediction of the chemical and physical properties of chemical substances.” 

So how are these parameters defined and how do they vary with atomic number?

Effective nuclear charge

The concept of the effective nuclear charge (often symbolized as Zeff or Z*) relates to the net positive charge experienced by an electron in a multi-electron atom. The term “effective” is used because the shielding effect of negatively charged inner electrons prevents higher orbital electrons from experiencing the full nuclear charge due to the repelling effect of the lower inner-layer electrons. 

In an atom with one electron, that electron experiences the full charge of the positive nucleus. In this case, the effective nuclear charge can be calculated from Coulomb’s law.

However, in an atom with many electrons the outer electrons are simultaneously attracted to the positive nucleus and repelled by the inner negatively charged electrons. The effective nuclear charge on such an electron is given by the following equation:

Zeff = Z – S

Z is the number of protons in the nucleus (atomic number), and
S is the shielding calculated from the electrons between the nucleus and the electron in question. A systematic method for determining this is given by “Slater’s rules”.

These can be summarised as follows:
Arrange the electrons into a sequence of groups in order of increasing principal quantum number n, and for equal n in order of increasing azimuthal quantum number l, except that s- and p- orbitals are kept together.
[1s] [2s,2p] [3s,3p] [3d] [4s,4p] [4d] [4f] [5s, 5p] [5d] etc.

Any electron higher in the sequence to the electron under consideration contributes nothing to the shielding, S, and is ignored.

For an electron in an ns or np orbital:

0.35 comes from each other electron within the same group except for the [1s] group, where the other electron contributes only 0.30.

0.85 for each electron with principal quantum number n one less than that of the group, i.e (n-1) shell

1.00 for each electron with principal quantum number two or more less, i.e (n-2) etc. shell

For an electron in an nd or nf orbital:
0.35 comes from each other electron within the same group

1.00 for each electron “closer” to the atom than the group. This includes electrons with the same principal quantum number but in s or p orbitals.
In tabular form, the rules are summarized as:

Group Other electrons in the same group Electrons in group(s) with principal quantum number n and azimuthal quantum number < l Electrons in group(s) with principal quantum number n-1 Electrons in all group(s) with principal quantum number < n-1
[1s] 0.30
[ns, np] 0.35 0.85 1
[nd] or [nf] 0.35 1 1 1


Example 1.
Consider a sodium cation, Na+, a fluorine anion, F, and a neutral neon atom, Ne. Each has 10 electrons, 1s2 2s2 2p6 so the shielding from the 1s and 2s/2p electrons is 2 * 0.85 + 7 * 0.35 = 4.15 but the effective nuclear charge varies because each has a different atomic number:

Zeff F = 9 – 4.15 = 4.85
Zeff Ne = 10 – 4.15 = 5.85
Zeff Na+ = 11 – 4.15 = 6.85

So the sodium cation has the largest effective nuclear charge, and can be expected to have the smallest radius.

Example 2.

Predict whether K would be more energetically stable with a configuration of 
1s2 2s2 2p6 3s2 3p6 4s1        or        1s2 2s2 2p6 3s2 3p6 3d1

For K, Z=19 and considering the 4s electron then the screening constant S can be calculated from:

S= (8 * 0.85) + (10 * 1.0) = 16.8

Zeff = 19 – 16.8 = 2.2

For the 3d calculation of S:

S= (18 * 1.0) = 18

Zeff = 19 – 18 = 1

Accordingly, an electron in the 4s (as opposed to the 3d) orbital would come under the influence of a greater effective nuclear charge in the ground state of potassium and so will be the orbital that is occupied.

Ionization Energy

The Ionization Energy (IEn) of an element is defined as the internal energy change associated with the removal of an electron from the gaseous atom, E, in its ground state, i.e. at 0 K. The first IE is therefore the energy required for the reaction:
E(g) → E+(g) + e   energy required = IE 

This energy change is generally considered equivalent to the enthalpy change at 298 K (ΔH298 K). Estimates of the error suggest < 10 kJmol-1 which when compared to typical IE values often in their thousands, is insignificant.

The diagrams above show the variation in the values of the 1st five IE’s as a function of Z up to Nd (60). 

