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Discovery of Element Thorium

Thorium Discovery

In 1815 J. Berzelius, the discoverer of the element, named it thorium in honour of Thor, the ancient Scandinavian god of thunder. But the famous Swedish chemist anticipated the events: no new element was discovered by him that year. He analysed a rare mineral from Falun mines in which he discovered what he  believed to be the oxide of an unknown element. Berzelius thought this justified the addition of one more name to the list of the existing elements. No contemporary dared even to doubt the discovery since in those times the scientists had boundless trust in Berzelius. However, Berzelius himself had doubt, and justifiably so: ten years later it was shown that the oxide observed by him was yttrium phosphate (yttrium had already been known for a long time). Thus, in 1825 the past triumph turned sour.

A year later F. Wohler reported the discovery of a new element in a rare Norwegian mineral now known under the name of “pyrochlore”. Wohler did not attach particular importance to this observation and, as it turned out, mistakenly so.

Meanwhile, G. Esmark found a heavy black mineral on the Leven island near the shores of Norway. The scientist sent a sample of the mineral to J. Berzelius who thoroughly analysed it. In 1828 Berzelius reported isolation of silicates of a new element from the mineral. The old name “thorium” proved useful. The mineral which had become the source of thorium–2 was named by J. Berzelius “thorite”.

When Berzelius studied the properties of thorium, Wohler paid attention to the fact they were similar to those of the element which he left without attention in 1826. Wohler was much more disappointed when six years later the famous German scientist and traveller W. Humboldt presented him with a sample of pyrochlore from Siberia. Wohler discovered thorium in it as a few years earlier he had found it in the Norwegian pyrochlore. Thus, thorium played a trick on Wohler.

  1. Berzelius tried to separate pure thorium but in vain. For very long the element was known in the form of its oxide and only in the 1870’s was it prepared in the metallic form. Thus, thorium became the second radioactive element (after uranium) to be discovered by the conventional chemical analysis having nothing to do with radioactivity.
Will cockroaches emerge as survivors of a nuclear war ?

Will cockroaches emerge as survivors of a nuclear war ?
Sufficient evidence to back the notion that cockroaches are radiation resistant is lacking
Everyone expects cockroaches to survive a nuclear war. If cockroaches are close to the ‘epicentre’ they will be incinerated because of the intense heat. The legend is built on the notion that cockroaches are radiation resistant.
Safe testing
Recently ‘myth-busters’ of the Discovery Channel decided to settle the issue once and for all. Kari Byron, the TV hostess, revealed that the radiation resistance of cockroaches was in their lost of myths from day one.
The TV team had to convince the discovery Channel that they can do the testing safely (Tri-city Herald, October 19, 2007). Kari Byron noted that people are just scared when they hear radiation. She attributes this to the availability of “too many zombie movies”.
Routine use
It appears that the Channel does not know that industry routinely uses hundreds of irradiators all over the world safely to expose thousands of samples of food, medicine and surgical products to precisely known radiation doses. Irradiating a few hundred cockroaches is no big deal! Byron’s team will expose 200 ‘farm fresh cockroaches, bought from a scientific supply company.
Different doses
The staff of Pacific North west National Laboratory will assist the TV team to expose the cockroaches to different does of gamma radiation using an irradiator located in the basement of a building at Hanford.
A group of 50 cockroaches, left unexposed will serve as control. The second group of 50 will receive 1000 rads. The third and fourth groups of 50 will have to suffer doses of 10,000 rads and 100,000 rads respectively. (Rad is a unit of radiation dose; a dose of 450 rads may kill 50 percent of the persons exposed).
To add insult to injury, the experimenters will confirm the beasts to small boxes to ensure uniform doses to each group! They plane to expose flour beetles and fruit-files to similar doses.
The team faces some logistic problems. The exposed insects must reach San Francisco for close observation.
They cannot fly them as airlines will not let them in the passenger cabin; they cannot be placed in the baggage hold without wrecking the experiment.
Final destination
Grant Imahara, electronics and radio-control specialist of the TV Channel revealed that they have to maintain reasonable temperature and humidity so that the cockroaches will not go into shock. A ‘mythbusters’ employee will drive them to the final destination in San Francisco.
Tri-city Herald quoted Michelle Johnson, a technical group manager of the national lab as saying that the show presents good examples of scientific method and encourages developing a questioning attitude.
Gamma irradiation
“I have been told that cockroaches are more resistant to radiation and in a world nuclear war, only the cockroaches would survive. But I have not seen any publication that discusses it with any credibility…. I have irradiated cockroaches and a constructed killing curves for them using gamma irradiation… my opinion is that insects in general would be relatively resistant to radiation compared to non-insects…. “Joseph G Kunkel, Professor of Biology, at the University of Massachusetts at Amherst, who maintains a cockroach home page, asserted.
The lives of insects revolve around their molting (periodical shedding and renewal of the outer skin) cycle. During a molting cycle the cells of the insert divide usually only once.
“This is encoded in Dyar’s Rule, i.e. insert double their weight at each molt and thus cells need to divide only once per molting cycle”, he wrote.
“Cells are most sensitive to radiation when they are dividing. Now if a typical cockroach molts at most once a week, its cells usually divide within a 48-hour period within that week…about ¾ of the cockroaches would not have cells that are particularly radiation sensitive at any one time.”
“If a killing radiation is endured by a cockroach and human population, then ¾ of the cockroaches might survive while none of the humans might survive since our blood stem-cells and immune stem-cells are dividing all the time”, he clarified.

So we can say cockroaches have better probability to survive in Nuclear war than humans.

Discovery of Aluminium Element

Aluminium

Aluminium is chemical element to which history was unjust. The third most abundant metal on Earth after oxygen and silicon, and found practically everywhere in the earth’s crust (in 250 minerals, at least) aluminium was discovered only in 1825. And still, this later discovery of aluminium is not accidental. It was due to the extreme stability of aluminium oxide. To separate metallic aluminium from it is a tall order even in our times, to say nothing of the last century. Such reducing agents as charcoal and hydrogen could not separate the metal from the oxide. Only alkali metals, first of all potassium, made it possible “to capture the fortress”. This shows how the discovery of some elements created the prerequisites for the discovery of others: free aluminium was first prepared with the help of potassium.

Man knew of various aluminium compounds in very remote times. Clay and brick are nothing but usual aluminosilicates. Alumina (aluminium oxide) was a constant companion of man but many centuries were required to prove the presence of a new metal in it. Aluminium is one of the main components in such precious stones known from the time immemorial as ruby and garnet, sapphire and turquoise. Alums were known for a long time. In Latin they were named alumen–the word which contained the root of the future “aluminium”. However, the composition of alums remained undetermined for a long time and they were often confused with other compounds.

In 1754 the German chemist Marggraf tried to shed light on the problem. Having added pure alkali to the alum solution, he obtained a dense white precipitate which he named “alum earth”. Then Marggraf observed that the addition of sulphuric acid to the “earth” yielded alum; thus, the composition of alum was established. And, finally, Marggraf demonstrated the presence of the “alum earth” in clays. Had history so willed it, Marggraf would have been acclaimed as the discoverer of this element, but history waited for somebody else to prepare pure aluminium. Only 30 years after Marggraf’s experiment did it become clear that alumina was an oxide of an unknown element. This was suggested by A. Lavoisier who placed “alum earth” into his “Table of Simple Bodies”. But no attempts were made for some time to separate the element in a free state.

The first attempt was made by H. Davy and J. Berzelius, who tried to decompose alumina with the aid of electric current, but in vain; it was only H. Davy’s proposal (1807) to name the element “aluminium” that had any practical importance. This name became internationally accepted although in Russia the name “glinium” (from the Russian word for “clay”) was used for a long time.

The first who managed to obtain metallic aluminium was the Danish scientist H. Oersted known in history as a physicist rather than as a chemist. He discovered the induction of magnetic field of an electric current, but preparation of pure aluminium showed him to be also a skillful chemist. Having red–heated a mixture of alumina with charcoal, Oersted passed chlorine through it; as a result anhydrous aluminium chloride was obtained. Then the scientist heated the new compound with potassium amalgam and obtained amalgam of aluminium for the first time. As soon as Oersted distilled off the mercury, he discovered pieces of metal that looked like tin. The product contained impurities but nevertheless, this was the birth of metallic aluminium. Oersted published an article in a little known Danish journal which passed practically unnoticed in the scientific circles. And news of Oersted’s achievement did not reach many chemists. Therefore, some historians believe that aluminium was discovered not by Oersted but F. Wohler.

