Metallic Bond

Electrostatic attractive force between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions causes a type of chemical bonding which is known as Metallic bonding. It may be thought of as the distribution of free electrons among a structure of positively charged ions in a solution (cations). Numerous physical properties of metals, such as strength, ductility, thermal and electrical conductivity, opacity, and lustre, are explained by the metallic bonding between the atoms and molecules. 

Even in its purest form, metallic bonding is not the only sort of chemical bonding that a metal may form with other substances. Elemental gallium, for example, is made up of pairs of atoms that are covalently bonded together in both the liquid and solid states; these pairs combine to create a crystal structure that has metallic bonding between them. The mercurous ion is another example of a covalent connection formed between two metals.

Properties of Metallic bonding 

Several significant characteristics of metals are imparted through their metallic bonding, which make them economically useful.

1. Malleability and ductility  

Metals are referred to as ductile and malleable (can be beaten into sheets) (can be pulled out into wires). This is due to the atoms’ capacity to roll over one another into different configurations without fracturing the metallic connection.

2. High melting and boiling points

The attractive force between the metal atoms is particularly high as a result of the strong metallic bonding that exists between them. The expenditure of considerable energy is necessary in order to counteract 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).

3. Electrical Conductivity

Electrical conductivity is a property of a substance that indicates its capacity to enable 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 positive  charge. The reason for this is that metals are typically considered to be strong conductors of electric current.

4. Metallic Lustre 

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. Thermal Conductivity 

Generally speaking, the capacity of a substance to transmit or transfer heat is measured in terms of its thermal conductivity. Increasing the kinetic energy of electrons in a metallic material 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 transmit their kinetic energy to other electrons.

Examples of Metallic Bonds

The metallic bond is a type of binding that is typically encountered in metals. Here are a few illustrations:

1. Sodium (Na)

Sodium contains a single electron in its outermost orbital, which is the 3s orbital, which is the most stable. In an arrangement of sodium atoms, the outermost electron of one atom exchanges space with the equivalent electron on a nearby atom, resulting in the formation of an electrostatic field. As a result, a molecular orbital with a 3s configuration is generated. Each sodium atom is surrounded by eight other sodium atoms in its immediate vicinity. The sharing occurs between a core sodium atom and the 3s orbital of each of its neighbours.

2. Magnesium (Mg)

Magnesium contains two electrons in its outermost shell, the 3s shell, which is the outermost of its three electron shells. Both of these electrons have a delocalized configuration. Magnesium forms metallic bonds in the same way as sodium does, with the exception that magnesium has a higher electron density than sodium. Furthermore, the charge on each magnesium nucleus is double that of the charge on each sodium nucleus. As a result, the attraction between the nuclei and the delocalized electrons will be larger than the attraction between the nuclei and sodium. Magnesium has a stronger bond strength than other metals, in general.

3. Aluminium (Al)

Aluminium contains three valence electrons in the 3s orbital, which means it has three valence electrons. A positive charge of +3 is achieved when all three electrons of aluminium atoms are lost. Aluminium ions are formed when this occurs. Despite the fact that these positively charged ions repel one another, the negative electrons in the block keep them all together in the block. A consistent pattern of arrangement is achieved as a result of the electrons being shared among the cations. 

Conclusion

A ‘cloud’ of free moving valence electrons bonded to the positively charged ions in a metal is a type of chemical bond known as metallic bond. Metallic bonds have a completely different structure than ionic and covalent ones. Only metals have a metallic connection. The electrons are separated from the atoms and dispersed throughout the metal, allowing them to travel around freely. Interactions between ions and electrons, on the other hand, are still common. A binding force is formed as a result of these interactions, which binds the metallic crystal together. A metallic connection is based on this force.