Features to note for IE1 are:

  • the values associated with the noble gases are the highest in each period
  • the values associated with the group 1 elements are generally the lowest in each period, (group 2 elements for the 2nd IE and group 3 for the 3rd IE etc.)
  • the gradual increase in values across a given period (applies to IE2-5 as well)
  • the drop in values on going from an element in group 15 to its neighbour in group 16 e.g. for N-O, P-S, As-Se
  • the drop in values on going from an element in group 2 or 12 to its neighbour in group 13, e.g. for Be-B, Mg-Al and Zn-Ga, Cd-In
  • the rather similar values for a given row of d-block elements

These observations can be accounted for in terms of the effective nuclear charge since the further away from the positively charged nucleus that a negatively charged electron is located, the less strongly that electron is attracted to the nucleus and the more easily it can be removed. So, as the atomic radius decreases from left to right across the Period the 1st Ionization Energy increases.


Electron Affinity

The electron affinity(EA) of an element E is defined as minus the internal energy change associated with the gain of an electron by a gaseous atom, at 0 K :
E(g) + e → E(g)      energy change = ΔEAU(0 K)

This energy change is generally considered to be equivalent to ΔH298 K, so:

EA = – ΔH298 K

Unlike ionization energies, which are always positive for a neutral atom because energy is required to remove an electron, electron affinities can be positive (energy is released when an electron is added), negative (energy must be added to the system to produce an anion), or zero (the process is energetically neutral).

Chlorine has the most positive electron affinity of any element, which means that more energy is released when an electron is added to a gaseous chlorine atom than to an atom of any other element, EA= 348.6 kJmol-1 and the group 17 elements have the largest values overall. The addition of a second electron to an element is expected to be much less favoured since there will be repulsion between the negatively charged electron and the overall negatively charged anion. For example, for O the values are:

O(g) + e → O(g)           EA = + 141 kJmol-1
O(g) + e → O2-(g)        EA = – 798 kJmol-1


The concept of Electronegativity originated with Linus Pauling in the 1930’s and was defined as “the power of an atom in a molecule to attract electrons to itself”.

The values proposed by Pauling were calculated based on differences in bond dissociation enthalpy values found when comparing homo-diatomic molecues with hetero-diatomic molecules. For example , the bond energy of chlorine monofluoride, ClF, is about 255 kJ mol-1 which is significantly greater than for either of the two homo-nuclear species Cl2 and F2 (242 and 153 kJ mol-1 respectively). Pauling attributed this to an electrostatic attraction between the partially charged atoms in the heternuclear species. That is the excess bond energy came from an ionic contribution to the bond.

The method of calculating the Pauling values is:

D(XY) = [D(XX).D(YY)]1/2 + 96.48 * (χY – χX)2      where the 96.48 factor means D values are in kJ mol-1

In Housecroft and Sharpe the average, rather than geometric mean is used, and this is rearranged to give:

ΔD = D(XY) – {½ * [D(XX) – D(YY)] } = (χY – χX)2 = (Δχ)2

or Δχ = √(ΔD)

As only differences in electronegativity were defined, it was necessary to choose an arbitrary reference point in order to construct a scale. Hydrogen was chosen as the reference, since it formed covalent bonds with a large variety of elements: its electronegativity was fixed at 2.20.

The Mulliken scale was calculated by taking the average of the Ionization Energy and the Electron Affinity (when both were given in units of eV).

χM = ½ (IE1 + EA1)      where both IE1 and EA1 are in eV

A variant of this (2006) that converts the values to roughly the Pauling scale is:

χM = 0.00197 * (IE1 + EA1) + 0.19 where IE1 and EA1 are now given in kJ mol-1 

The plots above indicate that while the absolute values are different, the trends are quite similar and the 2 curves are comparable when scaled appropriately.

The Allred-Rochow scale considered that electronegativity was related to the charge experienced by an electron on the “surface” of an atom: the higher the charge per unit area of atomic surface the greater the tendency of that atom to attract electrons. Their scale was dependent on Zeff and inversely proportional to the square of the covalent radius, rcov.