The second discovery of aluminium took place two years later, in 1827. Undoubtedly, F. Wohler was a more skillful experimenter than Oersted and his process of separating pure aluminium was more sophisticated. At first Wohlar’s attempt to obtain the metal using the Danish scientist’s method failed but soon he succeeded in preparing small amounts of anhydrous aluminium chloride. Wohler developed his own procedure for the process: (1) preparation of aluminium hydroxide; (2) preparation of a thick paste from aluminium hydroxide; charcoal and vegetable oil; (3) calcination of the paste and preparation of a mixture of aluminium with charcoal powder; (4) preparation of pure anhydrous AlCl3 by passing dry chlorine through the mixture. The complexity of this procedure was rewarded by the purity of the product. The scientist decomposed AlCl3 with potassium under conditions ensuring the highest possible purity of the metal. F. Wohler was the first chemist to describe the most important properties of metallic aluminium and in 1845 he prepared aluminium in the form of an ingot.

However, Wohler, like his predecessors, did not obtain pure aluminium. The decisive word was said by the French chemist A. Saint Claire Deville. In 1854 he prepared the samples of pure metal, using sodium instead of potassium for the reduction stage. Simultaneously with Bunsen he performed electrolysis of melted double chloride of aluminium and sodium: this was the first instance of producing aluminium electrochemically. A. Saint Claire Deville also pioneered the development of an industrial process of aluminium production.

It is difficult to believe that only one hundred years ago this silvery metal was extremely expensive and was even called “clay silver”. Things made of aluminium cost no less than gold ones. Only after the processes for producing cheap electric energy had been developed and rich deposits of aluminium ore had been found, did aluminium become a metal for everyday uses.

Silicon

Silicon

Silicon is the second most abundant element on Earth after of oxygen. Although is constitutes 28 per cent of the earth’s crust, its abundance did not make for its early discovery. The reason for this lies in the difficulty of reducing silicon from its oxide. 

Generally speaking, there is every ground to classify silicon as an element of antiquity. Its compounds were known and used from time immemorial (suffice it to mention silicon tools of primitive man). We classified carbon as an element of antiquity since it was known in a free state from very remote times. However, that carbon is a chemical element became clear only two hundred years ago. Glass, in the long run, is also a silicon material. However, the date of silicon discovery is the date of its preparation in a free state since such is the established practice in the history of science.

At the turn of the 18th century many scientists believed that silica, or silica earth, contained an unknown chemical element and tried to isolate it in a free state. H. Davy at attempted to decompose silica with an electric current–the method by which a number of alkali metals had already been prepared–but without success. The scientist’s attempt to prepare free silicon by passing metallic potassium vapour over red–hot silicon oxide also failed. In 1811 L. J. Gay Lussac and L. Thenard applied themselves to the problem. They observed a vigorous reaction between silicon tetrafluoride and metallic potassium; a reddish brown compound was formed in the reaction. The scientists could not reveal the nature of the product; most likely, it was contaminated amorphous silicon.

At last, in 1823, J. Berzelius had stroke of good luck. The Swedish chemist heated a ground mixture of silicon oxide, iron and charcoal to a high temperature and obtained an alloy of silicon and iron (ferrosilicium), the composition of which he was able to prove. To separate free silicon, J. Berzelius repeated L. Thenard and L. J. Gay Lussac’s experiments and also obtained a brown mass. Under the action of water, hydrogen bubbles were liberated and free amorphous silicon was formed as a dark brown insoluble powder which contained potassium silicofluoride as an impurity. Berzelius removed the impurity by washing the precipitate for a very long time.

Another method proposed by J. Berzelius–calcination of potassium fluorosilicate with an excess of potassium–proved to be more successful and straightforward. The sintered mass was decomposed with water and, as a result, pure amorphous silicon was obtained. J. Berzelius showed that upon calcination silicon was transformed into silica; this makes Berzelius the discoverer of silicon. Crystalline silicon was obtained in 1854 by A. Saint Claire Deville during separation of metallic aluminium (see p. 109). The Latin name “silicium” originates from “silex” meaning “a hard stone”.

Selenium

Selenium

Selenium is still another element that chemists had met long before its discovery, but failed to identify owing to its having been masked by the presence of other similar elements. Thus, selenium remained undiscovered, “hiding” behind Sulphur and tellurium. Only in 1817 did it surrender to the Swedish chemists–the famous J. Berzelius and his assistant G. Gahn. Inspecting a sulphuric acid factory in Gripsholm on September 23, they found a small amount of a precipitate, partially red and partially light brown, in sulphuric acid. On heating in the flame of a blowpipe, the precipitate emitted a weak smell of radish and transformed into a regulus with a leaden lustre. In Klaproth’s opinion the smell of radish pointed to the presence of tellurium. Similar smell was noticed in the Falun mine where pyrite required for the acid production was extracted. Curiosity and hope to find this rare metal in the brown precipitate forced Berzelius to investigate it. However, he did not discover tellurium. Then he collected the deposits formed after several months of Sulphur combustion for sulphuric acid production in the Falun factory and obtained a large amount of precipitate. Thoroughly analysing the precipitate, Berzelius came to the conclusion that it contained an unknown metal whose properties were similar to those of tellurium. By analogy, the new metal was named “selenium” from the Greek selenusfor the Moon (as tellurium is named after our planet). Berzelius studied many properties of selenium and described them in an article “The Study of a New Mineral Body Found in Sulphur Extracted in Falun” published in 1818 in the journal Annales de chimie et de physique.

Discovery of Elements : Lithium

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.

Concept of HARD AND SOFT ACIDS AND BASES

HARD AND SOFT ACIDS AND BASES

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.

(b)      

  • [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:

\[SO_{3}^{2-}+HF\xrightarrow{{}}HSO_{3}^{-}+{{F}^{-}}

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 Cadmium element

Cadmium

In 1817,  F. Stromeyer,a lecturer of the Chair of Chemistry at Göttingen University (Medical Department) and the chief inspector of chemist’s shops in Hanover, found that calcination of zinc carbonate, sold on chemist’s shops, produced a yellow compound although neither iron nor lead impurities were discovered in it.

This remarkable fact interested Stromeyer and he decided to visit a pharmaceutical firm in Salzgitter where he observed the same phenomenon. This prompted the scientist to study zinc oxide in more detail. To his surprise, Stromeyer discovered that the colour which zinc oxide acquired was due to a strange metal oxide never observed before. The chemist succeeded in separating this oxide from zinc oxide and reducing it to the metal.

His method consisted in the following: he dissolved contaminated zinc oxide in sulphuric acid and passed hydrogen sulphide through the solution; then he filtered off and washed the mixture of sulphides and dissolved it in concentrated hydrochloric acid. The acid was removed by evaporating the solution to dryness. Having dissolved the residue in water, F. Stromeyer added a large amount of ammonium carbonate. Since carbonate of the new metal did not dissolve in the presence of ammonium carbonate, Stromeyer filtered the precipitate off, washed it, and transformed it into oxide which he reduced to metal with charcoal upon heating. As a result, bluish grey metal was obtained. However, since Stromeyer had only three grams of this metal, he could not thoroughly study its properties. Only in 1818 did he succeed in investigating the new metal.

F. Stromeyer named the metal “cadmia”, in accordance with the method of its preparation (as a result of calcination of ZnCO3). “Cadmia” is the Greek for natural ZnCO3. Independently of F. Stromeyer but somewhat later cadmium was discovered by W. Maissner and K. Kersten in Germany (1818). Stromeyer’s priority was contested by the German physician K. Roloff who, by the way, was the first to pay attention to the strange behaviour of commercially available zinc oxide upon heating. K. Kerston suggested to name the new metal “melinum” because of the yellow colour of its sulphide. It was also proposed to name the new metal “klaprothium” (in honour of M. Klaproth) or “unonium” (after the asteroid) but none of the names found acceptance.

Discovery of Boron

 Discovery of Boron

People widely used borax, one of the boron compounds, back in the Middle Ages. Probably borax had been known much earlier; it was reported that in the first millenium A. D. borax was used for soldering metals. However, the composition of natural borax remained unclear for a long time. Boric acid was obtained for the first time in 1702 by the Dutch physician W. Homberg who heated borax with sulphuric acid. It was used in medicine as “Homberg’s sedative salt”. In 1747 the French chemist Th. Baron tried to determine the composition of borax. He found that it contained Homberg’s salt and soda; he was quite right; now we know that borax is sodium salt of boric acid (Na2B4O7).

The name of Swedish chemist T. Bergman deserves mention in the early history of Boron. He believed that Homberg’s salt was most likely not a salt but a compound resembling acid. As a matter of fact it was he who introduced the name “boric acid”. The term “boric radical” was mentioned in Lavoisier’s “Table of Simple Bodies” and meant boron oxide. However, twenty years had to pass before the new chemical element, boron, was discovered.

