Metallic bonding is a type of chemical bonding where metal nuclei share free valence electrons. These free electrons are called delocalized because they are not confined (localized) to one atom. In contrast, valence electrons are shared between two atoms in a covalent bond and spend more time near one atom than the other in an ionic bond.
- In metallic bonding, valence electrons are delocalized or free to flow between several atoms.
- Ionic and covalent bonds involve only two atoms.
- Metallic bonding accounts for many of the key properties of metals.
The Electron Sea Model
The electron sea model is a simplistic and somewhat inaccurate view of metallic bonding, but it’s the easiest to visualize. In this model, a sea of electrons floats around a lattice of metal cations.
The main problem with this model is that the metal or metalloid atoms are not, in fact, ions. If you have a chunk of sodium metal, for example, it consists of Na atoms and not Na+ ions. The electrons are not randomly floating around the nucleus. Rather, the electron that fills an atom’s electron configuration comes from that atom or one of its neighbors. In some cases, electrons float around clusters of nuclei. It’s much like resonance structures in covalent bonding.
How Metallic Bonds Form
Like covalent bonds, metallic bonds form between two atoms with similar electronegativity values. Atoms that form metallic bonds are metals and some metalloids. For example, metallic bonds occur in silver, gold, brass, and bronze. It’s also the type of bonding in pressurized hydrogen and in the carbon allotrope graphene.
What makes metallic bonding work is that the valence electron orbitals associated with the positively-charged nuclei overlap one another. In most case, this involves s and p orbitals. Metal atoms are bound to one another by attraction between the positive nuclei and the delocalized electrons.
Bonds Formed by Metals
Metal atoms form ionic bonds with nonmetals. They form either covalent or metallic bonds with themselves or other metals. Hydrogen and the alkali metals, in particular, form both covalent and metallic bonds. So, metallic hydrogen and lithium occur. So do H2 and Li2 gas molecules.
Metallic Bonding in Homework Questions
Type of Bond Formed
The most common homework question asks whether two atoms form metallic, ionic, or covalent bonds. Atoms form metallic bonds when they are both metals. They may also form covalent bonds in certain situations, but if you have to choose one type of bond, go with metallic. Ionic bonds form between atoms with very different electronegativity values (usually between a metal and a nonmetal). Covalent bonds usually form between two nonmetals.
Predicting Properties
You can use metallic bonding to compare properties of metallic elements. For example, metallic bonding explains why magnesium has a higher melting point than sodium. The element with a higher melting point contains stronger chemical bonds.
Determine which element forms stronger bonds by examining the electron configurations of the atoms:
Sodium: [Ne]3s1
Magnesium: [Ne]3s2
Sodium has one valence electron, while magnesium has two valence electrons. These are the electrons that are delocalized in metallic bonding. So, the “sea” of electrons around a magnesium atom is twice as large as the sea around a sodium atom.
In both atoms, the valence electrons are screened by the same number of electron shells (the [Ne] core or 1s2 2s2 2p6). Each magnesium atom has one more proton than a sodium atom, so the magnesium nucleus exerts a stronger attractive force on the valence electrons.
Finally, the magnesium atom is slightly smaller than the sodium atom because there is a greater attractive force between the nucleus and the electrons.
Putting all of these considerations together, it’s no surprise magnesium forms stronger metallic bonds and has a higher melting point than sodium.
Metallic Bonding and Metal Properties
Metallic bonding accounts for many of the properties associated with metals.
- High electrical and thermal conductivity: Free electrons are charge carries in electric conductivity and thermal energy (heat) carriers in thermal conductivity.
- High melting and boiling points: Strong attractive forces between delocalized electrons and atomic nuclei give metals high melting and boiling points.
- Malleability and ductility: Metallic bonding accounts for metal mechanical properties, including malleability and ductility. Because electrons slide past each other, it’s possible to hammer metals into sheets (malleability) and draw them into wires (ductility).
- Metallic luster: Delocalized electrons reflect most light, giving metals a shiny appearance.
- Silver color: Most metals appear silver because most light is reflected off the oscillating resonance electrons (surface plasmons). Absorbed light tends to be in the ultraviolet part of the spectrum, which is outside the visible range. In copper and gold, the absorbed light is within the visible range, giving these metals a reddish and yellowish color.
How Strong Are Metallic Bonds?
Metallic bonding ranges from very strong to weak. Its strength depends largely on how much electron shells shield valence electrons from nuclear attraction. Partly this is due to relativistic effects in large atoms, so metallic bonding in mercury and the lanthanides is weaker than in lighter transition metals.
There are too many individual variations to generalize about the relative strength of metallic, ionic, and covalent bonds.
References
- Brewer, Scott H.; Franzen, Stefan (2002). “Indium Tin Oxide Plasma Frequency Dependence on Sheet Resistance and Surface Adlayers Determined by Reflectance FTIR Spectroscopy”. The Journal of Physical Chemistry B. 106 (50): 12986–12992. doi:10.1021/jp026600x
- Daw, Murray S.; Foiles, Stephen M.; Baskes, Michael I. (1993). “The embedded-atom method: a review of theory and applications”. Materials Science Reports. 9 (7–8): 251–310. doi:10.1016/0920-2307(93)90001-U
- Okumura, K. & Templeton, I. M. (1965). “The Fermi Surface of Caesium”. Proceedings of the Royal Society of London A. 287 (1408): 89–104. doi:10.1098/rspa.1965.0170
- Pauling, Linus (1960). The Nature of the Chemical Bond. Cornell University Press. ISBN 978-0-8014-0333-0.
- Rioux, F. (2001). “The Covalent Bond in H2“. The Chemical Educator. 6 (5): 288–290. doi:10.1007/s00897010509a