χAR = (3590 * Zeff / r2cov) + 0.744      where rcov is in pm

The values range between 0 and 10. Once again a good correlation to the Pauling scale was found and this applies as well to other Electronegativity scales.

Atomic Radius

Have a look at an interactive visual display (JSmol) showing the 

periodic table of elements with atomic and ionic radii.

The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the center of the nucleus to the boundary of the surrounding cloud of electrons. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Three widely used definitions of atomic radius are Van der Waals radius, ionic radius, and covalent radius.

Covalent radius is defined as half the covalent bond length when the two atoms bonded are homonuclear (½ X-X bond).

Van der Waals radius is defined as half of the internuclear separation of two non-bonded atoms of the same element on their closest possible approach.

It is not possible to measure the sizes of both metallic and nonmetallic elements using a single technique and method. To get values for comparison, theoretical quantum mechanical functions have been used instead to calculate atomic radii.

In the periodic table, atomic radii decrease from left to right across a period and increase from top to bottom down the groups. As a result of these two trends, the largest atoms are found in the lower left corner of the periodic table, and the smallest are found in the upper right corner.

The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory. The atomic radii decrease across the Periodic Table because as the atomic number increases, the number of protons increases across the period, but the extra electrons are only added to the same quantum shell. Therefore, the effective nuclear charge towards the outermost electrons increases, drawing the outermost electrons closer. As a result, the electron cloud contracts and the atomic radii decreases.

The Lanthanide Contraction 
The chart on the right above can be used to explain why Zirconium and Hafnium are two of the hardest elements in the Periodic Table to separate. In addition why the teaching of Transition Metal Chemistry is often covered in 2 courses since the properties of the first row elements are expected to be quite different to those of the second and third row.

Considering that the size of Gallium is smaller than Aluminium suggests that the 3d contraction is having an impact as well.

Ionic Radius

Although neither atoms nor ions have sharp boundaries, they are sometimes treated as if they were hard spheres with radii such that the sum of ionic radii of the cation and anion gives the distance between the ions in a crystal lattice.

Ions may be larger or smaller than the neutral atom, depending on the ion’s charge. When an atom loses an electron to form a cation, the lost electron no longer contributes to shielding the other electrons from the charge of the nucleus; consequently, the other electrons are more strongly attracted to the nucleus, and the radius of the atom gets smaller. Similarly, when an electron is added to an atom, forming an anion, the added electron shields the other electrons from the nucleus, with the result that the size of the atom increases. Typical values range from 50 pm to over 220 pm.

Atomic and Ionic Radius Retrieved 24 November 2014





In the Valence Bond (VB) theory an atom may rearrange its atomic orbitals prior to the bond formation. Instead of using the atomic orbitals directly, mixture of them (hybrids) are formed. For carbon (and other elements of the second row) the hybridization is limited to mixing one 2s and three 2p orbitals, as appropriate.

We recognize three basis types of hybridization: sp3, sp2 and sp. These terms specifically refer to the hybridization of the atom and indicate the number of p orbitals used to form hybrids.

  • In sp3 hybridization all three p orbitals are mixed with the s orbital to generate four new hybrids (all will form σ type bonds or hold lone electron pairs).
  • If two p orbitals are utilized in making hybrids with the s orbital, we get three new hybrid orbitals that will form σ type bonds (or hold lone electron pairs), and  the “unused” p may participate in π type bonding.  We call such an arrangement sp2 hybridization.
  • If only one p orbital is mixed with the s orbital, in sp hybridization, we produce two hybrids that will participate in σ type bonding (or hold a lone electron pair). In this case, the remaining two p orbitals may be a part of two perpendicular π systems.

The most important rule is that the number of orbitals must be preserved

in the mixing process. The mixing principles can be illustrated on a simple

example of one s and one p orbital making two equivalent sp hybrids

(only one is shown here for clarity). The constructive interference of

the wavefunctions at the “right” half gives a large lobe of the hybrid

orbital, while the destructive interference on the left (opposite signs

of the s and p wavefunctions) yields a small “tail”.  The second hybrid

formed in this case is a 180°-rotated version of the one shown.  In

this case each of the two hybrids is constructed from ½ of s and ½ of p.