It so happened that boron was discovered by several scientists: the French chemists L. Thenard and L. J. Gay Lussac and the English chemist H. Davy. They named the new element “boron” and “boracium” (from the word “borax”). The method of preparing the new element was the same in both cases: reduction of Boric acid with metallic Potassium. Independent discovery of a new chemical element by several researchers within ten days was unique event in the history of chemical elements. Gay Lussac and Thenard announced their discovery on June 21, 1808, and Davy on June 30. Clearly, the priority of the French chemists in this case was ephemeral, especially because of the fact that it was Davy’s previous discovery (preparation of elemental potassium) that gave the means for the separation of free boron.

Effect of Temperature over rate of a Chemical Reaction

Factors affecting Reaction rate

The rate of a chemical reaction depends on the rate of encounter between the molecules of the reactants which in turn depends on the following things.

(1)     Effect of temperature on reaction rate : The rate of chemical reaction generally increases on increasing the temperature.

(2)     Nature of reactants : (i) Reactions involving polar and ionic substances including the proton transfer reactions are usually very fast. On the other hand, the reaction in which bonds is rearranged, or electrons transferred are slow.

(ii)    Oxidation-reduction reactions, which involve transfer of electrons, are also slow as compared to the ionic substance.

(iii)   Substitution reactions are relatively much slower.

(3)     pH of the medium : The rate of a reaction taking place in aqueous solution often depends upon the  ion concentration. Some reactions become fast on increasing the H+ ion concentration while some become slow.

(4)     Concentration of reactants : The rate of a chemical reaction is directly proportional to the concentration of the reactants means rate of reaction decreases with decrease in concentration.

(5)     Surface area of reactant : Larger the surface area of reactant, the probability of collisions on the surface of the reactant particles by the surrounding  molecules increases and thus rate of reaction increases.

(6)     Presence of catalyst : The function of a catalyst is to lower down  the activation energy. The greater the decrease in the activation energy caused by the catalyst, higher will be the reaction rate. In the presence of a catalyst, the reaction follows a path of lower activation energy. Under this condition, a large number of reacting molecules are able to cross over the energy barrier and thus the rate of reaction increases.  Fig. shows how the activation energy is lowered in presence of a catalyst.

(7)       Effect of sunlight :  There are many chemical reactions whose rate are influenced by radiations particularly by ultraviolet and visible light. Such reactions are called photochemical reactions. For example, Photosynthesis, Photography, Blue printing, Photochemical synthesis of compounds etc.

H2 + Cl2  \underrightarrow { \quad Sunlight\quad (hv)\quad } 2HI : The radiant energy initiates the chemical reaction by supplying the necessary activation energy required for the reaction.

 

Rate law, Law of mass action and Rate constant

(1)     Rate law : The actual relationship between the concentration of reacting species and the reaction rate is determined experimentally and is given by the expression called rate law.

For any hypothetical reaction, aA + bB → cC + dD

Rate law expression may be, rate = k[A]a[B]b

Where a and b are constant numbers or the powers of the concentrations of the reactants  and  respectively on which the rate of reaction depends.

(i)    Rate of chemical reaction is directly proportional to the concentration of the reactants.

(ii)     The rate law represents the experimentally observed rate of reaction, which depends upon the slowest step of the reaction.

(iii)    Rate law cannot be deduced from the relationship for a given equation. It can be found by experiment only.

(iv)    It may not depend upon the concentration of species which do not appear in the equation for the over all reaction.

(2)       Law of mass action : (Guldberg and Wage 1864) This law relates rate of reaction with active mass or molar concentration of reactants. According to this law, “At a given temperature, the rate of a reaction at a particular instant is proportional to the product of the reactants at that instant raised to powers which are numerically equal to the numbers of their respective molecules in the stoichiometric equation describing the reactions.”

Active mass = Molar concentration of the substance

=  \frac { Number\quad of\quad gram\quad moles\quad of\quad the\quad substance }{ Volume\quad in\quad litres } =\frac { W/m }{ V } =\frac { n }{ V }

Where W = mass of the substance, m is the molecular mass in grams, ‘n’ is the number of g moles and V is volume in litre.

Consider the following general reaction,

m1A1 + m2A2 + m3A3 → Products

Rate of reaction  [A1]m1[A]m2[A3]m3

(3)     Rate constant : Consider a simple reaction, A → B. If C4 is the molar concentration of active mass of A at a particular instant, then,  \frac { dx }{ dt } ∝ CA  or   \frac { dx }{ dt } = kCA ; Where k is a proportionality constant, called velocity constant or rate constant or specific reaction rate constant.

At a fixed temperature, if CA = 1, then Rate =  \frac { dx }{ dt } = k

“Rate of a reaction at unit concentration of reactants is called rate constant.”

(i)      The value of rate constant depends on, Nature of reactant, Temperature and Catalyst

(It is independent of concentration of the reactants)

(ii)     Unit of rate constant :  Unit of rate constant = \left[ \frac { litre }{ mol } \right] ^{ 1-n } × sec-1  or   \left[ \frac { mol }{ litre } \right] ^{ 1-n } × sec-1

Where n = order of reaction

Difference between Rate law and Law of mass action
Rate law Law of mass action
It is an experimentally observed law. It is a theoretical law.
It depends on the concentration terms on which the rate of reaction actually depends It is based upon the stoichiometry of the equation
Example for the reaction, aA + bB → Products Example for the reaction, aA + bB → Products
Rate = k[A]m[B]n Rate = k[A]a[B]b

 

Difference between Rate of reaction and Rate constant

Rate of reaction Rate constant
It is the speed with which reactants are converted into products. It is proportionality constant.
It is measured as the rate of decrease of the concentration of reactants or the rate of increase of concentration of products with time. It is equal to rate of reaction when the concentration of each of the reactants is unity.
It depends upon the initial concentration of the reactants. It is independent of the initial concentration of the reactants. It has a constant value at fixed temperature.

 

Discovery of Halogens Elements

Halogens

Man did not become properly acquainted with halogens until the 19th century although fluorine and chlorine were discovered in the seventies of the 18th century. But the fact that chlorine is a chemical element was understood only about forty years after its  discovery. Fluorine was “hiding” behind fluorine compounds for a whole century before it was, at last, obtained in a free state. But iodine and bromine were at once recognized as simple substances.

As we see, the fates of these elements, named halogens in 1811, were different in the history of science but they played a peculiar role, especially in chemistry.

All of them were produced by chemical analysis except ree fluorine which was prepared electrochemically.

 

Fluorine

The famous Soviet scientist A.E. Fersman called this chemical element “omnivorous”. And indeed, there are very few substances, both natural and man–made, that can withstand unprecedented chemical aggressiveness of fluorine. The story of fluorine is an illustration of this property. Fluorine proved to be the last (chronologically) non–metal to be separated in a free state (apart from inert gases). One hundred years passed from the time of the forecasting of the existence of fluorine to the moment when scientists succeeded in obtaining it in a gaseous state. Chemists tried to prepare it over fifteen times but every time the attempts failed. And in several cases they even lost their lives.

At the same time a common natural compound of fluorine (fluorspar of fluorite, CaF2) had been known from very remote times. This harmless mineral known to any stone collector was mentioned in manuscripts as early as the 16th century. But when hydrofluoric acid was first prepared, fluorite assumed new significance. It is difficult to establish who was the first to prepare hydrofluoric acid; all that is known is that in 1670 the Nurnberg craftsman H. Schwanhard observed its corrosive action on glass. Schwanhard and many after him erroneously believed that etching of glass was caused by silicic acid, while it was hydrofluoric acid that destroyed glass.

A century passed before fluorspar fell into the hands of C. Scheele. He studied two varieties of fluorite–green and white. The scientist heated powdered samples with sulphuric acid and noticed that the inner surface of the glass retort became opaque while a white mass precipitated on the bottom of the retort. Scheele assumed that fluorite consisted of lime earth saturated with unknown acid. He added lime water to this acid and obtained artificial fluorspar similar to the natural mineral.

The year when hydrofluoric acid was separated (1771) is considered to be the date of the discovery of fluorine although this is hardly justified. The nature of the acid obtained by Scheele (named “Swedish acid” at the time) remained unclear. There was a controversy in the scientific world about Scheele’s discovery but with every year it became increasingly evident that he was right.

Hydrofluoric acid entered the category of reliably classified chemical compounds and scientists gradually came to believe that it contained a new chemical element. This opinion was strengthened by A. Lavoisier who included the radical of hydrofluoric acid (radical fluorique) as a simple body into “The Table of Simple Bodies” .But Lavoisier was also wrong: he thought that the acid contained oxygen. His mistake was, however, understandable since at that time chemists believed that oxygen was an indispensable constituent of all acids.

The purity of the acid prepared by Scheele’s method left much to be desired. Not before 1809 did Gay Lussac and Thenard obtain a relatively pure hydrofluoric acid, heating fluorspar with sulphuric acid in a lead retort. Both scientists were severely poisoned during the experiments.

