Metallic Bonds

A ‘metallic bond’ is a word that refers to the sharing of a sea of valence electrons between many positively charged metal ions that occur in a collective fashion.

A sort of chemical bonding, metallic bonding is responsible for numerous of the qualities that are unique to metals, including their lustrous sheen, malleability, and heat and electricity conductivities, among others.

Some metal samples have both metallic and covalent bonding, which can be observed in some cases. Examples include gallium atoms that are covalently connected to one another and that form crystal formations that are kept together by metallic connections. The mercurous ion is also capable of forming metallic and covalent bonds.

The following are some of the factors that influence the strength of a metallic bond:

• The total number of electrons that have been delocalized.
• The amount of positive charge carried by the metal cation is measured in volts.
• The cation’s ionic radius is measured in microns.

We’ve included an illustration below that explains how electrons are delocalized over a stiff lattice of metal ions when they’re involved in a metallic connection.

When a metal is heated to the melt state, the bonds between the atoms of the metal are not destroyed. The links between metal ions are weakened instead, resulting in the ordered array of metal ions losing their defined, solid structure and becoming fluid. When the metal is heated to its boiling point, however, these connections are entirely destroyed.

Properties Attributed by Metallic Bonding

Several significant properties of metals are imparted through their metallic bonding, which make them commercially useful. Some of these characteristics are discussed in further detail later in this section.

1. Electrical Conductivity 

Electrical conductivity is a property of a substance that indicates its capacity to allow a charge to pass through it easily. Because the mobility of electrons in the electron sea is not regulated, any electric current that passes through the metal travels through it.

When a potential difference is applied to the metal, the delocalized electrons begin to move in the direction of the positively charged charge. The reason for this is that metals are generally considered to be good conductors of electric current.

2. Thermal Conductivity 

Generally speaking, the ability of a material to transmit or transfer heat is measured in terms of its thermal conductivity. Increasing the kinetic energy of electrons in a metallic substance at one end causes the kinetic energy of electrons in that area to grow. Collisions between these electrons and other electrons in the sea allow them to transfer their kinetic energy to other electrons.

The greater the mobility of the electrons, the greater the speed with which kinetic energy is transferred. These extremely mobile delocalized electrons are made possible by metallic bonding; as a result, they are able to transport heat through the metallic substance by interacting with other electrons.

3. The malleability and ductility of the material

Upon striking an ionic crystal (such as a sodium chloride crystal) with a hammer, it breaks into numerous smaller pieces. Due to the fact that the atoms in crystals are locked together in a hard lattice that is not easily distorted, this is the case. In the presence of a force (such as that applied by the hammer), this causes the crystal structure to fracture, ultimately culminating in the shattering of the crystal.

If we consider metals, the sea of electrons in the metallic bond allows for deformation of the lattice structure to take place. Consequently, when metals are struck repeatedly with hammers, the hard lattice is deformed rather than shattered. This is why metals can be pounded into thin sheets in order to save space. Metals are referred to as very ductile because their lattice structures do not shatter easily.

4. Metallic Luster 

In the presence of light incident on a metallic surface, the photon’s energy is absorbed by the sea of electrons that make up the metallic bond. The absorption of energy causes the electrons to become more excited, resulting in an increase in their energy levels. These excited electrons return to their ground states in a short period of time, generating light in the process. The emission of light caused by the de-excitation of electrons gives the metal a metallic lustre that is glossy and reflective.

5. High Melting and Boiling Points

The attractive force between the metal atoms is quite high as a result of the strong metallic bonding that exists between them. The expenditure of considerable energy is required in order to overcome this force of attraction. This is one of the reasons why metals have high melting and boiling points in the first place. Zinc, cadmium, and mercury are examples of exceptions to this rule (explained by their electron configurations, which end with ns2).

Metal-to-metal bonds can maintain their strength even while the metal is in the condition of melting. Gallium, for example, melts at 29.76 degrees Celsius but only boils at 2400 degrees Celsius. As a result, molten gallium is classified as a non volatile liquid.

Conclusion

Metallic bonding is the attraction between delocalized electrons and positively charged nuclei that occurs between these two states. It has a tremendous presence and may be found in all directions. Emission spectra are produced as a result of electrons transitioning from excited states to lower energy levels. Simple flame tests for metal salts are based on the emission spectra of the salts.