IB Chemistry - Bonding

IB Chemistry home > Syllabus 2016 > Structure and bonding > Covalent bonding

Syllabus ref: 4.2  Syllabus ref: 14.1

"Co" means sharing, "valent" refers to the electrons in the outer, or valence, shell. Hence, the term covalent bonding means to share electron pairs between two outer shells of atoms in order to bond the atoms together, making a more complex particle.

There are two ways that this can occur, either each atom provides one electron for the pair, or both of the electrons are provided (donated) by one of the atoms. In the following chapter we examine the process and the consequences of both variations of covalent bonding.

Nature of science:

Looking for trends and discrepancies. Compounds containing non-metals have different properties to compounds that contain non-metals and metals.

Use theories to explain natural phenomena. Lewis introduced a class of compounds which share electrons. Pauling used the idea of electronegativity to explain unequal sharing of electrons.

Principle of Occam's razor-bonding theories have been modified over time. Newer theories need to remain as simple as possible while maximizing explanatory power, for example the idea of formal charge.

Understandings - SL

A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.

Single, double and triple covalent bonds involve one, two and three shared pairs of electrons respectively.

Bond length decreases and bond strength increases as the number of shared electrons increases.

Understandings - HL

Essential idea: Larger structures and more in-depth explanations of bonding systems often require more sophisticated concepts and theories of bonding.

Covalent bonds result from the overlap of atomic orbitals. A sigma bond (s) is formed by the direct head-on/end-to-end overlap of atomic orbitals, resulting in electron density concentrated between the nuclei of the bonding atoms. A pibond (p) is formed by the sideways overlap of atomic orbitals, resulting in electron density above and below the plane of the nuclei of the bonding atoms.

Formal charge (FC) can be used to decide which Lewis (electron dot) structure is preferred from several. The FC is the charge an atom would have if all atoms in the molecule had the same electronegativity. FC = (Number of valence electrons)-½(Number of bonding electrons)-(Number of non-bonding electrons). The Lewis (electron dot) structure with the atoms having FC values closest to zero is preferred.

Exceptions to the octet rule include some species having incomplete octets and expanded octets.

Delocalization involves electrons that are shared by/between all atoms in a molecule or ion as opposed to being localized between a pair of atoms.

Resonance involves using two or more Lewis (electron dot) structures to represent a particular molecule or ion. A resonance structure is one of two or more alternative Lewis (electron dot) structures for a molecule or ion that cannot be described fully with one Lewis (electron dot) structure alone.

Applications and skills - SL


Applications and skills - HL

Prediction whether sigma (s) or pi (p) bonds are formed from the linear combination of atomic orbitals.

Deduction of the Lewis (electron dot) structures of molecules and ions showing all valence electrons for up to six electron pairs on each atom.

Application of FC to ascertain which Lewis (electron dot) structure is preferred from different Lewis (electron dot) structures.

Deduction using VSEPR theory of the electron domain geometry and molecular geometry with five and six electron domains and associated bond angles.

Explanation of the wavelength of light required to dissociate oxygen and ozone.

Description of the mechanism of the catalysis of ozone depletion when catalysed by CFCs and NOx.

In Section 2.1