A year later an event of extreme importance took place in the pre–history of fluorine. Two scientists–the Englishman H. Davy and the Frenchman A. Ampere–independently “banished” oxygen from hydrofluoric acid. They strongly believed that the acid was a compound of hydrogen with an unknown element and that it is similar to hydrochloric acid HCl. It was the second decisive intervention of H. Davy in the fate of halogens (shortly before he had established the elemental nature of chlorine).

It is therefore clear why Davy was the first who attempted to obtain free fluorine. By the way, the name was proposed by Ampere who borrowed it from the Greek ftoros for “destructive”. Ampere chose this name because of the hydrofluoric acid’s aggressiveness (chemists were still to see the fury of free fluorine!). But Davy was in a more peaceable mood and suggested the name of “fluorine” by analogy with “chlorine”.

Having named the element, Davy, nevertheless, did not succeed in preparing free fluorine. For two years (1813 and 1814) the scientist was storming the impregnable fortress. Two methods were used by H. Davy: the electrochemical method, which had already given the world sodium, potassium, calcium, and magnesium, and the reactions of chlorine with fluorides. Electrolysis of hydrofluoric acid gave no results; the second method was also fruitless. Severe illness caused by work with fluorine–containing compounds forced Davy to stop the experiments although he was one the first to determine the atomic mass of fluorine (19.06). Davy’s unsuccessful experiments and his illness seemed to serve as a warning for other scientists and for almost 20 years nobody tried to obtain free fluorine. Only M. Faraday, Davy’s famous pupil and assistant, whose contribution to science was no less important than that of his teacher, made an attempt in 1834 (after Davy’s death) to solve the riddle of free fluorine. However, even electrolysis of dry melted fluorides proved to be futile.

The chain of failures grew longer. In 1836 the brothers Knox from Ireland set out to solve the problem. During five years they were performing dangerous experiments, without success. The brothers were severely poisoned in the process and R. Knox died. In 1846 the Belgian P. Layette and then the French chemist D. Niklesse shared the dramatic fate of the Knox brothers. At last, in 1854–1856, E. Fremy, Professor of Ecole Polythechnique in Paris, seemed to succeed in preparing free fluorine. He electrolytically decomposedanhydrous melted CaF2. Metallic calcium deposited on the cathode, while on the anode a gas was liberated which could be nothing but fluorine. However, to observe a chain of bubbles is not enough–they had to be collected; in this, however, Framy failed. But, in our opinion, E. Fremy deserves the name of a co–discoverer of fluorine, at any rate, his right to it is no less than that of Scheele.

In 1869 the English chemist G. Gore obtained a small amount of free fluorine which at once reacted explosively with hydrogen. There were about ten other researchers who hoped to obtain free fluorine. History, of course, has their names but we shall not mention them here.

And at last the moment came when A. Moissan took resolutely in his hands the fate of fluorine. First of all, he analysed the errors of his predecessors and clearly realized that the attempts of Faraday, E. Fremy, and G. Gore had failed because they could not subdue the “fury” of fluorine which instantly reacted with the material of the apparatus. Moissan was also aware of the mistake of those investigators who tried to isolate fluorine by the action of chlorine on fluorides; chlorine had to be a weaker oxidizer than fluorine.

Moissan overcame the difficulty by using a U–shaped vessel. At first he used a platinum vessel but later decided that a copper one must be much more suitable since neither fluorine nor hydrogen fluoride reacted with copper fluoride being formed. Thus, a layer of copper fluoride prevented the vessel from destruction. Moissan filled the vessel with anhydrous hydrofluoric acid and added a small amount of potassium bifluoride to it for the solution to become electroconductive. The vessel was immersed in a cooling mixture at –25oC. Platinum electrodes were inserted through CaF2 plugs. Electrolysis liberated hydrogen on the cathode and fluorine on the anode; fluorine was collected in copper tubes.

On June 26, 1886, Moissan performed the first successful experiment, observing the flame produced by the reaction of fluorine with silicon. He sent a modest report to the Paris Academy of Sciences where he wrote that different hypotheses about the nature of the liberated gas were possible. The simplest of them was that fluorine is actually liberated, although the gas might also be hydrogen perfluoride or even a mixture of HF and ozone. The reactivity of this mixture is high enough to explain the strong action of the gas on crystalline silicic acid.

Since Moissan was not a member of the Academy, his report was read by A. Debray and a special committee was organized consisting of A. Debray, E. Fremy and “the Elder” of the French chemists M. Berthelot. On the first day Moissan’s attempt to prepare free fluorine failed but on the following day he succeeded and the committee witnessed his success. Thus, another date appeared in the biography of fluorine and, maybe, the most important one–the date of its preparation in a free state (1886). In 1887 Moissan obtained liquid fluorine.

 

Chlorine

In ancient times man knew of such chlorine–containing compounds as sodium chloride NaCl and ammonium chloride NH4Cl. Later hydrochloric acid (HCl) became known and widely used. Numerous chlorine compounds were subjected to the scrutiny of researchers and there is no doubt that during manipulations with them free chlorine was repeatedly obtained. Among those who observed free chlorine were such outstanding scientists as J. Glauber (of the Glauber’s salt fame), J. Van Helmont and R. Boyle. But even if this strange yellow–green gas had caught their attention, they would have hardly understood its nature.

The Swedish chemist C. Scheele was also mistaken. He prepared chlorine by the same method that is described in modern school textbook: by the reaction of hydrochloric acid with manganese oxide (Scheele made use of ground pyrolusite ,that is natural MnO2). It would be wrong to say that the scientist chose this method by chance. Scheele knew that the reaction of HCl with pyrolusite had to give rise as usual (see p. 47) to inflammable air (known subsequently as hydrogen). Some gas was, indeed, liberated but it did not bore even remote likeness to inflammable air. It had a very unpleasant smell and an unpleasant yellow–green colour. The gas corroded corks and bleached flowers and plant leaves. The new gas proved to be a highly active chemical reagent. It reacted with many metals and, when with ammonia, formed a dense smoke (ammonium chloride NH4Cl). Its solubility in water was poor. Scheele did not utter the words “a new chemical element”, although he had the discovery within his grasp and could follow the logical chain of arguments about its elementary nature. A zealous follower of the phlogistic theory, the Swedish chemist identified the gas discovered by him with hydrochloric acid that had lost phlogiston. He named it “dephlogisticated hydrochloric acid or dephlogisticated muric acid” (HCl was named muric acid after the Latin muria, “brine, Salt water”). At that time Scheele shared the opinion of H. Cavendish and other scientists that inflammable air (hydrogen) was actually phlogiston. It followed that the new gas had to be a simple substance (hydrochloric acid minus phlogiston) but Scheele did not make such seemingly obvious conclusion. Although 1774 is considered to be the new gas’s date of discovery, much time was to pass before its nature was properly understood.

  1. Lavoisier overturned the phlogistic theory. Even the name “dephlogisticated” muric acid” evoked a strong protest in him. In his opinion, the acid obtained by Scheele was a compound of muric (hydrochloric) acid and oxygen. Oxidized muric acid–that is how Lavoisier named what we know as elemental chlorine now. The French chemist believed that all acids must contain oxygen combined with some element. Lavoisier called this element “murium” in the case of muric acid and included it into his “Table of Simple Bodies” (murium radical –radical muriatique).

The result was paradoxical; trying to elucidate the nature of the gas discovered by Scheele, Lavoisier only complicated the issue. Probably, this development in the history of chlorine was simply inevitable in the light of new theoretical conceptions. Some chemists attempted to prepare free murium but the attempts were fruitless and the nature of the new gas did not become clearer.

In 1807 H. Davy tried to solve the problem, subjecting the notorious muric acid to various manipulations. He attempted to decompose it electrolytically, but no decomposition was observed. No matter how ingeniously he treated oxymuric acid, he could not succeed in preparing water of liberating oxygen. In a word, the acid behaved as if it were a simple substance. Moreover, its action on metals or their oxides yielded typical salts. Nothing else was left to Davy but to recognize that oxymuric acid consisted of only one simple substance, i.e. to recognize the elemental nature of the gas discovered more than 30 years earlier by Scheele. He reported on this to the Royal Society on November 19, 1810.

Davy proposed to name the element “chlorine” from the Greek chloros meaning “yellow–green”. Two years later, in 1812, the French chemist Gay Lussac proposed to change the name for “chlor” (which became generally accepted except in English–speaking countries).

Gay Lussac in cooperation with Thenard began to study oxymuric acid almost simultaneously with Davy; at first, they wanted to prove that it was oxygen–free. The two scientists passed the acid through a red–hot porcelain tube over charcoal. If there had been oxygen in the gas discovered by Scheele, it would have been absorbed by the charcoal. Although the composition of the gas at the inlet and outlet of the tube remained unchanged, this experiment did not shake the belief of the firm followers of A. Lavoisier about the composition of oxymuric acid.

