1.4:
Chemical Bonds
When atoms approach each other, their nuclei repel each other, as do the electrons. Concurrently, the electrons from each atom are attracted to each other's nucleus and arrange in such a way to maximize the attractive forces between the atoms while minimizing the repulsive forces.
If this results in a net attractive force, the potential energy of the system is lowered. This interaction is called a chemical bond.
However, not every combination of atoms can form stable chemical bonds. An atom’s chemical behavior is primarily determined by what forms bonds and what type of bonds they are. Understanding the bonding between atoms that make a compound helps to comprehend and predict their molecular behavior.
Bond energy is the energy required to break a bond in one mole of a gaseous compound. It depends on the type of bonded atoms and the number of electrons involved. Bond strength increases with bond energy.
The type of bonded atoms and number of electrons involved also influence the bond length, which is the average distance between the nuclei of two bonded atoms and is inversely proportional to the bond multiplicity and bond strength.
The most common bonding interactions fall on a spectrum ranging from ionic bonds, which involve electrostatic attractions between oppositely charged ions, to covalent bonds, which share valence electrons between atoms.
Ionic bonding is the primary interaction found in ionic compounds. The constituent cations and anions are held together by electrostatic forces of attraction. Often, the ion formation is preceded by the transfer of valence electrons between participating atoms.
The atom that loses electrons becomes the cation, while the atom that accepts electrons becomes the anion. The transfer of electrons typically gives each ion the nearest noble gas configuration, which is the configuration reached with the smallest change in the number of electrons, making the ions energetically more stable.
In covalent bonding, the shared electrons interact with the bonding atoms’ nuclei and lower the potential energy. Double and triple bonds share two and three pairs of electrons between atoms, respectively. As this increases the number of electrons subject to additional attractive forces from the other nucleus, bond strength increases with bond multiplicity.
1.4:
Chemical Bonds
Atoms participate in a chemical bond formation to acquire a completed valence-shell electron configuration similar to that of the noble gas nearest to it in atomic number. Ionic, covalent, and metallic bonds are some of the important types of chemical bonds. Bond energy and bond length determine the strength of a chemical bond.
Types of Chemical Bonds
An ionic bond is formed due to electrostatic attraction between cations and anions. Often, the ions are formed by the transfer of electrons from one participating atom to the other. However, these bonds do not have a defined directionality because the electrostatic force of attraction is distributed uniformly throughout the three-dimensional space.
A covalent bond is a chemical bond formed by the sharing of electron pairs between adjacent atoms. The shared pair of electrons is called the bonding pair. Covalent bonds are directional in nature.
A metallic bond is formed between two metal atoms. Metallic bonding is described by the “Electron Sea model”. Based on the low ionization energies of metals, the model states that metal atoms lose their valence electrons easily and become cations. These valence electrons create a pool of the delocalized electrons surrounding the cations over the entire metal.
Bond Energies and Bond Length
The strength of a covalent bond is measured by the energy required to break it—that is, the energy necessary to separate the bonded atoms. Separating any pair of bonded atoms requires energy. The stronger a bond, the greater the energy required to break it.
The energy required to break a specific covalent bond in one mole of gaseous molecules is called the bond energy or the bond dissociation energy. The bond energy for a diatomic molecule is defined as the standard enthalpy change for the endothermic reaction. Molecules with three or more atoms have two or more bonds. The sum of all bond energies in such a molecule is equal to the standard enthalpy change for the endothermic reaction that breaks all the bonds in the molecule.
The strength of a bond between two atoms increases as the number of electron pairs in the bond increases. Generally, the greater the number of bonds between two atoms, the shorter the bond length and the greater the bond strength. Thus, triple bonds are stronger and shorter than double bonds between the same two atoms; likewise, double bonds are stronger and shorter than single bonds between the same two atoms. When one atom bonds to various atoms in a group, the bond strength typically decreases as we move down the group.
This text is adapted from Openstax, Chemistry 2e, Section 7.1: Ionic Bonding, Openstax, Section 7.2: Covalent Bonding, Section 10.5: The Solid State of Matter, and Section 7.5. Bond Strength: Covalent Bonds.