Hydrogen Element, Chemistry tutorial

Hydrogen:

Lavoister gave the name hydrogen to inflammable gas collected by reacting iron by sulphuric acid. Hydrogen takes place in the Free State, in some volcanic gases and in the outer atmosphere of the sun. Other stars are comprised almost wholly of hydrogen. The major source of hydrogen is water, and petroleum and natural gas, where it takes place in combination by carbon. The element is an important ingredient in all living matter, being found in the proteins and fats.

The hydrogen atom comprises of one proton and one electron, having electronic configuration of 1S1. Most of the chemistry of hydrogen can be illustrated in terms of its tendency to obtain the electronic configuration of the noble gas helium, this it can do by gaining an extra electron to provide the hydride ion, H-, by sharing its electron and accepting a part share of an electron from atom as in the hydrogen molecule H -H, and by accepting a lone pair of electrons, which it does as the proton whenever combined with, for illustration water and ammonia to give the hydroxonium H3O+ and ammonium NH4+ ions correspondingly.

Hydrogen can readily be attained by the action of certain metal on water or steam and dilute acids. Nitric acid and concentrated sulphuric acid should be avoided.  

Na + H2O → 2NaOH + H2 (g)

And by action of steam on coke or hydrocarbon.

C(s) + H2O → 2NaOH + CO (g) + H2 (g)

There are mainly three isotopes of hydrogen of relative masses 1, 2 and 3. They are known as ordinary hydrogen, deuterium and tritium respectively and differ in that while ordinary hydrogen has no neutrons, deuterium consists of one and tritium consists of two neutrons in the nucleus. Tritium is the mere one which is radioactive. The ratio of ordinary hydrogen to deuterium in hydrogen compounds is around 6000:1; tritium takes place in even smaller amounts.

Deuterium is somewhat less reactive than ordinary hydrogen however or else its properties are nearly similar. Deuterium is employed as a tracer for describing a broad range of reaction mechanisms, and thus are its compounds. Most of these compounds can be readily obtained from deuterium oxide (D2O-heavy water), for illustration DCl, an equivalent of hydrogen chloride.

Properties of Hydrogen:

Hydrogen is a colourless gas with no taste or smell. This can be liquefied by compression and cooling in the liquid nitrogen, followed through sudden expansion. Liquid hydrogen boils at 25oC and becomes solid at - 259oC. Hydrogen burns in air and, under some conditions, reacts explosively by oxygen and the halogens example:

2H2 (g) + O2 (g) → 2H2O (l)

This reacts partially by boiling sulphur to give hydrogen sulphide. 

H2 (g) + S (l) ↔ H2S (l)   

And by nitrogen at elevated atmosphere and pressure in the presence of a catalyst to form ammonia. This forms ionic hydrides with most of the metals of Group 1A and 2A. It form covalent hydrides with the elements from Group 4B and 7B in the periodic table and are gaseous at ordinary temperature, with exemption of a few, and with transition metals a series of instead ill-defined compounds - interstitial hydrogen- is formed.

Active Hydrogen:

The atomic hydrogen can be produced by dissociating hydrogen molecules into atoms by employing high energy sources, like discharge tube having hydrogen at low pressure, or a high current density at high temperature. Therefore, dissociation is highly endothermic 

H2 (g) → 2H (g) ΔHo (298k) = + 435.9Kjmol-1

Most of the metals are capable to catalyze the recombination of hydrogen atoms example: platinum and tungsten, which yields in liberation of the similar quantity of energy as is required to effect the dissociation. This effect is employed in the atomic hydrogen blowlamp for welding metals. Hydrogen is a powerful reducing agent example: it will decrease metallic oxides and chlorides to metals, and oxygen to hydrogen peroxide. The nascent hydrogen is hydrogen at the instant of formation. The Nascent hydrogen can reduce elements and compounds which don't readily react by normal hydrogen.

Uses of Hydrogen:

Previously, merely small quantities of hydrogen were needed as a fuel in the form of town gas and water gas, for filling balloons and in the oxy-hydrogen blowlamp for welding. In recent times, though, large quantities of the gas are employed in the procedure listed below:

a) Manufacture of the ammonia through Harber process. 

b) Manufacture of the hydrogen chloride and hydrochloric acid.

c) Manufacture of organic chemicals example-methanol.

d) Manufacture of the margarine.

e) Extraction of a few metals from their oxide.

f) Liquid hydrogen has been employed as a rocket fuel. 

Study of group IA and group IIA elements:

The elements of Group IA comprises L, Na, k, Rb, Cs and Fr, the first four are metals. Certainly, from a chemist view-point, they made an exceptional set as they encompass a large number of properties in common. Lithium is the only member of the group which is not fully typical.

They are all highly electropositive metals. Certainly, the tendency for them to lose their outermost electron and change to a positive ion is the most significant feature of their chemistry. This is due to the fact that the outer 'S' electron is extremely well shielded by the inner electrons. The 'S' electron feels merely a fraction of the nuclear charge. As we do down the group, shielding wins over the effect of the increasing numbers of protons in the nucleus. Caesium, for illustration is a much more powerful reducing agent than the sodium. The metals are so reactive that in nature they are for all time found combined by other elements. They do exist as chlorides, nitrogen, sulphates and carbonize.  

It is hard to convert Group IA metal ions to neutral atoms, so if we require obtaining the pure metal we have to make use of electrolysis. The pure metals are silvery white and apart from Li, soft and simple to cu. Though, they quickly tarnish in air giving a layer of oxide, peroxide or at times super oxide. They will as well react violently with water. For both the reasons they are kept under a layer of oil.