Nevertheless, Davy’s experiments strongly impressed the contemporary scientific community which gradually came to the conclusion that murium was in fact chlorine. In 1813 Gay Lussac and Thenard agreed with Davy. Only Berzelius for a long time continued to doubt the elemental nature of chlorine but in the end he also had to accept the truth. The elemental nature of chlorine became an irrefutable fact only after the discovery and study of iodine and bromine.

In 1811 the German chemist I. Schweiger proposed to name chlorine a “halogen” (from the Greek for “salt” and “produce”, i.e. “salt–producing”) because of its ability to combine readily with alkaline metals. At the time the name was not accepted but later it became common for the group of similar elements: fluorine, chlorine, bromine, and iodine. Chlorine was obtained for the first time in a liquid form in 1823 by M. Faraday.

 

Iodine

Iodine was the second halogen to be obtained in a free state. Both the appearance and chemical properties of iodine are rather peculiar. Were it the only halogen in existence, chemists would have to think hard about its nature, but the elemental chlorine had already been known and this fact helped to understand the nature of iodine.

  1. Courtois, an entrepreneur from the French town of Dijon, was engaged, among other things, in the production of potash and saltpeter. He used ash of sea algae as the initial raw material. A mother solution of sea algae was formed under the action of water on the ash. To–day we know that the ash contains chlorides, bromides, iodides, carbonates, and sulphates of some alkali and alkaline–earth metals. However, when Courtois performed his experiments it was only known that the ash contained potassium and sodium compounds (chlorides, carbonates, and sulphates). Upon evaporation, first, sodium chloride precipitated and then potassium chloride and sulphate. The residual mother solution contained a complex mixture of various salts, including Sulphur–containing ones.

To decompose these Sulphur compounds, Courtois added sulphuric acid to the solution. One day it so happened that he added a greater amount of acid than was necessary. Suddenly something unexpected happened: amazingly beautiful clouds of violet vapour appeared whose magnificence was marred only by their unpleasant, even lachrymose smell. Then followed something even more surprising: on the surface of cold objects the vapour did not condense forming heavy drops of a violet liquid but precipitated at once as dark crystals with metallic lustre. Courtois discovered many other interesting and unusual properties of the new substance. He had every reason to announce the discovery of a new chemical element but, evidently, the researcher was not confident enough and his laboratory was too poorly equipped to perform further investigations. He, therefore, turned for help to his friends, Ch. Desormes and N. Clement, asking them for a permission to continue his experiments in their laboratory. He also asked them to report his discovery in a scientific journal.

Consequently, the report about “The Discovery of a New Substance Obtained from an Alkali Salt by Mr. Courtois” signed by N. Clement and Ch. Desormes appeared only in 1813 in the “Annales de chimie et de physique”, i.e. two years after the discovery of the element. To enable other chemists to investigate the substance, B. Courtois gave a very small amount of it to a pharmaceutical firm in Dijon. Clement himself prepared a certain amount of iodine, studied its properties and was, probably the first to advance an opinion that iodine, resembled chlorine. In 1813 J. Gay Lussac and H. Davy independently of each other proved the elemental nature of iodine. The French chemist suggested the name “iode” for the new element (from the Greek iodes meaning “violet colour”) and the English scientist suggested the name “iodine”. The first name found acceptance in the Russian language.

Iodine is a rare example of a chemical element whose properties were studied thoroughly during a short period of time after its discovery. Here a great contribution was made by Gay Lussac who even wrote a book on iodine which was in effect the first monograph in the history of science completely devoted to one element.

But the subsequent generations did not forget B. Courtois’s contribution. A street in Dijon is name after him; this honour was bestowed on very few discoverers of chemical elements.

 

Bromine

This element, unusual in many respects, was the last of the natural halogens to be discovered (if, of course, we accept the discovery of fluorine by Scheele in 1771).

On an autumn day in 1825, the following event took place in the laboratory of L. Gmelin, a professor of medicine and chemistry at Heidelberg University. A student by the name of C. Löwig brought to his teacher a thick–walled flask with an evil–smelling reddish brown liquid. Löwig told Gmelin that in his native town of Kreiznach he had studied the composition of water from a mineral spring. Gaseous chlorine turned the mother solution red. Löwig extracted with ether the substance that caused the colouring of the solution. It was reddish brown liquid known subsequently as bromine.

Gmelin showed great interest in his student’s work and advised him to prepare the new substance in greater amounts and to study its properties in detail. It was a reasonable piece of advice since Löwig had little experience as an experimenter; but the work required time and the time factor turned against the student.

 

Discovery of element : Ruthenium

Ruthenium

Ruthenium was the first chemical element discovered by a Russian scientist. It was Karl Klaus. The discovery of this last of the platinum metals was made forty years after the discovery of iridium.

In 1828 G.V. Ozann, Professor of the Tartu University, studied the residue obtained after the dissolution of crude Uralian platinum in aqua regia and found that it contained three new elements: pluranium, polonium and ruthenium. But Berzelius, to whom Ozann had sent a letter about his findings, did not support the discovery. Because of this significant fact the study of this platinum residue was not renewed until 1841. Berzelius’s prestige was so high that no chemist in the world would argue with him.

The second reason for such a late discovery of ruthenium is its great similarity to the other “brothers” in the family. Prior to Klaus in Russia, this problem was studied by the Polish scientist A. Snyadetskii who also reported the discovery of a new element which he named “West” after the asteroid of the same name. But his discovery proved to be false.

Klaus began his research in 1840. The then Minister of Finance of Russia E. F. Kankrin, a competent and energetic person, rendered him great assistance; Klaus obtained 2 pounds of crude platinum residue and extracted a considerable amount of iridium, rhodium, osmium and palladium from it, apart from 10% platinum. In addition, Klaus separated a mixture of metals which, in his opinion, had to contain a new substance.

First of all, Klaus repeated Ozann’s experiments. Then he continued the investigation according to his own plan. The results were striking. In 1844 he published a 188–page report with the following information: analytical results on the residue obtained after platinum dissolution in aqua regia; new methods of platinum metals separation; methods of studying lean residues; the discovery of a new metal–ruthenium; analytical results on lean residues; and the simple methods of separating platinum ores and residues; new properties and compounds of the previously known metals of the platinum group. This was a real encyclopaedia on chemistry of platinum metals.

1. Klaus separated six grams of the new element from its double salt with potassium. He sent a report about it to Berzelius but the letter was sceptical again. Great courage was required from Klaus to contradict the old and eminent scientist. The Russian chemist proved the genuineness of his discovery and in 1845 J. Berzelius recognized the new element. A special committee was formed in Russia consisting of Academicians H. Hess and Yu. F. Fritsshe to check the results obtained by Klaus. The committee confirmed the discovery and K. Klaus was awarded the Demidov’s prize (1000 roubles).

The name of the element is derived from the Latin for Russia (Ruthenia). Klaus gave this name to the element moved by his patriotic feelings and trying to show that all work in this field had been done in Russia (G. Ozann, A. Snyadetskii, K. Klaus).

Klaus spent a total of 20 years studying platinum metals. He deserves the right to be called the founder of the Russian school of studies of platinum and platinum metals.

Discovery of element : Osmium and Iridium

Osmium and Iridium

The discovery of four new elements with similar properties in one country (England) in the course of two years was unprecedented in the history of science. Another English chemist, S. Tennant, was studying platinum metals simultaneously with W. Wollaston, who discovered palladium and rhodium while extraction of osmium and iridium is associated with the names of other scientists, although the greatest contribution was made by S. Tennant.

As compared with other platinum metals, osmium and iridium have some specific features to which they owe their names. “Osmium” derives from the Greek osme for “smell” since osmium oxide is volatile and has a peculiar smell. Iridium got its name from the variety of colouring of its salts (from the Greek iris for “rainbow”). A painter could have prepared an entire palette from iridium paints if they were not so expensive. These unusual properties promoted the discovery of these platinum metals.

1. Tennant, like W. Wollaston, dissolved crude platinum in aqua regia. At the bottom of the retort he discovered a black precipitate with metallic lustre. This phenomenon had been observed previously in experiments with platinum, but the precipitate was believed to be graphite. In summer 1803 Tennant suggested that the precipitate most likely contained a new metal. In autumn of the same year the French chemist H. Collet–Descoties also concluded that the precipitate contained a metal that precipitated from ammonium platinum salts and yielded red colour. In his turn, L. Vauquelin heated the black powder with alkali and obtained a volatile oxide. Vauquelin believed that it was an oxide of the metal mentioned by H. Descoties. Tennant’s experiment set off a series of investigations. Tennant himself continued his research and in spring 1804 he reported to the British Royal Society that the powder contained two new metals which could be separated fairly easily. In 1805 he published the article “On Two Metals Found in the Black Powder Formed after Dissolution of Platinum”. The names “osmium” and “iridium” were mentioned in the article for the first time.

