Boron:

Boron belongs to the Group IIIB, elements that mark, the starting of the P-block elements. All such elements shows a group valency of three, but due to the very large input of energy that is essential to form the 3-valent ions - the sum of the first three ionization energies - their compounds whenever anhydrous are either necessarily covalent or comprises an appreciable amount of covalent character. Boron never forms a B³⁺ ion as the huge amount of energy needed to eliminate three electrons from a small atom can't be repaid by the formation of a stable crystal lattice, even by the most electronegative fluorine atom.

Boron takes place principally, as borates example - sodium borate in which the boron storm is part of the anionic complex. Boron can be attained as an amorphous brown powder via treating borax with hydrochloric acid, igniting the baric acid H₃BO₃, to give the oxide, B₂O₃ and lastly reducing the latter by magnesium at a high temperature.

B₂O₃(s) + 3Mg(s) → 2B(s) + 3MgO(s)

This is employed in the construction of high impact resistant steel and, as it absorbs neutrons, in reactor rods for controlling the atomic reactions.

A crystalline form of boron can be obtained via the thermal decomposition of the boron trioxide on a tantalum filament.

2BI₃(s) → 2B(s) + 3I₂(s)

Chemical Properties of Boron:

Amorphous boron is an extremely reactive element combining directly with oxygen, sulphur, nitrogen and halogens to give correspondingly an oxide, sulphide, nitride and halide by covalent bonds running entirely via the structure.

a) Halides of Boron:

The volatility of halides reduces with increasing relative molecular mass, therefore BF₃ and BCl₃ are gases, BBr₃ is a liquid and BI₃ is a white solid. They are covalent and exist as BX₃ molecules, their structures being planar.  

The halides react vigorously by water to provide the halogen hydride, with exception of boron trifluoride that provides fluoroboric acid, HBF₄, which in solution includes the tetrahedral BF₄^- ion, boric acid is as well formed, example:

BCl₃ (g) + 3H₂O (l) → H₃BO₃ (aq) + 3HCl (g)

4BF₃ (g) + 3H₂O (l) → 3BF₄ (aq) + H₃BO₃ (aq) 

Boron trifluoride is employed as a Friedel-crafts catalyst in the organic chemistry, specifically for polymerisation reactions.

b) Diborane:

This compound, which is an inflammable and extremely reactive gas, is the simplest hydride of boron. This can be generated by the reduction of boron trichloride by Lithium aluminium hydride and should be handled in the vacuum systems which make use of mercury valves, as it attacks tap-grease.

4BCl₃ (g) + 3LiAlH₄ → 2B₂H₆ (g) + 3LiCl (s) + 3AlCl₃ (s)

Physical proof exhibits that the structure can be considered to comprise somewhat resembling boron - boron double bonds, the extra two electrons needed for this being given by two of the six hydrogen atoms that bridge the boron atoms as protons.  

The other hydrides of boron are B₄H₁₀, B₃H₉, B₅H₉, B₅H₁₁ and B₁₀H₁₄, they all are electron deficient.

c) Boron Trioxide:

Boron trioxide can be obtained via burning boron in oxygen or by fusing the orthoboric acid.

2H₃BO₃ (l) → B₂O₃ (l) + 3H₂O (l)

It is generally obtained as a glassy material whose structure comprise of arbitrarily orientated three-dimensional networks of BO₃ groups, each and every oxygen atom uniting the two boron atoms.

Boron trioxide is an acid oxide and gradually reacts with water, making orthoboric acid. Whenever fused by metallic oxides it forms borate glasses which are frequently colored, this in the basis of borax based test in the qualitative analyses. 

d) Orthoboric Acid:

Orthoboric acid is made up whenever the boron halides are hydrolyzed or whenever dilute hydrochloric acid is added to the solution of borax.

B₄O₅^2-(aq) + 2H⁺ (aq) + 5H₂O (l) → 4H₃BO₃ (aq)

This is obtained as a white solid on the subsequent crystallization. The orthoboric acid is a weak monobasic acid and in aqueous solution the boron atom completes its octet via removing OH from the water molecule.

B(OH)₃(aq) + 2H₂O (l) → B(OH)₂(aq) + H₃O⁺(aq)  

It thus functions as a Lewis acid and not as the proton donor. The structure of orthoboric acid is mainly based on the planar B(OH)₃ unit.

e) Borates:

Boron, similar to silicon, has a great affinity for oxygen and a multitude of structures exist having rings of alternating boron and oxygen atoms. The single BO₃^3-ion is instead uncommon however does take place in (Mg²⁺)₃ (BO₃^3-), the ion, as expected, consists of a planar structure.

