Nature of chemical bonding

Three things hold this topic together.

Bonds inside a molecule are strong; bonds between molecules are weak. The first group are the intramolecular bonds: ionic, covalent and dative covalent. The second group are the intermolecular forces: Van der Waals forces, dipole-dipole interactions and hydrogen bonding.

A covalent bond is an overlap of two atomic orbitals. That single sentence is valence bond theory, and everything else in the topic follows from it. Head-to-head overlap gives a σ bond, sideways overlap of p orbitals gives a π bond, and overlap of two d orbitals gives a δ bond.

Whether a bond is covalent, polar covalent or ionic is a matter of degree, not of kind. The controlling quantity is the difference in electronegativity between the two atoms. Equal electronegativity gives a non-polar covalent bond, a small difference gives a polar covalent bond, and a large difference gives an ionic bond.

A chemical bond is the attractive force that holds atoms together in the form of compounds.

Bonding is an energetic process, and it runs in both directions:

• Bond making releases energy. It is exothermic.
• Bond breaking requires energy. It is endothermic.

Bonds between atoms within a molecule are called intramolecular bonds. There are three of them:

1. Ionic bonds, between ions.
2. Covalent bonds, in which the two atoms share an electron pair.
3. Dative covalent bonds, in which one atom donates a whole electron pair into a covalent bond.

Examples of the three intramolecular bonds

Ionic. Sodium chloride, NaCl, becomes Na⁺Cl⁻. This is a 1+ 1- pair.

Covalent. Two hydrogen atoms, each carrying one electron, come together and share that pair between them to give H₂.

Dative covalent. A water molecule, which has lone pairs on its oxygen, meets a hydrogen ion H⁺, which has none. The oxygen donates a lone pair to the hydrogen ion and the product is the hydronium ion.

Bonds between different molecules are called intermolecular forces, and they are much weaker than the intramolecular bonds above.

There are three of them:

1. Van der Waals forces. Found in non-polar molecules; a permanent electrostatic force.
2. Dipole-dipole interactions. Caused by the presence of polar bonds.
3. Hydrogen bonding. A very strong dipole-dipole bond, formed when hydrogen binds with an electronegative atom.

A covalent bond is a type of intramolecular bond in which two atoms share an electron pair. It is usually found between non-metals.

The simplest example is the hydrogen molecule. Each hydrogen atom brings one electron, the two electrons are shared, and the result is written H:H or H-H. This is also the standard illustration of molecular orbital theory.

Several quantities are used to describe a covalent bond:

• Bond energy.
• Bond length.
• Bond order.
• Bond angle, which is the angle between a pair of bonds that have an atom in common.
• Polarization.
• Electronegativity.

Covalently bonded substances share four properties:

1. Low melting and boiling points.
2. Poor electrical conductors.
3. Mostly insoluble in water.
4. Can be gases, liquids or solids.

Valence bond theory states that the overlap between two atomic orbitals forms a covalent bond, and that bond may be either an ordinary covalent bond or a dative one. Put the other way round, the theory describes a chemical bond as the overlap of atomic orbitals.

Covalent bonds come in three kinds, according to how the orbitals overlap.

Sigma (σ) bonds

Found in single bonds. Formed by head-to-head overlap of atomic orbitals.

The overlapping pair can be two s-orbitals, an s-orbital with a p-orbital, or two p-orbitals. Examples are H-H and Li-Li.

Pi (π) bonds

Found in double bonds, and formed by sharing two p orbitals sideways.

They are weaker than σ bonds, and they are found in unsaturated compounds such as the alkenes.

Delta (δ) bonds

Found in quadruple bonds. A delta orbital is formed when two d orbitals overlap.

Once the σ and π distinction is in place, the composition of every bond order follows:

The triple bond is the one worth having an example for: acetylene, H-C≡C-H.

Atoms hold their electrons in orbitals, which fill in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p and so on.

Hybridization is the phenomenon of promoting an electron up an orbital so that the orbital can form a bond. Described the other way, it is the joining of two or more orbitals to make a new one.

The classic illustration is carbon. In the ground state the 2s orbital holds a pair and the 2p orbitals hold single electrons. After hybridization an electron is promoted, and the resulting hybrid orbitals are the ones that go on to bond. This rearrangement of the electrons is what accounts for the shapes of the molecules actually observed.

Types of hybridization

1. sp - one electron moves from the s-orbital to a p-orbital.
2. sp² - one electron moves from the s-orbital to a p-orbital which already has 2 electrons.
3. sp³ - one electron moves from the s-orbital to a p-orbital which already has 3 electrons.

The number in the name is the number of p orbitals involved, so sp mixes one s with one p, sp² mixes one s with two p, and sp³ mixes one s with three p.

These are the two characteristic properties of a covalent bond, and both follow from the fact that a bond is an orbital overlap.

Direction. The atomic orbitals have appropriate shapes and a steric direction. Overlap of those orbitals leads to covalent bond formation, so the resulting bond has a characteristic direction. Atoms bonded to one another therefore have specific orientations relative to each other, and that is what gives a molecule its shape.

