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Metallic bonding is the type of bonding found in metals such as sodium, copper, magnesium, and iron. Consider sodium. We have already seen how it can form ionic bonds, such as in NaCl. We saw that it has one spare electron which it can transfer to an atom such as chlorine to become a sodium ion and achieve a stable noble gas configuration. Another way it could lose this electron is for the electron to dissociate from the atom and form a “sea” of electrons around other sodium ions that are behaving in an identical way. The positively charged sodium ions do not fly apart from electrostatic repulsion because of this intervening sea of negatively charged electrons. This is depicted in Figure 3.10. In this way, the atoms, by delocalizing the electrons (which is essentially like losing them), achieve a stable noble gas configuration.

Figure 3.10 Metallic bonding showing a “sea” of delocalized electrons around positively charged metal ions.

Metals form strong bonds. Again, using the reasoning discussed earlier, consider potassium. Its energy per bond is ∼5 × 10⁻¹⁹ J, which is equivalent to approximately 125 times the thermal energy in the bond at room temperature. Many metals are therefore stable at room temperature (e.g. iron and silver).

Interestingly, metallic bonding is not relevant for life. Life certainly does make use of metal ions in many diverse ways that will crop up throughout this book. Metal ions are particularly prominent in enzymes, in which they play a role in mediating the catalysis of chemical reactions as cofactors. These are found in many situations in cells where electron transfer is needed, such as iron–sulfur clusters used in energy-gathering processes. Examples of these ions include copper, iron, and magnesium. However, in all these examples, the metal ions are generally present as individual ions, not as bulk metallic deposits. So where are all the metal structures in life? You might like to read the Discussion Point.