Within each type of hybridization, one can produce infinite number of different hybrids (mixtures). The hybrids are defined by the p to s ratio of the contributing orbitals.  Thus, an spm hybrid is composed of m+1 parts: one part of s and m parts of p, and the p/s ratio is equal m, called the hybridization index.  For example, an sp3 hybrid has ¼ (25%) of s and ¾ (75%) of p. This fraction is called an s (or p) character of the orbital. Thus, an sp3 hybrid has 25% s character.

The hybrids with larger s character have bigger front lobes (and smaller “tails”) than the hybrids with smaller s character, as illustrated above for an sp (left) and sp3 (right) hybrids.  The s/p ratio is, thus, responsible for the bonding properties of the hybrid.  Increased s contribution brings electrons closer to the nuclei, increasing stabilizing Coulomb interactions. The more s character the hybrid orbital has the lower its energy, the better its overlap (with bonding partners), and the stronger (and shorter) bonds it can form.

It is important not to confuse the hybridization type that applies to the atom (see above) with the individual hybrid character that is described by the hybridization index m. The first (indirectly) indicates the number of p orbitals set aside (for participation in π systems), the second precisely describes the specific mixture of s and p used to construct the given hybrid.

The mixing of s and p orbitals in different ratios also results in changes of the angle between the resulting hybrids.  Since the atomic p orbitals are 90° from each other, their various degree of participation in the mixing will yield hybrids separated by different angles, depending on their p character.  For two identical hybrids, in general, the more p character in the hybrids the smaller the angle between them. Thus, two pure p orbitals (100% p) are 90° apart, two sp3 hybrids are 109.5° apart,  two sp2 hybrids are 120° apart, and two sp hybrids are 180° apart.  More generally, the angle (α) between any two hybrids (spm and spn) is given by cosα = –1/(m·n)0.5.

An atom will adjust its hybridization in such a way as to form the strongest possible bonds and keep all its bonding and lone-pair electrons in as low-energy hybrids as possible, and as far from each other as possible (to minimize electron-electron repulsions).  This adjustment  is accomplished by varying s and p characters of individual mixtures, but “moving” s character between hybrids (to lower energy of some) also changes the angles between them (potentially increasing electron-electron repulsion).  Thus, it is all a compromise game.

Let us look at some examples. Something simple to start: methane. The carbon atom forms four identical bonds using four identical hybrid orbitals. These orbitals are the result of sp3 hybridization (here we talk about the hybridization type), i.e. one s and three p orbitals are mixed to form four sp3 hybrids (here we talk about the composition, or character, of each hybrid).  Each of these hybrids is composed of ¼ of s and ¾ of p (the p/s ratio is 3, i.e. m = 3). The angle between any two such hybrid orbitals is (yes…  that’s the cosine formula above) 109.5o.  The table below gives more examples of different hybrid orbitals involved in making C-H bonds. Note that for the same hybridization type (sp3 in our example) one can have quite different hybrids involved in making C-H bonds.

Table 1. Hybrids Involved in Making C–H Bonds in Simple Hydrocarbons

Hydrocarbon Hybridization


Bond angle Hybrid  involveda s characterb Bond length (A) BDEc sp3 109.5o sp3 25% 1.100 105 sp3 107.3o sp3.36 23% 1.100 100 sp3 115o sp2.37 30% 1.089 106 46 sp2 116.6o sp2.23 31% 1.076 106 45 sp 180o sp(e) 50% 1.060 132 25

a. Hybrid orbitals on carbon involved in the formation of the indicated (in green) bonds with hydrogen atoms calculated from the formula: m = -1/cosα

b. Percent s character in the hybrid orbitals [(1/(m+1)) ×100]

c. Homolytic Bond Dissociation Energies (BDE). BDE’s depend to a much larger degree on the stability of the radicals formed than on the hybridization type of the bond broken (see below).

d. Acidity of the C–H bonds (the pKa values for methane and ethane are very approximate).

e. Strictly speaking, in this case the hybrid cannot be determined just from the bond angle. An angle between any two sp-type hybrids is always 180°, regardless of their specific s characters. In this case, the two hybrids made by mixing s and p orbitals are not equivalent, one makes bond to carbon, one to hydrogen.

