Metallic+Bonding

= = =** Physical Properties of Metals **= Solid metals typically have a characteristic luster and shiny appearance. At standard conditions, metals are also cold to the touch which relates to their high heat conductivity. Metals readily conduct electricity by allowing current to flow through them without electron displacement. The heat conductivity of a metal also generally increases with increasing electrical conductivity. Metal solids are malleable, meaning they can be hammered permanently out of shape without breaking or cracking. Metals are also considered ductile, or are capable of being drawn out into a wire. = =

=** Models Used to Describe Metals **=

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The electron-sea model suggests that metal atoms in a metallic solid contribute their valence electrons to form a uniformly distributed "sea" of electrons. The electrons present in the outer energy levels of the bonding atoms are not held by any one atom, and can move easily around the structure. The centralized metal cations holds onto the negatively-charged electrons through electrostatic attraction. ======

In general, this electron-sea model is able to account for mostof the physical properties described above. For instance, t he sea of delocalized valence electrons allow a metal to conduct electricity. In a metal wire connected to the terminals of a battery, electrons freely enter the wire from the negative terminal and are pulled through the metal toward the positive side. Mobile electrons are also what enable metallic solids to easily transfer kinetic energy and thus conduct heat throughout the solid. Metal structures can easily be stretched and flattened as a result of a metal atom’s numerous bonds with each neighboring atom. Most metals are in close-packed arrangements and are typically surrounded by 12 other atoms.

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However, as simple and useful as the electron-sea model is, it does not account for all of a metal’s physical properties. According to the model, the strength ofmetallic bonding should increase as the number of valence electrons increases. This would in turn increase the melting and boiling points. However, the actual periodic trend observed is an initial increase in boiling points and then a decrease as more valence electrons are added. ====== = =

Molecular-Orbital Model for Metals
Molecular orbitals are formed from the overlap of atomic orbitals. When atomic orbitals overlap, two types of molecular orbitals will be created:
 * One is the bonding orbital, whose overall energy is less than that of the parent atomic orbitals.
 * The other is an antibonding orbital, whose energy is higher than the parent orbitals. Large numbers of atoms within a solid will create a continuous bands of allowed energy states.

The 4s and 4p orbitals overlap more effectively than the 3d orbital because they are larger in size. The greater overlap results in a wider range in energy along the band, spanning from the bonding to the antibonding end. The 4s and 4p bands cover a large range of energy values, but due to the Pauli exclusion principle such bands are only abl e to hold two and six electrons per atom, respectively. The 3d orbital does overlap less effectively, which reduces the overall energy range of its band, but these orbitals can hold up to ten electrons per atom.

Many fundamental properties of metals arise from electrons available for bonding not completely filling a metal’s molecular orbitals. This, co mbined with an overlapping gap between the highest filled orbital and lowest unfilled orbital, enables electrons to movebetween the two with ease. The electrons in higher levels only require a small input of energy to move to previously unoccupied orbitals. This frees electrons, allowing them to move through the structure and conduct electrical or thermal energy.

The molecular-theory model can also be used to explain why the boiling point and bond strength of metals increases then decreases as more valence electrons are added, a property which the electron-sea model cannot account for. Elements from the first groups, such as zinc or cadmium, have few bonding orbitals filled which results in low bond strength. Adding more valence electrons, the bonding strength will increase. Further addition of electrons will of course increase the number of bonding orbitals, but it will also increase the number of antibonding orbitals, decreasing the metal’s bond strength and boiling point overall.

=Summary Video= media type="youtube" key="6h5chgMiRkk" height="351" width="576"

=Sources= = =
 * Brown, Theodore, H. Eugene LeMay, Bruce Bursten, and Catherine Murphy. "23.5 Metallic Bonding." // Chemistry: The Central Science //. 11th ed. Upper Saddle: Pearson College Div, 2011. Print.
 * "Metallic Bonding." // Chemical Education at University of Wisconsin //. Web. 20 Mar. 2011. <[]>.
 * "Chemical Bonding" Encyclopedia Britannica. Encyclopedia Britannica Online School Edition. Encyclopedia Britannica, 2011. Web. 20 Mar. 2011 <[]>.
 * // Bonding in Metals //. Perf. Dr. Enderle. // YouTube //. 29 Dec. 2009. Web. 20 Mar. 2011. .