The notorious black powder was, evidently, a natural alloy of osmium with iridium, the so–called osmiridium. Iridium is known to be chemically stable and in the compact form does not dissolve even in aqua regia. On the contrary, osmium is readily soluble in aqua redia; in general among platinum metals osmium has the most atypical chemical properties. That is why iridium and osmium were relatively quickly and easily separated.

In 1817 the English chemist and mineralogist W. Brande justly noted in his lecture devoted to the discovery of platinum metals that if one tried to analyse the entire development of chemistry from the standpoint of contemporary analytical accuracy, the history of the discovery and separation of platinum metals would, probably, be the most striking one.

But had all of platinum metals been discovered? The question was posed again and again. Years passed but they brought nothing new, at any rate, no reliable answer. Only in 1844 was ruthenium, the last of the platinum metals, discovered; ruthenium is as abundant in nature as platinum, which, with its greatest atomic mass, was the first to be discovered. Why it was so remains a mystery. It may have been pure chance since the study of platinum metals was extremely difficult and required great analytical skill and profound knowledge of chemistry.

Discovery of element : Palladium and Rhodium

Palladium

Back in the late 17th century Brazilian miners frequently ran into a strange naturally occurring alloy. It had different names and was believed to contain gold and silver. It could be an alloy of palladium and gold. But the real discovery of the second of the platinum metals took place in 1803 owing to the work of the English chemist W. Wollaston. Studying crude (unpurified) platinum, he dissolved it in aqua regia, removed the excess of the acid, and added mercury cyanide to the solution. A yellow precipitate was formed. Heating the solution with Sulphur and borax, he obtained bright metal balls. Wollaston named the new metal “palladium” (after the asteroid discovered a year earlier by the astronomer W. Olbers). Wollaston’s success was largely owing to the fact that he had found a proper precipitating agent for palladium, mercury cyanide, which does not precipitate other platinum metals.

The discovery of palladium received publicity in a rather peculiar way. In 1804 the young Irish chemist R. chenevix put an advertisement in the Journal of Chemical Education about a “new metal for sale” which was an alloy of platinum with mercury. W. Wollaston, naturally, was of a different opinion and defended his discovery. He published the article “On a New Metal Found in Crude Platinum” in which he underlined that the metal “for sale” named palladium is contained in platinum ores although in small amounts.

Contemporary scientists (and L.  Vauquelin among them) valued highly the Wollaston’s achievement, and the more so since soon he discovered another platinum metal, rhodium. The fact that palladium was the first platinum metal to be extracted may be explained by its greatest abundance as compared with other platinum metals. In addition, it exists in nature in a native state as proved by Wollaston in 1809 and by A. Humboldt in 1825 (for Brazilian platinum ores which had been the only source of material prior to the discovery of Uralian Platinum).

 

Rhodium

The discovery of palladium became the key to the discovery of rhodium at the turn of 1803, i.e. before the news about palladium was widely spread. 

Crude platinum from South America was also a source of rhodium. It is, however not known whether it was the same sample in which Wollaston had discovered palladium. Having dissolved a certain amount of crude platinum in aqua regia and neutralized the excess of the acid with alkali, Wollaston first added an ammonium salt to precipitate platinum as ammonium chloroplatinate. Mercury cyanide was added to the remaining solution (here the experience in separating palladium proved useful) and palladium cyanide precipitated. Then Wollaston removed the excess of mercury cyanide from the solution and evaporated it to dryness; a beautiful dark–red precipitate was formed which, in the scientist’s opinion, was double chloride of sodium and of the new metal.

This salt decomposed readily upon heating in a hydrogen flow, as a result of which metallic powder was formed (after removal of sodium chloride). The scientist also obtained the new metal in the form of pellets. The name “rhodium” was given to the new element because of the red colour of its first salt to be produced (the Greek rodon means “a rose”). This element is the least abundant of the platinum metals. The only rhodium mineral known is rhodite, found in gold–bearing sands of Brazil and Colombia, whereas several minerals are known for each of the other platinum metals.

Platinum Metals : It’s Discovery of Element

Platinum Metals

The history of platinum metals (ruthenium, rhodium, palladium, osmium, iridium, and platinum) is full of false discoveries of chemical elements made in the studies of these metals and due to the great difficulties involved in studying natural ores containing platinum and accompanying metals. Platinum that mankind had come to know prior to the real discovery of this element contained different impurities. Among the platinum metals platinum occupies the second place after palladium in terms of abundance. The content of platinum metals in the minerals may vary considerably from deposit to deposit. Therefore, there were many chance events in the history of platinum and its analogues and much is still unclear. The date of the discovery of platinum is rather vague. For a long time it was not clear how many platinum metals really exist. In many cases the confusion arose because of similar properties of the platinum metals. Four of them–palladium, rhodium, osmium, and iridium– were discovered in the early 19th century owing to the considerable progress in chemical analysis. However, it is quite possible that is was just a chance that prevented earlier discovery of platinum metals, at any rate of the sufficiently abundant palladium.

 

Platinum

Platinum was the first to be discovered among the platinum metals. 1748 is considered to be its date of birth. But is it the real date? 

Ancient Greeks and Romans mentioned “electrum” an alloy that some scientists identify with platinum. Others believe that “electrum” was the Egyptian alloy of gold with silver. Pliny the Elder described a white heavy compound found in the sands of Galicia and Portugal but it was, most likely, a tin ore. A box made of platinum was found in the tomb of Queen Shapenapit (the 7th century B.C.).

In 1557 the Italian scientist G. Scaliger described a new white metal discovered in South America. It was the first definite mention of platinum. Another two centuries passed. The Paris Academy of Sciences sent an expedition to the Spanish colonies. Among its participants was a young lieutenant Don Antonio de Ulloa. Having safely returned home, he wrote the book Historical Report about the Trip to South America which was published in Madrid in 1748. He wrote that in the region of Choko he had seen many gold–bearing mines but some of them had been abandoned because of a high content of platinum in the ore. A. Ulloa was the first to note that this metal had an extremely high melting point and that it was very difficult to extract it from the ores. Two years later the English chemists W. Watson and W. Brownrigg set out to study the new metal and gave the first scientific description of it. In November 1750, W. Watson reported the discovery of a new sem–metaln called “platino–del–pinto” which had hitherto been unknown to mineralogists.

This work prompted further study of the new metal. In 1752 the Swiss chemist H. Scheffer published a detailed report about his investigation of platinum or white gold. After that a series of similar papers appeared. Two of them were particularly interesting. In 1772 C. von Sickingen extensively studied the properties of platinum, looking into the possibility of alloying platinum with silver and gold, its solubility in aqua regia, and, what is most important, he was the first to use the method of precipitating platinum from solutions with ammonium chloride. This reaction played a great role of studying the platinum metals. But the results obtained by C. von Sickingen were not published until 1782.

The second round of studies is associated with the name of P. Chabanean. He was the first to pay attention to the fact that experiments with platinum from different deposits yielded contradictory results. With hindsight this has a very single explanation: Chabanean was working not with pure platinum but with a mixture of six elements–the platinum metals that had not yet been discovered. For instance, in the absence of osmium, platinum was non–volatile and did not ignite whereas the presence of osmium made the alloy volatile and combustible.

What is the exact date of platinum’s discovery? The metal had to go a long way before it was given the right to its own title. 1750 seems to be a major landmark in the history of platinum: in that year it was studied and described in detail.

Niobium and Tantalum

Niobium and Tantalum

The early histories of these elements are so intertwined that it is hardly worthwhile to consider them separately. Their common history begins on November 26, 1801, when Ch. Hatchett made a report to a session of the Royal Society about the discovery of a new element. Communications of this type had long ceased to be a sensation. Hatchett’s report “Analysis of a Mineral from North America Containing an Unknown Metal” was received quietly. True, Hatchett got his sample not from the New World but from a place much nearer–from the British Museum. The Museum’s catalogue described the mineral as “a black ore sent to the Museum by Wintrop from Massachusetts”.

At first Ch. Hatchett assumed that the object of his study was a variety of Siberian chromium ore and tried to isolate chromic acid from it. But things took a different turn. Now it is known that the mineral from Massachusetts contained a variety of metals and it was not easy to extract a new element from it. There was no chromium in the mineral and Hatchett concluded that the compound which he had separated was not chromic acid but an oxide of an unknown metal. In honour of its place of origin, the English scientist named the mineral “columbite” (from Columbus and Columbia, the former name of America). The element was named “columbium”. A year later, in 1802, an event took place which added a little zest to the trivial discovery of columbium. In December 1802 the Swedish chemist A. Ekeberg, who analysed some minerals found near the village of Itterbul, described the discovery of an oxide of a new metal. The white oxide mass did not dissolve even in a great excess of strong acids.