Saturation. Only two atoms can participate in one covalent bond. This is what is meant by saying the covalent bond has a characteristic saturability.

Between the same pair of atoms, the molecules of a compound can hold a single bond, or a multiple bond which may be double or triple:

• Single: H₃C-CH₃
• Double: H₂C=CH₂
• Triple: H-C≡C-H

A dative covalent bond, also called a coordinate covalent bond, is a covalent bond in which one atom donates the whole electron pair. Both shared electrons come from the same atom.

It is an intramolecular bond, like the ordinary covalent bond.

Two conditions and two consequences are worth holding on to:

• The acceptor atom must be electron deficient for a dative bond to form at all.
• In the arrow notation, the tail of the arrow indicates the source of the pair of electrons: C: + D gives C:D, written C→D.
• Once formed, the dative bond is the same as a covalent bond.
• All the bonds in the product have the same length and the same average bond enthalpy, so the dative bond cannot be picked out afterwards.

The two standard examples

Water and the hydrogen ion. Water has two lone pairs on its oxygen. The hydrogen ion H⁺ has no electrons at all, so it is electron deficient. Oxygen donates a lone pair and the product is the hydronium ion, which carries a single positive charge.

Ammonia and the hydrogen ion. Ammonia has three N-H bonds and one lone pair on its nitrogen. The lone pair is donated to H⁺ and the product is the ammonium ion, NH₄⁺. Its displayed formula shows three ordinary N-H bonds and one drawn as an arrow from nitrogen to hydrogen, which is the dative bond.

Electrons that belong to a molecule as a whole, rather than being attracted to particular atoms or particular bonds, are said to be delocalized. It helps to imagine them as clouds of electrons spread over the molecule.

Delocalization is stabilising: in some molecules the electrons delocalize and the molecule is more stable as a result.

Delocalized bonds are described as being associated with several atoms or several bonds at once, and the electrons in them are free to move.

The two standard examples are benzene, whose ring carries a delocalized electron cloud above and below it, and methanoic acid, where the delocalization runs over the carbon and both oxygens.

Conjugation is the overlap of two p-orbitals across a sigma bond. A conjugated system is therefore a structure which consists of alternating multiple and single bonds. Benzene is the cyclic example.

Conjugation matters biologically because it controls what light a molecule can absorb:

The greater the number of conjugated bonds, the lower the energy and the longer the wavelength of light that may be absorbed.

The lipids provide the two examples to remember, and both are acyclic:

• Carotene has 10 conjugated bonds.
• Lycopene has 11 conjugated bonds.

Whether a covalent bond is polar depends on whether the two atoms pull on the shared pair equally.

Non-polar covalent bonds

A non-polar covalent bond is a covalent bond between the same atoms, or more generally between two non-metals of the same electronegativity. Because the atoms attract the electrons to an equal degree, there is equal sharing of the bonding electron pair.

Examples: H₂, O₂, N₂, F₂ and Cl₂.

Polar covalent bonds

A polar covalent bond is formed between different atoms, with different electronegativities, so there is unequal sharing of the bonding electron pair. The shared electron pair sits towards the atom of greater electronegativity.

The example is H-Cl, in which chlorine has the higher electronegativity and the electron pair is drawn towards it.

Bond polarity is not the same as molecular polarity

A molecule can contain polar bonds and still be non-polar overall, because the bond dipoles can cancel. Three cases show the whole pattern:

Carbon dioxide is the instructive one: its two C=O bonds are genuinely polar, but the molecule is linear, so the two dipoles point in opposite directions and cancel exactly.

No real bond is purely covalent or purely ionic, so the degree of ionicity is used to say where a bond falls between the two. It may range from 0 to 1.

It is dependent on the size and the electronegativity of the atoms forming the bond.

A related but separate quantity is the degree of ionization, which is the proportion of neutral particles that are ionized, whether in aqueous solution or in the gas phase.

An ionic bond results from the transfer of an electron from a metal to a non-metal. It arises from elements with low electronegativity reacting with elements with high electronegativity.

• Cations are formed by the donation of electrons. They are the metals, written (Metal)⁺.
• Anions are formed by the acceptance of electrons. They are the non-metals, written (Non-metal)⁻.
• The two are held together by the electrostatic attraction of opposite charges.

The example is table salt, sodium chloride. Sodium gives up its single outer shell electron completely to chlorine, which needs exactly one electron to fill its shell. The result is written [Na]⁺ [Cl]⁻, with chlorine's completed octet drawn around it.

Ionic compounds are made of ions and are called salts or crystals. The transfer forms a crystal structure held by strong forces of attraction between oppositely charged ions.

1. High melting and boiling points.
2. Hard solids at 22 °C.
3. Non-conductors of electricity in the solid state.
4. Good conductors in the liquid (molten) state, or when dissolved in water.
5. In aqueous solution they conduct electricity.

The pattern behind all five is the same: the ions are locked in place in the crystal and can only carry a current once the lattice is broken up, either by melting it or by dissolving it.