Molecular Structure & Bonding

Elements can combine to form compounds. Atoms in compounds are held together by chemical bonds. Bonds are formed by the interaction of valence electrons. The physical and chemical properties of compounds can vary from their constituent elements. Process that involve the alteration of chemical bonds are called chemical processes, while those do that do not alter chemical bonds are called physical processes. Furthermore, weaker intermolecular forces also exist between molecules.

In forming chemical bonds, most elements follow the octet rule: Atoms tend to bond until they have eight electrons in their outer shell, forming a stable configuration akin to the noble gases. However, elements beyond the second period retain d orbitals and can thus extend beyond the eight electrons. Hydrogen also can only have two valence electrons are opposed to eight. Boron also tends to form six electrons in its outer valence shell.

Ionic Bonds

An ionic bond is a type of chemical bond that involves a metal and a nonmetal ion (or polyatomic ions such as ammonium) through electrostatic attraction; it is a bond formed by the attraction between two oppositely charged ions. Thus, atoms of low electronegativity tend to ionically bond with those of high electronegativity. Compounds with ionic bonds include sodium chloride (table salt), ammonium perchlorate (rocket propellant), and iron (II) sulfide (fool’s gold).  Compounds formed with ionic bonds are often known as salts.

Covalent Bonds

When atoms with similar electronegativities attract, they form covalent bonds, in which the atoms share electrons to achieve a noble gas configuration.  Compounds formed with covalent bonds are known as molecules.

Bond Energy

Bond energy is the energy required to separate two bonded atoms. The bond strength increases as the number of shared electron pairs increases – though not in a direct proportion. Thus, for instance, a double bond is stronger, but less than twice as strong as, a single bond. Additionally, a triple bond is less than three times stronger, but still stronger, than a single bond.

Bond Length

Bond length is the average distance between the nuclei of the atoms involved in the bond. The atoms are pulled closer together as the number of shared electron pairs increase, lowering bond length, so a triple bond is shorter than a double bond, which is shorter than a single bond.

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms of a molecule, and the lone pairs of electrons that may exist in the molecule. Valence electrons in lone pairs are represented as dots, but they also may contain lines to represent shared pairs in a chemical bond:

Examples of Lewis Structures, Source: Wikimedia User Dynablast

Examples of Lewis Structures

Some molecules retain resonance, meaning that several Lewis structures may work.

Formal Charge

Formal charge is defined for each atom as:

(Valence Electrons) – (0.5) (Bonding Electrons) – (Non-bonding Electrons)

VSPER Theory

Valence shell electron-pair repulsion (VSPER) theory is based on the notion that electron pairs repel each other and thus separate as far from each other as possible. Listed here are the shapes that molecules can retain based on their electron pair domains and number of bonded domains:

Bonding Electron Pairs Lone Pairs Electron Domains Shape Ideal Bond Angle (example’s bond angle) Example Image
2 0 2 linear 180° BeCl2 Linear-3D-balls.png
3 0 3 trigonal planar 120° BF3 Trigonal-3D-balls.png
2 1 3 bent 120° (119°) SO2 Bent-3D-balls.png
4 0 4 tetrahedral 109.5° CH4 AX4E0-3D-balls.png
3 1 4 trigonal pyramidal 109.5° (107.5°) NH3 AX3E1-3D-balls.png
2 2 4 bent 109.5° (104.5°) H2O Bent-3D-balls.png
5 0 5 trigonal bipyramidal 90°, 120° PCl5 Trigonal-bipyramidal-3D-balls.png
4 1 5 seesaw 180°, 120° (173.1°, 101.6°) SF4 AX4E1-3D-balls.png
3 2 5 T-shaped 90°, 180° (87.5°, < 180°) ClF3 AX3E2-3D-balls.png
2 3 5 linear 180° XeF2 AX2E3-3D-balls.png
6 0 6 octahedral 90° SF6 AX6E0-3D-balls.png
5 1 6 square pyramidal 90° (84.8°) BrF5 AX5E1-3D-balls.png
4 2 6 square planar 90° XeF4 Square-planar-3D-balls.png
7 0 7 pentagonal bipyramidal 90°, 72° IF7 Pentagonal-bipyramidal-3D-balls.png

Hybridization

Different types of orbitals can combine to form hydridized orbitals corresponding to certain geometries, as detailed in the chart below. Hybridization accounts for how many elements have expanded octets such as PCl5. It also accounts for why the bonds in molecules such as methane (CH4) are all the same; the single s and three p orbitals combine to form four equivalent hybridized “sp3” orbitals. The number of hybridized orbitals corresponds to the number of electron domains of the central atom of a molecule.

Pi and Sigma Bonds

Single, double, and triple bonds all have at least one sigma bond, which lies on the internuclear axis; the other bonds are all pi bonds, which lie above and below the internuclear axis:

Sigma (left) and Pi (right) bonds

 

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