Group IV-A Elements
This group of elements is placed in middle of periodic table. Group 4A (or IVA or group 14) elements of the modern periodic table includes the nonmetal carbon (C), the metalloids silicon (Si) and germanium (Ge), the metals tin (Sn) and lead (Pb), and ununquadium (Uuq).
Melting and Boiling Point:
Carbon , Silicon and Germanium have giant atomic structures and they have strong covalent bonds so their melting points are high. A decrease in melting point from C to Ge is due to increasing atomic size because of which covalent bonds are weakened. Sn and lead are metals and their atoms are bigger so their metallic bonding is weak hence their melting points are low.
Metallic character decreases down the group. In group IV-A Change from non-metal occurs. Carbon and Silicon are non-metals, Ge is metalloid while Sn and Pb are metals.
Carbon shows +4 (as in CCl4), -4 (as in Be2C), -1 (as in Na2C2) oxidation states.
Silicon shows a +4 oxidation state.
Ge, S, and Pb show +2 and +4 oxidation states.
So the most common oxidation state of group elements is +4.
Inert Pair Effect:
“The tendency of an element not to involve its pair of electrons of s orbital in bond formation is named as inert pair effect.”
If hybridization occurs then electrons of the s-orbital will not be inert and if no hybridization occurs then electrons of the s-orbital will be inert.
- Promotion of one of s-electron to p-orbital followed by sp3 hybridization of orbitals needs energy to be provided for the process.
- The formation of four covalent bonds involving the hybrid orbitals releases energy.
If the energy released in the formation of four covalent bonds is more than the energy absorbed in the promotion of an electron from s-orbital to p-orbital, then hybridization will occur and the element will not show an inert pair effect.
On the other hand, If the energy released in the formation of four covalent bonds is less than the energy absorbed in the promotion of an electron from s-orbital to p-orbital, then no hybridization will occur and the element will show an inert pair effect.
Inert pair effect increases down the group. As atomic radius increases down the group and bigger atoms make weak covalent bonds so energy released during bond formation is less than the energy absorbed during the promotion of an electron from s to p so no hybridization occurs and s-electrons remain inert.
Elements on the top of group 4 show no inert pair effect thus their oxidation state in the compounds will be +4 while lead being a larger atom often shows inert pair effect and in its most compounds its oxidation state is +2. Higher oxidation state tends to the covalent bond formation while a lower oxidation state tends to ionic bond formation. So Carbon forms covalent bonds while lead(Pb) forms ionic bonds.
Carbon, silicon and lead form tetrachlorides (CCl4, SiCl4 and PbCl4). As these elements are sp3 hybridized so their tetrachlorides are tetrahedral. The stability of chlorides decreases from CCl4 to PbCl4. Thus PbCl4 decomposes to give PbCl2. PbCl4 → PbCl2 + Cl2.
Stability of +4 oxidation state decreases down the group so C and Si form only CCl4, SiCl4 while Pb often forms PbCl2.
CCl4 does not react with water as water cannot reach carbon as carbon is a small atom and four big Cl atoms are around it. SiCl4 to PbCl4 reacts violently with water producing their oxides.
SiCl4 + 2H2O → SiO2 (White) + 4HCl
PbCl4 + 2H2O → PbO2 + 4HCl
PbCl2 is ionic so it just gets dissolves in water.
PbCl2 ⇌ Pb2+ + 2Cl–
Group IV-A elements make two types of oxides i.e. monoxides (CO, SnO and PbO) and dioxides (CO2, SnO2 and PbO2).
Nature of oxides:
Carbon and Silicon are non-metals so their oxides are covalent. Whereas Tin and Lead are metals so their oxides are ionic.
Acid-base behaviour of oxides:
As metallic character increases down the group so acidic behavior decreases down the group. Thus among dioxides CO2 and SnO2 are acidic while are GeO2, SnO2 and PbO2 are amphoteric oxides. Among monoxides, CO is neutral oxide while GeO, SnO and PbO are amphoteric.
CO2 + H2O → H2CO3
CO2 + 2NaOH → 2Na2CO3
SnO + HCl → SnCl2 + H2O
SnO + NaOH + H2O→ Na2Sn(OH)4
Structure of carbon monoxide CO:
Carbon monoxide(CO) is a divalent molecule. Carbon and oxygen contribute an unequal number of electrons in the bond formation. There is a triple bond between the two atoms including two covalent and a coordinate covalent bond.
Structure of carbon dioxide (CO2):
It is the most stable oxide of carbon. It has a linear structure having two carbon-oxygen double covalent bonds. It has a linear structure. The bond length is 1.15oA or 1.15 x 10–10m. The carbon dioxide molecule is non-polar due to its linear structure having zero dipole moment.
In solid-state (i.e. dry ice) carbon dioxide has a face-centered cubic structure.
Carbon in CO2 is sp hybridized. One half-filled sp hybrid orbital of carbon overlaps with one of the half-filled p-orbital of one oxygen atom. This overlapping takes place on the bond axis hence sigma bond is formed. The other half-filled sp hybrid orbital of carbon overlaps with one of the half-filled p-orbital of another oxygen atom. This overlapping also takes place on the bond axis hence sigma bond is formed. The remaining two half-filled unhybridized p-orbitals of carbon each laterally overlaps with the half-filled p-orbital of each oxygen atom forming Pi-bonds with them.
Structure of Silicon dioxide (SiO2):
Silicon atoms are bigger than carbon atoms so the silicon-oxygen bond length is more than that of carbon-oxygen bond length. Hence in this case lateral overlapping of orbitals will not be effective so only sigma bonds are formed. Silicon dioxide has a giant molecular structure in which each silicon atom is bonded to four oxygen atoms and each oxygen atom is bonded to two silicon atoms.
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