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The Trend from Non-Metal to Metal in Group 4 Elements

This page explores the trend from non-metallic to metallic behavior in the Group 4 elements (carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb)). It describes how this trend is evident in the structures and physical properties of the elements, and attempts to explain the trend.

Structures of the elements

The trend from non-metal to metal down the group is evident in the structures of the elements themselves. Carbon, at the top of the group, forms large network covalent structures in its two most familiar allotropes: diamond and graphite. Diamond has a three-dimensional structure of carbon atoms each bonded covalently to 4 other atoms. This diagram shows a representative portion of that structure:

 

 

This structure is also found in silicon and germanium and in one of the allotropes of tin, "grey tin" or "alpha-tin". The more common allotrope of tin ("white tin" or "beta-tin") is metallic, with its atoms held together by metallic bonds. The structure is a distorted close-packed arrangement. In a close-packed structure, each atom is surrounded by 12 neighboring atoms.

In lead and the heavier elements, the atoms are arranged in a 12-coordinated metallic structure.

From this information, it is clear that there is a trend from the typical covalency found in non-metals to the metallic bonding in metals, with an obvious inflection point between the two common tin allotropes.

Physical properties of the elements

Melting points and boiling points

If the trends in melting and boiling points down Group 4 are examined, it is difficult to comment on the shift from covalent to metallic bonding. The trends reflect the increasing weakness of the covalent or metallic bonds as the atoms get bigger and the bonds get longer. This trend is shown below:

 

The low value for tin's melting point compared with that of lead is presumably due to the distortion in tin's 12-coordinated structure. The tin values in the chart refer to metallic white tin.

Brittleness

A much clearer distinction between nonmetals and metals is shown when the brittleness of the elements is considered.

  • Carbon in its diamond allotrope is very hard, reflecting the strength of the covalent bonds. However, if a diamond is hit with a hammer, it shatters.
  • Silicon, germanium and grey tin (all with the same structure as diamond) are also brittle solids.
  • However, white tin and lead have metallic structures. The atoms can move around without any permanent disruption of the metallic bonds; this leads to typical metallic properties like malleability and ductility. Lead in particular is fairly soft.

Electrical conductivity

  • Diamond does not conduct electricity. In diamond the electrons are all tightly bound and not free to move.
  • Silicon, germanium and grey tin are semiconductors.
  • White tin and lead are metallic conductors.

This information shows clear trend between the typically non-metallic conductivity behavior of diamond, and the typically metallic behavior of white tin and lead.

Explaining the trends

One important characteristic of metals is that they form positive ions. This section examines factors which increase the likelihood of positive ions being formed down Group 4.

Electronegativity

Electronegativity measures the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, in which the most electronegative element (fluorine) is assigned an electronegativity of 4. The lower the electronegativity of an atom, the less strongly the atom attracts a bonding pair of electrons. That means that this atom will tend to lose the electron pair towards whatever else it is attached to. The atom we are interested in will therefore tend to carry either a partial positive charge or form a positive ion.

Metallic behavior is usually associated with low electronegativity. The trend in electronegativity in Group 4, and its implications for metallic behavior, can be examined using the figure below:

 

 

Electronegativity clearly decreases between carbon and silicon, but beyond silicon there is no definite trend. There therefore seems to be no relationship between the non-metal to metal trend and electronegativity values.

Ionization energies

When considering the formation of positive ions, a good start includes describing how ionization energies change down Group 4. Ionization energy is defined as the energy required to carry out each of the following changes (reported in kJ mol-1):

First ionization energy:

\[ X(g) \rightarrow X^+(g) + e^-\]

Second ionization energy:

\[ X^+(g) \rightarrow X^{2+}(g) + e^-\]

and so on for subsequent ionizations.

None of the Group 4 elements form 1+ ions, so looking at the first ionization energy alone is not helpful. Some of the elements do, however, form 2+ and (to some extent) 4+ ions. The first chart shows how the total ionization energy needed to form the 2+ ions varies down the group. The values are all reported in kJ mol-1.

The ionization energies decrease down the group, although there is a slight increase at lead. The trend exists because:

  • The atoms are getting bigger because of the extra layers of electrons. The farther the outer electrons are from the nucleus, the less they are attracted; therefore, they are easier to remove.
  • The outer electrons are screened from the full effect of the nucleus by the increasing number of inner electrons.
  • These two effects outweigh the effect of increasing nuclear charge.

Examining the ionization energy required to form 4+ ions, the pattern is similar, but not as simple, as shown below (values again reported in kJ mol-1):

 

 

 

Large amounts of ionization energy are required to form 2+ ions, and even more energy is required for 4+ ions. However, in each case there is a decrease in ionization energy down the group; this implies that tin and lead could form positive ions. However, there is no indication from these figures that this is likely.

Carbon's ionization energies are so large that there is essentially no possibility of it forming simple positive ions.

Contributors

Jim Clark (Chemguide.co.uk)