The futility of all attempts to dissolve the oxide prompted Ekeberg to name the new metal “tantalum” after the “torments of Tantalus” which means useless and futile work. The mineral was named “tantalite”. A. Ekeberg was firmly convinced that he had discovered a new element and this conviction was shared by many scientists. The more surprising were the results of the English chemist V. Wollaston who announced in 1809 that he found no difference between columbium and tantalum and that the two were one and the same element. Oxides of these metals had similar densities and seemed to Wollaston to be rather similar in their chemical properties. His article was titled “On the Identity of Columbium and Tantalum”. This meant that A. Ekeberg only rediscovered columbium, confirming the discovery made by Ch. Hatchett.

Berzelius held a different opinion. He supported the name “tantalum” given to the new element by Ekeberg and believed that the names of the English and Swedish chemists must stand together in history. In autumn 1814, Berzelius wrote in a private letter to the Scottish chemist Th. Thomson (the first advocate of Dalton’s atomic theory) that he by no means wanted to belittle Hatchett’s achievement but deemed it his duty to note that the properties of tantalum and its oxide had been almost unknown before Ekeberg’s work. Berzelius thought that the columbic acid of Hatchett was a mixture of tantalum oxide and tungstic acid, but soon it became clear that there was no tungsten in columbite.

Three decades later one of Berzelius’ pupils, H. Rose, resolved the dispute once and for all. He proved that tantalum and columbite were not identical; hence, Hatchett and Ekeberg had discovered two different elements.

Rose analysed columbites and tantalites from different deposits. And every time he found that, along with tantalum, they contained another element whose properties were close to those of tantalum. Rose named the “stranger” “niobium” (Niobe was tantalus’ daughter). In the summer of 1845, the scientist studied the same mineral in which Hatchett once detected columbium and isolated niobium oxide from it, which proved to be similar to columbium oxide.

At last the confusion was cleared. It had arisen because niobium and tantalum have very similar properties and are always present together both in columbites and tantalites. As a matter of fact, Hatchett and Ekeberg discovered both elements simultaneously and could not detect any difference between them. In the mineral studied by Hatchett niobium (columbium) undoubtedly predominated. Therefore, the most important event in the biography of both elements was the development of a method for separating niobium and tantalum. This was done in 1865 by the Swiss chemist J.C. Galissard de Marignac who found the difference in solubilities of potassium fluotantalate and fluoniobate in hydrofluoric acid. In the same year de Marignac correctly determined the atomic masses of niobium and tantalum for the first time. Many chemists tried to obtain them in a pure state but, as a rule, wound up with contaminated metals. Not before the beginning of the 20th century did W. von Bolten of the USA obtain niobium and tantalum of higher than 99 per cent purity.

Discovery of element : Beryllium

Beryllium

Academician A.E. Fersman, the outstanding Soviet geochemist, called beryllium one of the most remarkable elements having tremendous theoretical and practical importance. However, beryllium is not outstanding in any one of its qualities; it has typical properties of metals. What is really remarkable, is the extremely fortunate combination (as if purposely invented by nature) of different properties. Beryllium clearly illustrates how the history of a chemical element is affected by its properties. As regards its chemical behaviour, beryllium has much more in common with aluminium (its diagonally neighbouring element in the periodic table) than with magnesium, its direct analogue in the same group. That is why aluminium was masking the presence of beryllium (as well as of zirconium) in natural minerals for such a long time.

Because of a pronounced amphoteric nature of  beryllium, all attempts to obtain beryllium compounds in a sufficiently pure form were unsuccessful for a long time. As a result, many properties of the element and especially its valence and atomic mass were determined incorrectly. Consequently, the place of beryllium in the periodic table was not definitely found for a very long time. Only after it had been firmly established that beryllium is bivalent, that the formula of its oxide is BeO, and atomic mass is 9.01, was it once and for all placed in the upper box of the second group. A great contribution to that was made by the Russian scientist I.V.Avdeev.

The history of beryllium minerals goes far back into the past when such precious stones as beryls and emeralds were already known.

One of the first scientists to begin the study of beryls in 1779 was F. Achard, Professor of Chemistry at the Berlin Academy of Sciences. Before that time he had become famous for developing an industrial method of making sugar from sugar beet. The German chemist performed six analyses. His results recalculated in modern terms show that beryls contain 21.7% silicon oxide, 60.05% aluminium oxide, 5.02% iron oxide, and 8.3% calcium oxide. The total was only 95.07% (five per cent was missing!) but F. Achard had no comment on this. Similar figures obtained in 1785 by J. Blindheim: in his case the “calculations” yielded the sum of the components of 101 per cent. So, nothing particular was found about beryls.

In 1797 M. Klaproth, who by that time had already discovered uranium, titanium, and zirconium proving himself an outstanding analyst, received from the Russian diplomat and author D. Golitsyn samples of Peruvian emeralds and analysed them. But M. Klaproth did not wind up with 100 per cent either (66.25% silica, 31.25% alumina, 0.5% iron oxide, total 98%). The scientists did not know where 2 per cent had disappeared and did not try to explain. So he was not fated to add the discovery of the fourth element to his record.

At the same time, in French, another analyst L. Vauquelin, no less skillful than M. Klaproth, was at work. Beginning with 1793 he continued to study berlys and emeralds. But Vauquelin found nothing expect ordinary components (silica, alumina, lime, iron, oxide).  Later Vauquelin recalled how difficult it had been to recognize a new substance when its properties were so similar to those of already known ones. The scientist meant a close similarity between oxides of aluminium and unknown beryllium.

Anticipating the events a little, we shall call Vauqueline the real discoverer of beryllium. The logic of discovery was not simple and it, undoubtedly, does justice to the scientist. He reasoned in the following way: beryl and emerald are very much alike as regards their composition and the shape of crystals. The crystal shape is absolutely the same but what about composition? Vauquelin’s predecessors found the same components (alumina, silica, lime) in both minerals but their content varied.

After the first unsuccessful experiments L. Vauquelin decided to see why the components content varied so widely. Could it be that the minerals contained “something” else which was either lost in the course of the reaction or, figuratively speaking, was “hiding behind the backs” of one of the components (for instance, alumina).

  • Vauquelin had a certain psychological advantage. In 1797 he discovered chromium, which imparts a greenish colour to emerald and is absent in beryl. Hence, the difference between beryl and emerald is an established fact. But not only chromium could be responsible for the difference. February 14, 1798, should be considered as the birthday of beryllium. On that day Vauquelin made a report to the Paris Academy of Sciences, “About Aquamarine, or Beryl, and the Discovery of a New Earth in This Mineral”. He told the audience how he had performed five analyses and how he had become more and more convinced of the existence of the new earth. The first results were as follows:

Beryl: 69 parts of silica, 21 parts of alumina, 8-9 parts of lime, and 1  parts of iron oxide.

Emerald: 64 parts of silica, 29 parts of alumina, 2 parts of lime, 3-4 parts of chromium oxide, and 1-2 parts of water.

Whether it was intuition or something else, but Vauquelin suspected that in both case alumina contained an impurity. It resembled alumina very much and, therefore, it was rather difficult to detect it. The brilliant intuition of an analyst helped the scientist to discover that the impurity (the new earth), unlike alumina, did not form alum. Later he found other differences. But similarity prevailed over difference enabling beryllium to hide for so long behind aluminium. If beryllium earth is not alumina, L. Vauquelin thought it is the known earths since it differs from them much more than alumina. The scientist proposed to name the new element “glucinium” (symbol Gl) from the Greek glykys which means “sweet”. The present name “beryllium” was proposed by M. Klaproth who justly noted that some compounds of other elements are also sweet.

As an interesting historical detail we should like to mention that Vauquelin analysed Altaian beryls presented to him by French mineralogist and traveller E. Patren. The discovery of L. Vauquelin was confirmed by I. Gmelin, the German chemist, a professor of chemistry in Göttingen. He analysed Siberian beryls from Nerchinsk and made the same conclusions as Vauquelin. Metallic beryllium was isolated in 1828 by F. Wohlar and E. Bussy who treated beryllium chloride with potassium metal. It was thirty years after the discovery of beryllium.

Discovery of element : Chromium

 

Chromium

Siberia may be said to be the birthplace of chromium as we shall see later; in the 18th century the mineral crocoite, known at the time as red lead ore, was found there. Some other chromium ores had been known much earlier. And this is not surprising since chromium is one of abundant elements (0.02 per cent of the total mass of the earth’s crust). But it is not easy to separate chromium even in the form of oxide and for the time being this task was beyond the power of chemists. Although chromium compounds have different colours, this peculiar fact did not attract the attention of scientists to chromium minerals.