The elements of Group IIA: Be, Mg, Ca, Sr and Ba are all metals. They exhibit the properties we would expect, example: they are good reducing agents; they provide ionic compounds, their oxides and hydroxides are basic, and they provide hydrogen with acids. The alkaline nature of the elements is responsible for them being acknowledged as the alkaline earth metals. The exception of the common pattern is the first member, beryllium. One reason why beryllium is different is that its elements are not strongly shielded from the nucleus. The radius of Be2+ ion is very small, and it represents an extremely dense centre of positive charge by an immense polarizing power. This capability to draw electrons towards itself is responsible for the covalence of most of its compounds. The other characteristic of chemistry of beryllium is that in solution its compounds tend to suffer from hydrolysis, and a few are amphoteric instead of fully basic. Similar to the Group IA metals, the reaction of the elements makes it hard to extract them via chemical means.

Chemistry of Group IA:

a) Reaction with oxygen:

Lithium oxidizes less rapidly than the other metals; however they all offer ionic oxides and peroxides. In an abundant supply of oxygen the reactions can be violent.

2K (s) + O2 (g) → K2O2 (g)

They all are basic. They dissolve in water to give strongly alkaline solutions having hydroxide ions for illustration:

Na2O (s) + H2O (l) → NaOH (aq)

b) Reaction with Water:

Li, Na and K all float on water. Li reacts only slowly, however Na and K reacts more rapidly. Hydrogen is given off and the solution remaining is alkaline. The reactions of Rb and Cs by water must not be attempted, due to explosions.

2Na (s) + 2H2O (l) → 2HaOH (aq) + H2 (g)

c) The Hydroxides

The hydroxides of Group IA metals are among the strongest bases acknowledged.

They exist as ionic solids and are extremely soluble in water. With the exception of LiOH, which is slightly soluble, it and is as well the only one that will transform to an oxide on heating.

d) The Carbonates and Hydrogen Carbonates:

The carbonates are all soluble in water and their hydrogen carbonates exist as the solids. The exemption once again is lithium, which doesn't give a hydrogen carbonate.

Sodium carbonate is a helpful substance; it is sold as washing soda crystals Na2CO3.10H2O. In water it provides a slightly alkaline solution owing to salt hydrolysis.

Na2CO3 + H2O → NaHCO3 (aq) + NaOH (aq)

The case by which hydrogen carbonates provide all carbon-dioxide is made use of in fire extinguishers and baking powders.

e) The Halides:

All the metals provide fluorides, chlorides, bromides and iodides. Apart from caesium they encompass the similar crystal structure as the sodium chloride.

f) The Nitrogen and Nitrides:

Sodium nitrate (NaNO3), and sodium nitrites (NaNO2) are the most significant salts.  In common by all other nitrates, sodium nitrate is soluble in water. Chemically, the Group IA nitrates are a little different to those of other metals. In specific, whenever they are heated, they give off oxygen and change to a nitrite.  

2KNO3 (s) → 2KNO2 (s) + O2 (g)

Most of the nitrates are energetically stable. Though, the nitrogen in a nitrate ion is in a high oxidation state (+5) and the ions have a high percentage of oxygen. With the right chemicals, the ions will exhibit a considerable capability to act as oxidising agents. Particularly, KNO3 mixed with sulphur and carbon is employed as a gun powder. Sodium nitrite is employed in the manufacture of dyes and in increasing the shell life of raw meat sold in the supermarkets.

g) The Sulphates, Hydrogen Sulphases and Sulphites:  

All the members of the group provide sulphates and hydrogen sulphates. They all are soluble in water.

Sulphites, like sodium sulphite, Na2SO3, are more reactive than either sulphates or hydrogen sulphates example, if you warm a sulphite by an acid, you will find out sulphur dioxide is given off.

Na2SO3 (aq) + 2HCl → 2NaCl + SO2 (g) + H2O (l)

Sodium thiosulphate is generated by boiling a solution of sodium sulphate with powdered sulphur. Sodium thiosulphate is employed as a hypo in photography. In the laboratory it is employed in the iodine titrations.

h) The Hydrides:

All the hydrides of group are ionic, having the metal being positive and the hydrogen being negative.

Chemistry of Group IIA:

a) Beryllium oxide, BeO, is more similar to the oxide of aluminium in Group III instead of the oxides of the other element in Group II. It consists of a high degree of covalence, which is lacking in the other oxides. This is insoluble in water and it will dissolve only by great difficulty in acids. The reactivity of BeO based on its treatment. If it is heated to a high temperature (that is, above 800oC) it becomes nearly fully inert.

The other oxides will dissolve in water by increasing ease down the group. The resultant solutions are slightly alkaline owing to the reactions between the oxides and water example: 

MgO (s) + H2O (l) ↔ MgOH (aq)

b) The Halides:

The elements all provide fluorides, bromides and iodides and also chlorides. They all are soluble in water; however the fluorides are much less soluble than the others.

c) The Carbonates and Hydrogen Carbonates:

The Group IIA carbonates are dissimilar to those of the alkali metals of Group IA in two main respects. Primarily they are merely very slightly soluble in water, with the solubility reduces down the Group. Secondly they are decomposed via heat, giving off CO2 and leaving an oxide.

MgCO3 (s) → MgO (s) + CO2 (g)

d) The Sulphates:

The solubility of the sulphates reduces down the group. Be, Mg and Ca sulphates are often found as hydroxide crystals, example: BeSO4.4H2O, MgSO4.7H2O, CaSO4 .2H2O. The crystals of magnesium sulphates encompass a rather unfortunate reputation. They are better termed as Epsom salts (employed as laxative). The crystals of CaSO4.2H2O are found in nature as the mineral gypsum. Anhydrous calcium sulphate as well takes place naturally as anhydrite. Whenever gypsum is heated to around 100oC, it loses three quarters of its water of crystallization. The powder remaining is the plaster of paris.

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