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Module 1

Module 2

Module 3

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Final Exam

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handbook

Periodic table

s-elements

p-elements

d-elements

Fotogallery

Video

N [He]2s22p3 P [Ne] 3s23p3 As [Ar] 4s24p3 Sb [Kr] 5s25p3 Bi [Xe] 6s26p3

Halogens

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Properties of Halogens

Chemical

Physical

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History Of Discovery

Electrode potentials

Oxidation States

Compounds of Halogens. Chemical and Physical Properties Overview

N

P

As

Sb

Bi

-3

NH3

PH3

AsH3

SbH3

BiH3

NH4+(aq)

PH4+(aq)

Li3N

BN

-2

N2H4

P2H4

-1

NH2OH

0

N2

P(red),P(white), P(black)

As(grey), As (black)

Sb(yellow)

Bi

+1

N2O

H3PO2

+2

NO

+3

N2O3

P4O6 H3PO3

As2O3 H3AsO3

Sb2O3

Bi2O3 Bi(OH)3

+4

NO2 N2O4

+5

N2O5

P4O10

As2O5

Sb2O5

Bi2O5

HNO3

HPO3

H3AsO4

H3SbO3

HBiO3

NO3-(aq)

H4P2O7

HAsO3

K[Sb(OH)6]

H3PO4

POCl3

Electron configurations of nitrogen subgroup

Nitrogen

Nitrogen with four valence orbitals can be in the states of sp³-, sp²- and sp-hybridization like other second period elements|. There is a possibility of one or two -bonds (together with - bonds) formation in case of sp²- and sp- states,|formation| respectively. |figure,state,camp,mill|

The formation of nitrogen molecule is accompanied with release of significant amount of heat:

2N = N₂ + 940 kJ/mol

The molecule of N₂ is so extraordinarity stable that a lot of energy is required for its destruction. Additionally, internuclear distance is very short (0.1094 nm|) since the triple NN bond forms in the molecule. |curiously|The bond energy (940 kJ/mol) is equivalent to six single N—N bonds|truss| (160 kJ/mol), while|where as| the energy of triple СС bond |in the similar compound (С₂Н₂, 830 kJ/mol) is equivalent only to 2.5 single C—C bonds|truss| |. Whereas, a single N—N bond is considerably weaker (160 kJ/mol) than a single С—С bond (345 kJ/mol). Such a decrease in the strength of single homoatomic| bond for these elements is explained by the presence of antibonding| electrons in the atom of nitrogen increasing|aggravates| interatomic| repulsion. Therefore, chains —N—N—N— are so unstable unlike carbon chains.

The strongest in N₂ molecule is the first bond (510 kJ/mol) that prevents|prevent| destruction of triple NN bond (the |changing|order of bond strength change in the series —С—С—, —С=С— and —СС— is|truss| contrary|return|). The energy diagram|chart| of N₂ molecule is shown below:

2N^o = 2[1s²2s²2p³] = N₂[KK(^bonding2s)²(*2s)²(^bonding2p)⁴(^bonding2p)²]

The energy difference between s- and р-orbitals| is increased across the period towards |b group VIII elements. In accordance with this rule energies of 2s- and 2р-orbitals are nearly identical but electrons on 2s- and 2р- orbitals| have mutual repulsion, therefore ^bonding2р_х and ^bonding2р_y orbitals| become|stands| more favorable in energy and they are filled first,|earlier| not ^bonding2р_z. Out of ten electrons of N₂ molecule eight are located on bonding|coupling| orbitals|, and two - on antibonding ones. Six bonding|coupling| electrons are in excess|overabundance|, forming triple bond in N₂.

Triple covalent| bond in N₂ is absolutely|quite| unique. There are no other examples of diatomic homoatomic| molecules where bond of similar|like| type could be realized and could become so strong. From the point of view of the Valence Bond method | the structure|building| of N₂ molecule can be represented as follows:

Thus,|on this grow| six unpaired| electrons of 2р-sublevel of both atoms of nitrogen form three electronic dublets| and consequently three identical covalent| bonds|truss|.

The last energy level electronic configuration of this element is 2s²2p³. The other elements have vacant d-orbitals and excitement of one of the paired electrons of s-orbital is possible:

N     P, As, Sb, Bi    

This means that maximal covalence of nitrogen (the number of covalent bonds) is four (for instance, there are three -bonds and single delocalised -bond in NO₃^-ion; the bond order of each N—O bond is less than two and more than one). The oxidation states of nitrogen in compounds are|halving,compound,junction,joint,coupling|: -3, -2, -1, 0 +1, +2, +3, +4, +5.

Phosphorus. The outermost electron shell configuration of the ground state of a phosphorus atom is 3s²3p³. As shown above, the availability of vacant d-orbitals changes properties of P significantly although phosphorus is an electronic equivalent of nitrogen (both have identical outermost shells). Phosphorus displays -3, -2 (P₂H₄), 0, +1, +3, +5 oxidation states. Note that even oxidation states of phosphorus are not typical unlike nitrogen.

Electronegativity (EN) of phosphorus is 2.06. The lower EN, the less typical is the formation of ionic compounds with metals.

Arsenic, Antimony, Bismuth. These elements have the last energy level with 5 electrons - ns²np³, i.e. two paired ns- electrons and three unpaired np-electrons. However, there are also differences in electronic structure between them. Unlike phosphorus, they contain completed (n-1)d-sublevel, and in case of Bi, which is situated after lanthanides, 4f¹⁴-sublevel appears.

Forming a chemical bond with three unpaired electrons and depending on the EN of a partner, these elements may display oxidation states -3 and +3. The formation of compounds with the highest oxidation state +5 is attributed to the excited state of the elements. Like other subgroups of p-elements, it becomes less stable with the growth of the period number, i.e. in the series As—Sb—Bi.

Moreover, the secondary periodicity appears in the series As—Sb—Bi due to the presence of inner d- and f-subshells of electrons and their screening effect on outermost shell electrons.

For instance, oxidation state +5 in the middle of the series (Sb) is more stable than in case of As and Bi. The reason of this phenomenon is the increased energy gap between ns- and np-sublevels from As to Bi, and reduced ability to hybridisation of s- and p-orbitals.

Especially big is the difference in case of 6s²- electron pair. Significant penetration of their electron density to the nucleus through the screen of filled completely d-and f-subshells exists there. It is precisely this fact that explains the effect of 6s² electron pair inertness and therefore, the variety of Bi^{3 +} compounds and very limited number of Bi⁵⁺ compounds (they all are very strong oxidants).