Metallic Bonds: Properties, Theories, and Characteristics

Posted on Dec 12, 2024 in Chemistry

Properties of Metals

Metals, identified by their group number in the periodic table, share a set of similar properties:

  • Electrical and thermal conductivity
  • Metallic luster
  • Malleability
  • Ductility

Metals are electropositive, forming cations. They have few electrons in their valence shell. Their valence load results in positive oxidation numbers, and they are easily oxidized. All metals are solid at ambient temperature and pressure, except for mercury, which is liquid. These properties are due to their structure and the type of bonding: metallic bonding.

Metals form compact lattices with a coordination number of 8 or more. The metallic bond can be explained by three theories:

Sea of Electrons Theory

This theory states that metal atoms are ionized, and the electrons that have lost their valence become spherical cations. These cations form a three-dimensional network of tidy, compact areas, creating a field of electronic attraction. Common structures include the compact hexagonal lattice and the cubic compact lattice. The valence electrons flow freely through the lattice, forming a “sea of electrons” that neutralizes the positive charge and maintains the crystal lattice. This explains electrical conductivity and the photoelectric effect.

Delocalization Theory

Metals cannot form covalent bonds due to a lack of localized electrons; therefore, the bonds are delocalized. This theory proposes that an atom bonds with one atom at a time, then with another, rapidly changing the situation of electrons forming the bond. The valence electrons belong to all atoms in the network and can move freely.

Band Theory

According to this theory, the electrons of the bonds are located in molecular bonds formed from atomic orbitals. Half of these molecular orbitals will have more energy than the atomic orbitals, and half will have less. In the case of many atoms, there would be many orbitals. The orbital energy bands are considered the set of molecular orbital energies very close together, experimentally indistinguishable, and considered together.

There are three types of bands:

  • Occupied Bands: Formed from filled electron orbitals, they are full of electrons that cannot move because all levels within the band are filled.
  • Valence Bands: Formed from partially filled atomic orbitals. Electrons can move when a magnetic field is applied because the band is not full. It contains valence electrons. This band is formed by molecular orbitals of lower energy than atomic orbitals.
  • Conduction Bands: Formed from empty atomic orbitals. They facilitate electrical conduction, and valence electrons can jump to them and move freely around the metal crystal. They are atomic orbitals with higher energy.

Orbitals are filled from low to high energy, first filling the valence band and then the conduction band. The following can occur:

  1. The valence band is partially occupied, and the conduction band is empty.
  2. The valence band is full, and the conduction band is partially occupied.
  3. The bands overlap.
  4. A filled valence band and an empty conduction band with a forbidden energy band. No conduction is observed because it requires too much energy.
  5. A filled valence band and an empty conduction band with a small forbidden energy band. It is an insulator, but if the forbidden energy levels are covered, the atoms become conductors. These are semiconductor materials.

Key Properties of Metals

  • Electrical Conductivity: Metals are excellent conductors due to the mobility of their valence electrons. Conductivity decreases with increasing temperature as the oscillation of the nuclei increases, hindering the free movement of electrons.
  • Thermal Conductivity: Heat causes an increase in the kinetic energy of electrons, which is transmitted throughout the metal due to their mobility.
  • Luster: Metals can absorb and subsequently reflect almost all wavelengths of visible light.
  • High Density: Their crystals have a high coordination number.
  • Deformability: Any plane of atoms in the crystal structure can be moved without altering the forces holding the atoms together.

Resonance Structures

Many times, a Lewis structure can be written for the same molecule or ion.

Resonance Hybrid

A structure that would have a mix of features from all resonance forms.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

The arrangement of electron pairs around the central atom in a molecule depends on:

  1. Pairs of electrons that form bonds and those that do not are located as far apart as possible from each other and repel each other electrically. This determines the arrangement of electron pairs.
  2. The repulsive effect of a lone pair is higher than that of a bond pair, whose charge is neutralized and located between two nuclei (two atoms).
  3. Electron pairs in a double bond or three pairs in a triple bond hold atoms together in the same positions as a single bond.

Covalent Bonds: Bond Parameters

Bond Enthalpy

It is the enthalpy variation when a bond is dissociated. It is always positive, as it is an endothermic process.