The more complex borates are mainly based on the triangular BO₃ units example: (Na⁺)₃B₃O₆^2- has the structure as illustrated below :

Fig: Borates

f) Borazine and Boron Nitride:

Whenever ammonia and Diborane, in the ratio of two molecules to one, are reacted altogether at high temperature the volatile compound Borazine is formed, the molecule consists of a cyclic structure reminiscent of the benzene

Fig: Borazine

The structure of Borazine is taken to be a resonance hybrid of the two structures.

Fig: Borazine-resonance hybrid

This is isoelectronic with benzene and looks like the latter in some of its physical and chemical properties; its boiling point is 55^oC and forms addition compounds more eagerly than benzene example: it forms addition compound by hydrogen chloride, while benzene is unreactive towards this reagent.

Boron nitride (BN) is made up by the direct union of boron and nitrogen at white heat, it consists of a structure identical to that of graphite and is therefore a giant molecule however differ from graphite in only being the semiconductor of electricity.

Carbon and Silicon:

Carbon and silicon belongs to the Group IVA elements. They exhibit most of the properties that are feature of non metals, however as we move down the group, the metallic nature of the elements rises.

The normal valency of elements is four, however apart from carbon; the elements can make more than four bonds. This is due to the reason that they make use of a set of d-orbitals in bonding that is the empty d-orbitals of the outer electron shell, example: the 3d set for silicon. The availability of the d-orbitals is responsible for the capability of silicon to make complex ions like SiF₆^2-, with the exception of carbon. The other characteristic of the chemistry of Group IV is that some carbon compounds are less reactive than the corresponding compounds of the outer members of the group.

Chemistry of Carbon and Silicon:

a) The hydrides:

Carbon and silicon provides a variety number of hydrides like CH₄, C₂H₂, S₁H₄, S₁₂H₆ and so on. The geometries of hydrides follow those of methane and are mainly dependent on a tetrahedral arrangement around the central atom.

The carbon hydrides will not ignite in air except a flame is put to them. Apart of silicon, SiH₄, the silicon hydrides are less well behaved. For illustration, S₁₃H₈ is spontaneously flammable in the air.

Si₃H₈ (l) + 5O₂ (g) → 3SiO₂ (s) + 4H₂O (l)

Similar to the carbon hydride, silicon hydrides are not hydrolyzed by water alone. Though, traces of alkali will transform them into hydrated silica, SiO₂nH₂O and hydrogen gas. Carbon hydride is not hydrolyzed via alkali.

b) The Halide:

Carbon and silicon forms different halides like CCl₄, SiCl₄, and SiF₆^2-. There is a tendency for the elements to form four bonds and by a tetrahedral arrangement. As by the hydrides there is a tetrachloromethane and silicon tetrachloride. CCl₄ will not react with water however SiCl₄ is immediately transformed into silica.

SiCl₄ (l) + 2H₂O (l) → SiO₂ (s) + 4H⁺ (aq) + 4Cl^- (aq)   

c) The Oxides:

The oxides of carbon and silicon are mainly covalent; however the chief oxide of silicon, SiO₂, dissimilar to the small gaseous molecules CO and CO₂, consists of a giant molecular structure that is better symbolized by the formula (SiO₂)n.

Carbon and silicon oxide are definitely acidic. For illustration, silica behaves similar to CO₂ whenever it reacts with hot and concentrated alkali

SiO₂ (s) + 2OH^- (aq) → SiO₃^2- (aq) + H₂O (l)

Silica will as well react by metal carbonates giving off CO₂.

SiO₂ (s) + Na₂CO₃ → Na₂SiO₃ (s) + CO₂ (g)

Silica is not destroyed via hydrogen ions as are carbonates. Instead if you add dilute acid to the sodium silicate solution, the solution will turn to a gel. This is colloidal and an illustration of hydrosed silica, SiO₂.2H₂O.

d) Organic Compound:

The carbon and silicon form numerous interesting organic compounds, made up of chains of representing units example: organo-silicon compounds.

-(Si(CH₃)₂-O)_n-

Nitrogen and Phosphorous:

As a member of Group VA elements, nitrogen and phosphorus exhibit the typical properties of non-metals. For illustration, they are poor conductors of heat and electricity and provide acidic oxides. Their compounds are mainly covalent. Phosphorus consists of allotrope. 

Chemistry of Nitrogen and Phosphorus:

a) The Hydrides:

Both of them form hydrides by the unpleasant sources. They contain pyramidal shape; however the bond angle in hydrocarbon and the hydrides of group IV differ from that of the ammonia. 

