Chemical Bonds: Types, Properties, and Theories

Posted on Nov 4, 2024 in Chemistry

Chemical Bonds

Chemical Bond: The noble gases and metal fumes are composed of isolated atoms. Atoms can be closely allied and may be atoms of the same element or belong to different elements. A chemical bond is any mechanism of ligation or chemical bonding between atoms. These chemical bonds between atoms are formed and broken in chemical reactions. The formation of bonds results from a favorable energy balance; the bonded atoms form a system with less energy (more stable) than separated atoms.

Atoms Separated → Atoms + Energy

The breaking of bonds requires energy input.

Forming Links and Energy Stability

When two (or more) atoms are joined, the energy of the system changes. It occurs as follows:

  • Initially, the atoms are separated and do not interact.
  • As they approach each other, attractive forces cause the system to stabilize, and the energy decreases until a minimum distance is reached, which corresponds to the bond length.
  • If they continue to approach, repulsive forces are exerted between the nuclei, leading to instability and an increase in energy.

Two (or more) atoms unite to form a more stable system with lower energy than the separate atoms.

Types of Chemical Bonds

  1. Ionic Bond: Due to electrostatic attraction between ions. Occurs when metals combine with non-metals by electron transfer from the metal to the nonmetal (e.g., NaCl).
  2. Covalent Bond: Occurs between nonmetallic elements. Involves the sharing of pairs of electrons between adjacent atoms (e.g., H2O).
  3. Metallic Bond: Occurs in metals. Electrons are shared collectively among all the atoms that make up the metal.

Lewis Theory of Chemical Bonding

Lewis’s theory, based on the electronic nature of the bond, was the first modern theory of chemical bonding.

Principles of the Lewis Theory

The crucial idea comes from the Bohr atom model and involves dividing electrons into two groups: internal and valence. Only the valence electrons contribute to the bond. The theory is based on the following principles:

  1. In some cases, electrons are transferred from one atom to another, forming positive and negative ions that are attracted by electrostatic forces, resulting in an ionic bond.
  2. In other cases, one or more pairs of electrons are shared, resulting in the formation of a covalent bond.
  3. The transferred or shared electrons stabilize the electron configuration of atoms, as they achieve the configuration of noble gases, which have 8 electrons in the outermost shell (octet rule).

C = N – D and S = D – C, where:

  • C: Number of shared electrons
  • N: Number of electrons needed
  • D: Number of available valence electrons
  • S: Number of lone pairs of electrons (non-bonding electrons)

Number of bonds (E): E = (N – D) / 2

Multiplicity and Covalent Bond Order

Electrons are always shared in pairs, represented by a dash for simplicity. This theory allows for multiple covalent bonds. The bond order indicates its multiplicity. The bond order was initially considered to be a whole number and not greater than four.

1. Limitations and Improvements of the Lewis Theory

The octet rule is useful, especially for the bonds of organic molecules. The basic elements of organic chemistry (C, H, O, and N) are well suited to Lewis’s predictions. However, the original theory is not a universally valid rule. Transition elements systematically violate this rule because the sublevel being filled is the d sublevel, which can hold up to 10 electrons.

Incomplete Octet (Hypovalency)

Some elements, such as Be and B, tend to be hypovalent, with an electronic structure that does not reach the noble gas configuration. They violate the octet rule by default. Al and Ga show a similar trend.

Expanded Octet (Hypervalency)

Elements can have 10 or 12 valence electrons; the octet is expandable. The atom is hypervalent. Hypervalency is only possible in elements of the third period or higher (P, S, Cl, etc.). Elements of the first and second periods never become hypervalent.

Species with Odd Numbers of Electrons

Some structures cannot fulfill the octet rule, such as NO (11 electrons) and NO2 (17 electrons).

Resonance

Several equivalent Lewis formulas, but not identical, are possible for a single molecular species. The correct structure is an intermediate between all possible equivalent formulas. None of the individual formulas is valid. The correct formula is a blend or hybrid combination of all of them, called a resonance hybrid. Each hybrid is usually represented in brackets, and in the case of ionic species, the ion charge is given as a superscript outside the bracket. Resonance is an oscillation between two structures. The order of a given bond is the average value between the resonant formulas. The concept of resonance introduces the possibility of fractional bond orders.

2. Quantum Theories of Covalent Bonding

The concept of orbitals plays a key role. The two main quantum theories are the molecular orbital theory (MO theory) and the valence bond theory (VB theory).

Molecular Orbital Theory

MO theory is more complete but also more complex. It studies the molecule as a set of nuclei and electrons, for which the Schrödinger equation must be solved. Molecular orbitals are obtained, which indicate the regions of space where it is most likely to find electrons in the molecule. The simplest case is the hydrogen molecule. Of the molecular orbitals formed, one has lower energy and is more stable than the atomic orbitals. This is the bonding MO, and the electrons that occupy it favor bond formation. MOs can hold up to 2 electrons. The other MO is less stable, and the electrons that occupy it hinder bond formation. This is called the antibonding MO. This theory also predicts the existence of delocalized MOs in other cases.

Valence Bond Theory (VB Theory)

VB theory presents a more intuitive approach to the bond. The molecule is understood as a set of bonded atoms. Bonds are formed due to the overlapping of atomic orbitals, which must have unpaired electrons with opposite spins. Electrons with all equal quantum numbers cannot occupy the same orbital.