The only exception was crocoite. For the first time it was analysed in 1766 by the German chemist I. Lehmann who lived at the time in St. Petersburg. Treating the mineral with hydrochloric acid the chemist obtained a beautiful emerald–coloured solution. But his conclusion was erroneous: crocoite contained lead contaminated with impurities. These impurities could only be chromium since crocoite is lead chromate PbCrO4. I. Lehmann was not destined to establish the composition of the mineral.

For the second time crocoite became the object of study in 1770 when P. S. Pallas, a St. Petersburg Academician, was describing the Berezov mines in the Urals: “This lead ore comes in different colours but more often looks like cinnabar. The crystals of this heavy mineral shaped as irregular pyramids are imbedded in quartz like little rubies”.

  1. S. Pallas was a traveller, geographer and mineralogist, and not a chemist. But it was he who introduced crocoite to laboratories in Western Europe. A sample of the mineral fell into the hands of the well–known chemist L. Vauquelin. Three decades passed since I. Lehmann had studied crocoite. During this time the scientists repeatedly tried to determine its composition but failed to find any new elements in it. The results obtained were very contradictory. For instance, there was an analyst who reported that lead ore contained molybdic acid, nickel, cobalt, iron and copper. In his first experiments L. Vauquelin also made mistakes and found lead dioxide, iron and alumina in crocoite.

In 1797 the French chemist decided to study crocoite more thoroughly. Step by step Vauquelin refuted the results of all the previous analyses and at last drew a conclusion that crocoite contained a new metal with properties quite different from those of other metals.

  1. Vauquelin boiled powdered crocoite with potassium carbonate. The product was lead carbonate and a yellow solution which contained, in the scientist’s opinion, a potassium salt of an unknown acid. The solution acquired bright and diverse colours when various reagents were added: mercuric salts yielded a red sediment, lead salts gave a yellow sediment, tin chloride turned the solution green all these results convinced Vauquelin that he was dealing with a new element. Its separation in the form of oxide was rather simple after that.

Many years later D.I. Mendeleev wrote in his Principles of Chemistry that the Uralian red chromium ore, or chromium–lead salt, had given Vauquelin the means to discover chromium. Vauquelin derived this name from the Greek chroma meaning “colour” because of the bright colouring of its compounds. For the sake of justice we should note that the name “chromium” for the new element was proposed by Vauquelin’s compatriots A. Fourcroy and R. Haüy. Independently of Vauquelin and almost simultaneously with him the presence of a new metal in crocoite was established by M. Klaproth who, however, did not prove it as clearly as his French colleague. Numerous attempts to obtain pure chromium were unsuccessful. L. Vauquelin himself tried to prepare it but most likely it was chromium carbide that he obtained.

 

Discovery of element : Titanium

Titanium

W. Gregor was not a chemist. But sometimes this English clergyman did chemical experiments since his hobby was mineralogy. From time to time W. Gregor studied the composition of various minerals and so succeeded in the work that afterwards J. Berzelius respected him as a prominent mineralogist. One day W. Gregor became interested in the composition of black sand whose deposit he found in the Menaccin valley on the territory of his parish. This black sand, resembling very much gunpowder, mixed with dingy–white sand of a different kind attracted W. Gregor’s attention. Having separated specks of black sand, he analysed them; you will judge the carefulness of this analysis from the following figures:

per cent (these 9/16 is especially impressive) is iron oxides; 3  per cent is silica, and 45 per cent is accounted for by the compound described by Gregor as reddish–brown lime. And 4  per cent was lost during the analysis. In this list it is the reddish–brown lime that is of interest. It dissolved in sulphuric acid yielding a yellow solution. Under the action of zinc, tin, or iron, the solution turned purple. Gregor wrote an article, reporting his findings. Being very modest, he believed that his investigation was incomplete. He only set forth some facts the explanation of which was the privilege of more knowledgeable scientists.

His friend, mineralogist D. Hawkins, convinced Gregor that the black sand was a new unknown mineral. Such an opinion from a man who knew about mineralogy not less than Gregor, suggested to the latter that the black sand contained a new metallic substance. Gregor proposed to name it “menaccin” in honour of the place where the sand had been found, and the sand itself menaccite (or menacconite). Now this black sand is named ilmenite and has the formula FeTiO3. All this goes to show that titanium was discovered in 1791 by W. Gregor.

But many historians of science believe that M. Klaproth was the discoverer of titanium although the merit of Gregor’s work is unquestionable. But the English clergyman was too unambitious. Klaproth chose another way. Of course, he read Gregor’s report but did not immediately grasp its meaning. In 1795 Klaproth succeeded in separating an oxide of the new element from the mineral brought from Hungary. Now this mineral is known as rutile (TiO2). The oxide separated by Klaproth and the menaccin earth found by Gregor turned out to be very much alike. Soon Klaproth established that he and Gregor had discovered the same element.

The German scientist named the element “titanium” from mythological “Titans”– the sons of Ge (the goddess of Earth). Pure metallic titanium was obtained only in 1910.

Discovery of elements : Uranium

 

Uranium

There is hardly another chemical element that from near oblivion sprang to instant fame. This is uranium occupying box No. 92 in the periodic table. Discovered in 1789, uranium did not interest chemists for a long time and even its atomic mass was determined incorrectly. Its practical use was confined to making coloured glass. But in 1906 in the eighth edition of Principle of Chemistry Mendeleev appealed to those who were searching for new subjects of investigation to pay close attention to uranium compounds. The reason Mendeleev gave was that two most important events at the end of the 19th century science were related to uranium: the discovery of helium and the discovery of radio activity. And, finally is it a mere chance that uranium is the last in the series of naturally occurring chemical elements, the heaviest of them?

Some scientists have referred to the ninety second element as element No. 1 of our century.

And yet, there was nothing extraordinary about the  discovery of uranium about two hundred years ago. It was like many others during the emergence of analytical chemistry. There is no doubt about the name of the discoverer–M. Klaproth. True, the actual extraction of uranium is associated with the name of another scientist (we shall come back to it later).

Pitchblende had been known to man for ages. When chemical analysis was still in its infancy, pitchblende was considered to be an ore of zinc and iron. More accurate knowledge of its composition was to come later.

When a pitchblende sample fell into the hands of Klaproth, he dissolved a piece of the mineral in nitric acid and added potash to the solution. Yellow precipitate was formed which was soluble in the excess of potash. The precipitate was small greenish–yellow crystals in the form of hexagonal prisms. Gradually the scientist made the conclusion that he had obtained a salt of a new element. Having prepared an oxide, the scientist tried to separate pure metal. And when a lustrous black powder was formed on the bottom of the crucible, the German scientist decided that the aim was attained. But Klaproth was mistaken. At the most he obtained a mixture of oxide with a small amount of the metal. Indeed, chemists were yet to see how difficult it is to extract pure uranium.

Confident of success, M. Klaproth proposed the name “uranium” for the element discovered. The chemist wrote: “In old times only seven planets were known and thought to correspond to seven metals, and according to this tradition the new metal should rightfully be named after the planet which has been recently discovered”. It was the planet Uranus discovered in 1781 by the English astronomer Herschel. After that it became fashionable to name newly discovered chemical elements after celestial bodies. Uranium had been included in the list of simple substances and made its way to chemical textbooks, but metallic uranium remained unobtainable for a long time to come. There were even scientists who were doubtful about the discovery of the German chemist. Six years after Klaproth’s death (1817), J. Arfvedson, the pupil of Berzelius, decided, perhaps, following his teacher’s advice, to remove these doubts. He tried to reduce dark–green uranium oxide with hydrogen. Arfvedson believed that the initial material was the lower oxide (we know now that the Swedish scientist worked with U3O8). The reaction yielded a brown powder (UO2). J. Arfvedson, however, thought that he extracted metallic uranium.

It was only in 1841 that the French chemist E. Peligot succeeded with the aid of a new reduction method. He heated anhydrous uranium chloride mixed with metallic potassium in a closed platinum crucible and obtained a black metallic powder. Its properties noticeably differed from those which M. Klaproth used to ascribe to metallic uranium. Therefore, some historians of science associate the real discovery of uranium with name of E. Peligot.

Ingots of the metal were produced by the French chemist A. Moissan who melted it in an electric furnace invented by him in which a very high temperature could be attained. The scientist produced the first ingot in May. 1896, and gave it to Bacquerel. With the aid of the sample A. Bacquerel established that radioactivity is a property of the elemental uranium. This property attracted everybody’s attention to uranium for the first time.

At one time uranium gave a lot of trouble to D.I. Mendeleev when the scientist was working on his periodic table. The atomic mass of uranium was considered to be 120 and, therefore, uranium was placed in the third group as a heavy analogue of aluminium. But this allocation by no means agreed with the properties of uranium. Mandeleev concluded that the atomic weight had been determined incorrectly and proposed to increase it by 100 per cent. This put uranium in Group VI under tungsten and made it the last element in the periodic table.