Similar to ammonia, phosphine, PH₃, can accept a proton out its lone pair and give the phosphorium ion PH₄⁺ and it will join to make phosphornium iodide, PH₄I. Similar to the analogous ammonium salts it is ionic.

Though, phosphine will not accept protons as readily as ammonia.

b) The Halides and Oxohalides:

Nitrogen and phosphorus form trihalide with F, Cl, Br and I and pentahalides like PF₅, PCl₅, PBr₅, whereas those of nitrogen don't exist. Phosphorus penta chloride fumes in the air. This reacts in water to gives oxochloride.

PCl₅ (s) + H₂O (l) → PoCl₃ (l) + 2HCl (g)

And in surplus of water it gives:

PCl₅ (s) + 5H₂O (l) → H₃PO₄ (aq) + 5HCl (g)

c) The Oxides:

Nitrogen mainly forms different oxides like N₂O, NO, N₂O₃, NO₂, N₂O₄ and N₂O₅. The oxides of phosphorus are P₄O₆ and P₄O₁₀, which was once given as P₂O₃ and P₂O₅ correspondingly, before their structure, were found via x-ray diffraction. 

d) The Oxoacids:

Phosphorous forms different oxoacids such as H₃PO₄, H₃PO₃, H₃PO₂, H₃PO and so on. The structure of phosphoric (V) acid is illustrated below:

Fig: Oxoacids

e) The Sulphides:

The sulphide of nitrogen is N₄S₄, whereas those of phosphorus are P₄S₃, P₄S₅, P₄S₇ and P₄S₁₀. Specifically P₄S₃, have been employed in making matches. To make the match head the sulphide is mixed by an oxidising agent, example: potassium chloride (v), and a little ground glass and the mixture is bound altogether by glue. The match boxes generally encompass a strip of sand paper all along the side.

Oxygen and Sulphur:

Oxygen exists as the diatomic molecules, O₂. It consists of three isotopes, ¹⁶₈O or the main one, and the others are ¹⁷₈O and ¹⁸₈O, both of around 0.3%. Oxygen as well exists in triatomic molecules as ozone, O₃ having a triangular shape.

The one chemical property which dominates the chemistry of oxygen is its capability to combine by both metals and non-metals to make oxides. Oxides can be of four kinds: neutral, basic, acidic and amphoteric.

The Group I and II metals join directly by oxygen to provide basic oxides. Particularly, Na and K have to be kept under oil in order to stop them transforming into oxides. The reactivity of Group II metal is less marked, however a coating of oxides will soon provide the or else shiny metal surfaces a dull grey appearance. The oxides of sulphur and phosphorus are typical acidic oxides in that they all react by H₂O to give acidic solutions example:

P₄O₁₀ + 4H₂O → H₃PO₄ (aq) ↔ 2PO₄^3- (aq) + 6H⁺ (aq)

The most significant peroxide is hydrogen peroxide, H₂O₂. If pure, it is a colourless liquid, however it is too dangerous to use in this form in the laboratory, rather, it is kept in solution with water.

Sulphur posses allotropes and the other main difference to oxygen come about via the use of d-orbitals in bonding. In most of the sulphur compounds the bond to sulphur are shorter than expected. This subjects a degree of double character, which can take place if sulphur makes use of its empty 3d orbitals and also its S and P orbitals. Dissimilar oxygen, sulphur and other Group VI elements can make up to six covalent bonds by utilizing d-orbitals example: SF6.

Chemistry of Sulphur and Oxygen:

The majority of sulphur is employed to make sulphuric acid. Sulphuric acid is regarded as a strong acid in water. It dissociates in two phases:

H₂SO₄ (aq) + H₂O (l) → HSO^-₄ (aq) + H₃O⁺ (aq)

HSO₄^- (aq) + H₂O ↔ SO₄^2- (aq) + H₃O⁺ (aq) 

The acid exhibits its oxidising nature when it is concentrated, for illustration:

a) The Hydrides:

Of the hydrides of Group VI, water is by far the most significant, and is not typical of the others. Water is liquid at room temperature, due to the hydrogen bonding, whereas others are gases.

Hydrogen sulphide is extremely poisonous, having a rotten egg smell. The gas can be made by mixing hydrochloric and by a metalsulphide, frequently iron (II) sulphide.

FeS (s) + 2HCl (sq) → FeCl₂ (aq) + H₂S (g)

Dissimilar to water, but similar to ammonia, hydrogen sulphide can be a good reducing agent.

b) The halides and Oxohalides:

Oxygen forms halide having all the halogens, whereas sulphur forms halides with all the halogens with exception of iodine.