When we say that bonds made out of hybrids containing more s character are stronger, we have to be very precise in our meaning. Bond strength are commonly measured by bond dissociation energies (BDE’s) that reflect enthalpies of homolysis (i.e. the energy required to break a bond, forming two radicals).  What counts in such considerations is the actual strength of the bond broken (that is related to the s character of the hybrid) and the stability of the radicals formed which is influenced strongly by other effects (such as hyperconjugation or resonance) that may not even be present in the molecule before homolysis.  In general, the stability of the radicals is a more important contribution to BDE’s than the hybridization effect (i.e. s character).

Similarly the s character of the orbital containing the lone electron pair will influence the stability of the anion, and therefore, the pKa value of the hydrocarbon precursors of this anion.  If the geometry of the hydrocarbon is similar to that of the anion, the more s character in the hybrid involved in making the “acidic” C-H bond, the lower the pKa value of the hydrocarbon (and the more stable the anion).  But, again caution is required in making comparisons:  the stability of the anion also strongly depends on other effects (such as inductive and resonance stabilization, ion pairing and solvation, i.e. interactions with solvent molecules). In particular, if the resonance stabilization is involved, the hybridization of the hydrocarbon and the anion derived from it are going to be quite different.

In general, very rarely all hybrids formed by an atom are exactly equal.  We may have an infinite number of combinations of mixtures.  For example, let us assume that carbon forms two sp2 hybrids (with an angle of 120°) to two identical substituents and additional two (equal to each other) spx hybrid orbitals that are going to be used to form two additional bonds, as shown in A below. How can we find x? It is very simple, just a little fraction arithmetic! Each of the two sp2 hybrids contains 1/3 of s (total of 2/3 s). The remaining 1/3 s must be divided between the two identical spx hybrids; i.e. 1/6 s per orbital. To make the “full” hybrid the “missing” 5/6 of the hybrid must be composed of p. So, the ratio of p/s = 5 (or x = 5) and we have our answer (sp5). If you do not believe it, you may check the math doing the balance for the p orbitals.   Here, how it goes: the two sp2 hybrids contain 2/3 of p each (total 4/3 p).  Since all three p orbitals participate in hybridization, we have 3 – 4/3 = 5/3 of p left to be used in the two spx orbitals. That leaves (5/3)/2, or 5/6 of p per hybrid; exactly the same answer we got doing the balance for the s orbital.

Now we can look at the hybridization of nitrogen in ammonia (B) or oxygen in water (C) with more precision. Since each of these central atoms has four electron pairs around, and no π bonds (that information is available from a simple Lewis structure) we may say that oxygen and nitrogen are sp3 hybridized. We mean that each atom uses three p’s and one s atomic orbitals to make four hybrids. But none of these hybrids is an sp3 hybrid! The angle between hydrogens in : NH3 is 107°  Since all the hydrogens are identical, we can calculate that the hybrids used to make N-H bonds are sp3.42 (cos(107°) = –1/m or m = 3.42); i.e they have 1/(1+3.42) = 22.6% s character each). That leaves 32.2% s character for the hybrid containing the lone pair, or the lone pair is an sp2.10 hybrid (the p character in the lone-pair hybrid must be 100 –-32.2 = 67.8%, or m/(m+1) = 0.678 from where m = 2.10). For a moment, forget the arithmetic! The point is that the lone pair hybrid has increased its s character in comparison to what it would be in the pure sp3 hybrid (32.2% vs 25% s) stabilizing the lone pair (more s character more stabilization). And the lone pair needs more stabilization: these are, by definition, unshared electrons! The price to be paid is the decrease in the H-N-H angle from the ideal tetrahedral 109.5o to 107o and the increased repulsion between the bonding pairs. The observed situation is the compromise between these two trends. You may recognize this concept as being equivalent to saying that the lone pair needs “more space” than the bonding pairs. Similar situation is found in H2O. Here, the H-O-H bond is 104.5°. Now, you do the math! The O-H bonds use sp4 hybrids and the lone pairs (two here) are sp2.3, or have about 30% s character each. Still, that is better than the 25% s of the pure sp3 hybrids.