Chemical bonding
Chemical bonding is an energy problem. Electrons and nuclei rearrange until attractions and repulsions produce a lower-energy structure. Ionic, covalent and metallic bonding are different useful descriptions of that electron distribution—not three unrelated forces.
Energy and bonding models
The energy minimum defines the bond
At large separation two atoms interact weakly. As they approach, each nucleus attracts electrons associated with the other atom and the total energy can fall. If the nuclei are forced too close, nucleus–nucleus and electron–electron repulsions rise sharply.
The equilibrium bond length is the distance at the minimum of the potential-energy curve. Bond dissociation requires energy because separating the atoms moves the system out of that minimum.
Bond length and bond energy are therefore connected but not interchangeable: one is a distance, the other an energy difference.
Ionic bonding is an extended electrostatic lattice
In an ionic crystal, cations and anions attract many oppositely charged neighbors. Solid NaCl is not built from isolated Na–Cl molecules; it is an extended lattice whose total electrostatic stabilization involves the entire crystal.
Lattice energy grows in magnitude when ionic charges increase and ion–ion distances decrease. This helps explain why MgO is much more strongly bound than NaCl, although structure and ion sizes also matter.
When an ionic solid melts or dissolves, mobile ions can carry current. The solid itself usually conducts poorly because the ions are fixed in the lattice.
Covalent bonding: orbital overlap and shared density
In valence-bond language, a covalent bond forms when atomic orbitals overlap and two electrons occupy the internuclear region while being attracted to both nuclei.
Head-on overlap produces a σ bond. Side-by-side overlap of parallel p orbitals produces a π bond. A double bond normally contains one σ and one π bond; a triple bond contains one σ and two π bonds.
For the same pair of elements, higher bond order usually means a shorter and stronger bond because more electron density stabilizes the region between the nuclei.
Metallic bonding and bands
Metallic valence electrons are not confined to one atom pair. Their states extend across many atoms, so electrons can move through the solid when an electric field is applied.
The familiar “electron sea” is a useful first picture. Band theory gives the quantum-mechanical version: the enormous number of closely spaced atomic states in a solid forms bands of allowed electronic energies.
Delocalized bonding also helps explain why metals can often deform without the complete bonding network snapping, giving ductility and malleability.
Bond polarity makes ionic and covalent character continuous
If one atom is more electronegative, shared electron density shifts toward it and the bond becomes polar. Very large charge separation gives strong ionic character, but there is no universal Δχ threshold at which a bond instantly changes category.
Small highly charged cations can strongly polarize large anions, increasing covalent character. AlCl₃, for example, is much more covalent than a simple metal/nonmetal rule would suggest.
The right bonding model is the one that explains the property being studied: electron density, lattice energy, molecular geometry, conductivity or spectroscopy.
Quantum descriptions and delocalization
Molecular orbitals and bond order
Molecular-orbital theory combines atomic orbitals into orbitals extending over a molecule. In-phase combinations give bonding orbitals; out-of-phase combinations give antibonding orbitals, marked with an asterisk.
bond order = (bonding electrons − antibonding electrons)/2
Filling an antibonding orbital cancels some of the stabilization from bonding electrons. MO theory explains the paramagnetism of O₂, which a simple Lewis structure does not predict.
Localized bonds are sometimes only part of the picture
Lewis structures often draw electron pairs between specific atoms, but some molecules cannot be represented accurately by one localized structure. In carbonate, nitrate or benzene, several Lewis structures differ only in the placement of π electrons.
The real electron distribution is delocalized. Resonance structures are not rapidly interconverting molecules; they are alternative drawings used to represent one quantum-mechanical state whose electron density is spread over several bonds.
Delocalization can equalize bond lengths and stabilize a molecule or ion. Molecular-orbital descriptions include this naturally by allowing orbitals to extend over several atoms.
From bonding model to properties
Bonding model and material properties
Ionic solids
Often high-melting and brittle; conduct when ions become mobile.
Molecular substances
Strong covalent bonds inside molecules; bulk melting and boiling often controlled by intermolecular forces.
Network covalent solids
Extended covalent networks such as diamond or silica; often hard and high-melting.
Metals
Delocalized electrons support electrical and thermal conduction and often ductility.
Hydrogen bonding, dipole–dipole forces and dispersion forces act between molecules. They are essential to bulk properties but are not the covalent bonds holding the atoms of one molecule together.
Bond energies in reactions
Breaking a covalent bond requires energy; forming a bond releases energy. Average bond energies can therefore estimate a gas-phase reaction enthalpy by comparing the energy needed to break reactant bonds with the energy released when product bonds form.
These are average values. A C–H bond does not have exactly the same dissociation energy in every molecule because the surrounding electronic structure changes the bond.
Lewis, valence-bond and molecular-orbital models answer different questions
| Model | Best at | Important limitation |
|---|---|---|
| Lewis structures | Electron counting, connectivity, formal charge and resonance bookkeeping | Do not contain a full three-dimensional or energetic description |
| Valence-bond model | Localized bonds, orbital overlap, σ/π language and many molecular geometries | Can hide delocalization and some magnetic behavior |
| Molecular-orbital model | Delocalization, bond order, spectra and magnetism such as O₂ paramagnetism | More abstract and usually requires a larger orbital framework |
These models are complementary. A Lewis structure is not “wrong” because MO theory is deeper; it is a simpler tool designed to answer a different set of questions. Good chemical reasoning changes model when the question changes.
Exercises
Bond order
A diatomic has 8 electrons in bonding orbitals and 4 in antibonding orbitals. Find its bond order.
Solution
(8 − 4)/2 = 2.
NaCl crystal
Why is it misleading to draw solid NaCl as a collection of separate Na–Cl molecules?
Solution
Each ion interacts with several oppositely charged neighbors in an extended electrostatic lattice; there are no discrete NaCl molecular units in the crystal.
Single and triple bonds
Between the same two elements, which is usually shorter?
Solution
The triple bond, because higher bond order gives stronger stabilization between the nuclei and usually a shorter equilibrium distance.