Fluorine brings out the highest oxidation state as in SF₆. The fact that such atoms can in reality make six bonds is due to their use of d-orbitals in bonding. Of the oxohalides, the most significant are those of sulphur, example: thionylchloride, SOCl₂ and sulphurylchloride, SO₂Cl₂. The previous is a colourless liquid that is simply hydrolyzed.

SOCl (l) + 2H₂O (l) → H₂SO₃ (aq) + 2HCl (aq)

c) The Oxides:

SO₂ and SO₃ are gaseous in state. They are both highly soluble in water, with the reaction between SO₃ and water being explosive.    

d) Sulphides, Sulphates and other Oxoanions:

Sulphates consist of SO₃^2- ion. Most of the sulphides are soluble in water, and act as reducing agents. Whenever they are warmed with acid, SO₂ is given off.

SO₃^2- (aq) + 2H⁺ (aq) → SO₂ (g) + H₂O (l)

Sulphates consist of the SO₄^2-ion. The sulphates of Group II metals tend to be insoluble example: CaSO₄.Sulphates generally decomposes whenever heated to an adequately high temperature.

Fe(SO₄) → FeSO₃(s) + 3SO₃(g)

Thiosulphates consist of the ion, S₂O₃^2-. The structure of the ion is similar to that of a sulphate ion, apart from that one of the oxygen atoms is substituted by a sulphur atom. Sodium thiosulphate solution is broadly employed as a fixing agent in photography. It has the capability to dissolve the silver salts that have not been influenced by light. In the laboratory, thiosulphate solutions are employed in the iodine titrations.

I₂ (aq) + 2S₂O₃^2- (aq) → 2I^- (aq) + S₄O₆^2- (aq)

Peroxodisulphates encompass the ion S₂O₈^2- and are found in salts like K₂S₂O₆. They are oxidising agents and act according to the half-equation

S₂O₈^2- (aq) + 2e^- → 2SO₄^2- (aq)

They will for illustration, oxidize iodide to iodine and Iron (II) to iron (III).

e) The Sulphides:

The sulphides of Group I metals are ionic, example:  (Na⁺)₂S^2-. The sulphides of other metals, particularly the B metals are covalent to a lesser extent.

The Halogens:

Halogens are the members of Group VIIA elements and comprise: fluorine, chlorine, bromine and iodine. Fluorine and dioxide are gas; Bromine is a liquid, whereas iodine at room temperature occurs as a solid.

The Chemistry of Halogens:

Fluorine consists of the lowest energy as a result it reacts more readily than the other halogens. Reactively reduces down the group. 

a) Halogens as Oxidising Agents:

Each and every halogen encompasses a tendency to accept electrons. The only kind of negative ion they all give is the halide ion X^-.

They all readily form alkali halides having group IA metals. The alkali metals towards the bottom of Group I can react violently by fluorine and chlorine. 

They reaction readily by hydrocarbon, generating carbon

C₁₀H₂₂ (s) + 11Cl₂ (g) → 10C (s) + 22HCl (g)

b) Reactions by Water and Alkali:   

Both fluorine and chlorine are capable to oxidize the water. Fluorine can provide a mixture of oxygen and trioxygen.

2F₂ (g) + 2H₂O (1) → O₂ (g) + 4HF (aq) 

Chlorine doesn't liberate oxygen, rather, solution having a mixture of hydrochlorous and chloric (l) acid (hypochlorous acid) is formed.

Cl₂ (g) + H₂O (l) → HCl (aq) + HClO (aq)

Chlorate (l) ions, ClO^-, in a solution of chlorine are accountable for its bleaching action.

The halogen reacts by cold dilute alkali as represented below:

X₂ (g) + 2OH^- (aq) → X^- (aq) + XO^- (aq) + H₂O (l)

And if heated with concentrated alkali as:

3X₂ (g) + 6OH^- (aq) → 5X^- (aq) + XO₅^- (aq) + 3H₂O (l)

c) The Halide Ions:

Frequently, whenever a halogen reacts, each and every atom gains an electron to provide a halide ion. This is possible to differentiate between chloride, bromide and iodide ions. The simplest test comprises adding silver nitrate solution to a solution of the halide. This must be done in the presence of dilute nitric acid; or else other ions might provide precipitates. Silver ions react by halide ions to give precipitate. These in turn can be recognized by their colour, or by their reaction by ammonia solution.

Chloride ions are white. 

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