Valence Chemistry

Electrovalence

The ionic or electrovalence is the valence of an element when forming ionic compounds. It is unsigned. Often, the ionic valence indicates how many electrons an atom needs to gain or lose to achieve the noble gas configuration. Transition metals usually do not follow this criterion.

Covalence

The covalent valence indicates how many covalent bonds an element forms, or how many pairs of electrons it shares. It is unsigned and coincides with the number of unpaired electrons the atom has before bonding.

Multiple Bonds in VB Theory

The shape and size of atomic orbitals allow for more than one type of overlap. Frontal overlap is the most energetically favorable and the only one present in single bonds. Other overlaps give rise to pi MOs, which are responsible for multiple bonds.

Dative or Coordinate Covalent Bonds

When the shared pair of electrons is contributed by only one of the atoms, the bonding MO arises from the overlap of atomic orbitals, one of which is initially filled and the other empty. The difference between a normal covalent bond and a dative or coordinate bond is conceptual and depends on how we imagine the bond formation. The important thing is that in both cases, the bond is due to the sharing of electron pairs.

3. Properties of Covalent Bonds

Bond Length and Multiplicity

Bond length is the average distance of separation between the nuclei of bonded atoms. The higher the order of a bond, the shorter the bond length.

Bond Energy

Energy is released when atoms bond and is consumed when they separate. The bond dissociation energy is the energy required to break one mole of gaseous covalent species. This is an important quantity because it reveals the strength of each bond. Bond lengths and energies are average values. Sigma bonds are stronger than pi bonds.

Polarity of Covalent Bonds and Electronegativity

In homopolar covalent bonds, the atoms that are bonded are identical, and the electrons are shared equally. The electron cloud is distributed symmetrically between the nuclei, and no electrical poles are created. The bond is nonpolar. If the bonded atoms are not equal, the electrons cannot be shared equally. The electron cloud is deformed, becomes asymmetric, and there are excesses and deficiencies of negative charge, creating electrical poles. The bond is polar. The atom that is more attractive to the electron pair(s) acquires an excess of negative charge, and the atom that attracts less acquires a deficiency of negative charge. The polarity of a covalent bond is measured by a physical quantity called the dipole moment. It depends on the value of each charge of the dipole and the separation between them. The electronegativity of each element allows us to predict whether a covalent bond will be polar or not. If the atoms that bond have similar electronegativities, the bond is nonpolar, but if there is a significant difference in electronegativity, the bond will be polar.

Dipole

A dipole is a system of two point charges of equal magnitude but opposite sign, separated by a distance. The net charge of the dipole is zero, but the separation of charges creates an electric field whose intensity depends on the dipole moment.

4. Ionic Bond

The ionic bond is the chemical bond formed by electrostatic attraction between oppositely charged ions.

Formation of Ion Pairs

When elements with a large difference in electronegativity combine, electron transfer occurs from the more electropositive atom to the more electronegative atom, an extreme case of polar covalent bond formation. The formation of the cation consumes more energy than is generated in the formation of the anion.

Lattice Energy and Born-Haber Cycle

While the formation of an ion pair is energetically favorable, under normal conditions, there are no ion molecules, meaning ion pairs are not isolated. An ionic bond is collective, resulting in an ordered three-dimensional structure called an ionic crystal. Ionic compounds are solid, and the formulas that represent them are empirical. Their outward appearance reveals the crystal structure of their interior. This behavior occurs because the electric force is long-range and acts in all directions. The cohesive energy of a solid is the energy that holds together the particles that compose it. In ionic compounds, it is called the lattice energy and is defined as the energy required to break apart one mole of an ionic crystal and transform its ions into the gaseous phase. The lattice energy is the fundamental quantity that indicates the strength of an ionic crystal, and it determines the main properties of ionic solids. The Born-Haber cycle is used to calculate the lattice energy. The lattice energy depends on three factors: the charge of the ions, the separation distance between charges, and the spatial distribution of ions. The first two factors can be estimated in a first approximation by Coulomb’s law. Therefore, small ions with high charges have large lattice energies. For the same charge, larger ions generate lower lattice energies.

Properties of Substances According to Their Bond

Ionic Compounds

  • No molecules; crystalline structure held together by electric forces.
  • High melting and boiling points.
  • Hard and brittle.
  • Electrical insulators in the solid state and poor conductors of heat; when molten, they conduct electricity.
  • Soluble in highly polar liquids and conduct electric current when dissolved.

Metals

  • Eject electrons when heated or exposed to light.
  • Show a characteristic brightness when polished.
  • Excellent conductors of electricity and heat.
  • Medium or low hardness with good mechanical properties (elastic, ductile, and malleable).
  • High melting points.
  • Do not dissolve in ordinary molecular solvents (polar or nonpolar). Dissolve well in each other, forming alloys and amalgams.

Covalent Substances

Divided into two groups with different properties:

  1. Molecular Substances: The majority of covalent substances.
  2. Covalent Network Solids: Pure carbon (graphite and diamond), quartz, and corundum. No molecules; atoms are held together by strong, localized covalent bonds in a crystal lattice. No free electrons. Properties:
    • Solid at room temperature with very high melting points.
    • Very rigid; do not deform easily and fracture under pressure.
    • Hard.
    • Very good electrical insulators and poor conductors of heat.
    • Insoluble.