Course Syllabus

Module 1

Module 2

Module 3

Module Exams

Final Exam

Lab

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handbook

Periodic table

s-elements

p-elements

d-elements

Fotogallery

Video

C [He]2s22p2 Si [Ne] 3s23p2 Ge [Ar] 4s24p2 Sn [Kr] 5s25p2 Pb [Xe] 6s26p2

IVth group of elements

labs

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photos

Properties Of the fourth group of elements

Chemical

Physical

Uses

Occurrence

Preparation

History Of Discovery

Electrode potentials

Oxidation State

Compounds Of Nitrogen subgroup. Chemical And Physical Properties Overview

C

Si

Ge

Sn

Pb

-4

CH4

SiH4

GeH4

SnH4

PbH4

-3

C2H6

-2

-1

0

1

2

CO

SiF2

GeO

SnO

PbO

GeS

SnF2

PbF2

GeF2

SnCl2

PbCl2

3

H2C2O4

Sn2+ salts

[Pb(H2O)n]2+ complexes

4

CO2

SiO2

GeO2

SnO2

PbO2

CO32-

H4SiO4

GeS2

SnF4

Pb3O4

CF4

SiF4

GeCl4

SnCl4

PbF4

CCl4

SiF62-

GeF62-

SnF62-

PbCl4

GeCl62-

[Sn(OH)6]2-

PbF62-

[Pb(OH)6]2-

Summary of general characteristics

Si is the closest analogue of carbon in the group. Silicon compounds are the mineral basis of Earth and play a role, similar to the role of carbon compounds in the development of life.

Ge, Sn, and Pb are more similar in their chemical behavior than typical elements (C and Si). In the series Ge—Sn—Pb the number of energy levels, R_a increase and their metallic nature grows, respectively: Ge is a non-metal, Sn and, especially, Pb have metallic properties.

Electron configurations of carbon subgroup

IV group p-elements include C, Si, Ge, Sn, Pb.

General formula of outermost shell configuration of these elements in the ground state is ns²np². Having two unpaired p-electrons, they can display the valence II (this trend is strengthened towards Pb). However, the excitation of the 4 unpaired valence electrons takes place easily. The valence IV occurs in this case. It becomes the most typical state of C and Si and its stability diminishes toward Pb.

The carbon subgroup includes two elements (C, Si), which are, respectively, the core of living organisms and inanimate nature.

As the size of the atom increases the ionisation energy and other physical constants of the elements alter (see Table below) and these changes are reflected in the gradual change from a typical nonmetallic element, carbon, to the weakly metallic element, lead.

Elemental C and Si are typical non-metals that have the atomic lattice with covalent bond. They have elevated hardness, melting and boiling points.

The values of these constants are reduced significantly in case of Ge, but they are still relatively large. Together with the fragility, it characterizes Ge as a diamond-like crystal with covalent bonding mechanism. At the same time, Ge has some contribution of metallic bonding. Its forbidden band energy gap essentially diminishes and conductivity elevates.

Semiconductor properties of Sn exist only below 13,2 C (-Sn,  E = 0.08 eV), at higher temperatures it has transition into the metallic ß-Sn allotrope. Pb is a metal, which does not display semi-conductive properties.

Growth of metallic properties is accompanied by a consecutive reduction in elements’ ionization energy, electronegativity.

Silicon. The valence electrons configuration of silicon is 3s²3p² (3d⁰). It differs from C atom due to 3d- vacant sublevel that leads to significant differences in behavior of these elements.

Compare bond energies of silicon with other elements:

Element	Bond energy, кJ/mol
	C	Si	H	O	F	Cl	Br	I	N	S
C	348	328	414	359	488	340	286	214	207	272
Si	328	177	320	445	568	382	310	236	331	264

Bond energy of valence electrons with Si nucleus is considerably smaller than in C due to larger R_A. Therefore, Si—Si bond is much weaker comparing with C—C (177 and 348 kJ / mol). Silicon p-orbitals do not form -bonds, and multiple bonds Si = Si or Si  Si, respectively. For these reasons, to break Si atoms chains is easier than carbon ones.

However, additional -bonding is realised in Si at the expense of the overlapping of its own free d-orbitals with twoelectronic orbitals of partners, as a result their energy is higher. Another reason for this growth in energy is attributed to the growth of ionic bonding contribution (EN difference of interacting elements is increased).

Thus, unlike carbon, Si has no so strong bonds with hydrogen. On the contrary, its bonds with O are much stronger. Therefore, there is a sharp decrease of the number of stable Si compounds with hydrogen compared to hydrocarbons and, simultaneously, diversity of Si compounds with O. These distinguishing features of silicon influence upon comparative variety of organic and mineral world: more than a million species of living organisms, on the one hand, and about three thousand different minerals, on the other hand.

Selected properties of the elements and their simple substances are shown in the Table below.

Germanium, tin, and lead. Electronic configuration of atoms in the ground state is ns²np², like typical group elements (C and Si). On the other hand, Ge and Sn have these sublevels after completed (n-1)d¹⁰-sublevels, Pb has it after 4f¹⁴5d¹⁰-sublevels. This circumstance greatly affects the properties of elements. Since 3d-and 4f-sublevels are cainosymmetric, there is a significant contraction of atoms in elements where they are formed.

Cainosymmetric (Greek “kaynos” — new, that is a new type of symmetry of orbitals) are called orbitals, which appear for the first time. They include 1s, 2p, 3d, 4f-orbitals etc. The radial distribution of electronic density for them contains single characteristic peak (without many peaks of the second order of magnitude!).

Unlike cainosymmetric, the other orbitals of the same symmetry have additional maxima on the electronic density radial distribution dependence. This leads to larger attraction of cainosymmetric electrons to the nucleus due to a significant weakening of screening effect, reducing orbital atomic radii, increase of ionization energy and, consequently, weakening metallic properties cainosymmetric elements compared with the other elements.

Contraction effects make difficult transition into the excited state:

ns²np²  ns¹np³

and reduces ns-electrons capability of forming chemical bonds. Therefore, in the series Ge—Sn—Pb decreases the highest oxidation state (+4) stability and increases stability of oxidation state (+2). Oxidation state +4 of germanium is significantly more characteristic; both oxidation states of Sn: +2 and +4 are approximately equal in their stability.