9.14:
Bonding in Metals
A metallic bond is a bond between two metal atoms. Compared to nonmetals, metals have low ionization energies, allowing them to lose valence electrons easily. This gives the metallic bond distinct properties in contrast to ionic and covalent bonds.
Metallic bonds and most of their properties can be explained using the simple electron sea model. Consider the metal potassium. Due to the low ionization energy, each potassium atom can easily lose its valence electron to become a cation.
These potassium cations are held together in close-packing because of their attraction to the negatively charged sea of electrons. These electrons are not confined to any single ion but are evenly distributed and relatively free to move within the metal.
The electron sea model accounts for several important characteristics of metals. For example, when a voltage difference is applied to a metal wire, like copper wire, the negatively charged electrons move freely towards the wire’s positive end, generating an electric current. This is why most metals are excellent conductors of electricity.
In contrast, ionic compounds are nonconductors of electricity in their solid form but can conduct electricity when dissolved in water. This is because, in a crystalline ionic bond, electrons are transferred from the metal to the nonmetal, but remain localized to one ion.
However, when dissolved in water, the cations and anions dissociate and can move when subjected to a potential difference, creating an electric current.
Metals are also excellent thermal conductors. According to the electron sea model, when heat is applied to one end of the metal, the electrons move freely and quickly disperse the heat throughout the metal.
Metals can easily be pounded into sheets due to their malleability or into wires due to the ductility property. Since there are no localized bonds in metals, the metal atoms can slide past each other allowing easy deformity. Electrons then flow into the new shape to accommodate the deformity.
9.14:
Bonding in Metals
Metallic bonds are formed between two metal atoms. A simplified model to describe metallic bonding has been developed by Paul Drüde called the “Electron Sea Model”.
Electron Sea Model
Most metal atoms do not possess enough valence electrons to enter into an ionic or covalent bonding. However, the valence electrons in metal atoms are loosely held due to their low electronegativity or attraction with the nucleus. The ionization energy of metal atoms (energy required to remove an electron from the atom) is low, facilitating the easy removal of valence electrons from the parent atom. The atom forms a positively charged metal ion, while the free outer electrons exist as negatively charged delocalized electron clouds. These electrons can be shared by multiple neighboring metal-cations through a strong, attractive force between these negatively and positively charged species. Such an attractive force between the negatively charged electrons and metal cations is called metallic bonds, holding the atoms together. This electron sea model accounts for most physical properties of metals such as conductance to heat and electricity, high melting and boiling points, malleability, and ductility.
Metallic Solids
The electron sea model accounts for several metallic properties, including high thermal and electrical conductivity, metallic luster, ductility, and malleability. The delocalized electrons can conduct both electricity and heat from one end of the metal to another with low resistance. The metallic bond is not between two specific metal atoms, but between metal ions and many delocalized electrons, allowing metals to deform under pressure and heat without shattering or breaking. Different metals, such as iron, mercury, or copper, differ in their physical properties, reflecting the difference in metallic bond strength among the metals.
Metallic solids such as crystals of copper, aluminum, and iron are formed by metal atoms: all exhibit high thermal and electrical conductivity, metallic luster, and malleability. Many are very hard and quite strong. Because of their malleability (the ability to deform under pressure or hammering), they do not shatter and, therefore, make useful construction materials. The melting points of the metals vary widely. Mercury is a liquid at room temperature, and the alkali metals melt below 200 °C. Several post-transition metals also have low melting points, whereas the transition metals melt at temperatures above 1000 °C. These differences reflect differences in the strength of metallic bonding among the metals.
This text is adapted from Openstax, Chemistry 2e, Section 10.5: The Solid State of Matter.