Chemical Bonding

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Chemical Bonding

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• Ionic bonding, metallic bonding, covalent bonding and co-ordinate (dative covalent) bonding

• Shapes of simple molecules and bond angles

• Bond polarities and polarity of molecules

• Intermolecular forces, including hydrogen bonding

• Bond energies and bond lengths

• Lattice structure of solids

• Bonding and physical properties

Learning Outcomes

Candidates should be able to:

(a) show understanding that all chemical bonds are electrostatic in nature and describe:

(i) ionic bond as the electrostatic attraction between oppositely charged ions

(ii) covalent bond as the electrostatic attraction between a shared pair of electrons and positively charged nuclei

(iii) metallic bond as the electrostatic attraction between a lattice of positive ions and delocalised electrons

(b) describe, including the use of ‘dot-and-cross’ diagrams,

(i) ionic bonding as in sodium chloride and magnesium oxide

(ii) covalent bonding as in hydrogen; oxygen; nitrogen; chlorine; hydrogen chloride; carbon dioxide; methane; ethene

(iii) co-ordinate (dative covalent) bonding, as in formation of the ammonium ion and in the Al2 Cl6 molecule

(c) describe covalent bonding in terms of orbital overlap (limited to s and p orbitals only), giving σ and π bonds (see also Section 11.1)

(d) explain the shapes of, and bond angles in, molecules such as BF 3 (trigonal planar); CO 2 (linear); CH4 (tetrahedral); NH 3 (trigonal pyramidal); H2 O (bent); SF 6 (octahedral) by using the Valence Shell Electron Pair Repulsion theory

(e) predict the shapes of, and bond angles in, molecules analogous to those specified in (d)

(f) explain and deduce bond polarity using the concept of electronegativity [quantitative treatment of electronegativity is not required]

(g) deduce the polarity of a molecule using bond polarity and its molecular shape (analogous to those specified in (d));

(h) describe the following forces of attraction (electrostatic in nature):

(i) intermolecular forces, based on permanent and induced dipoles, as in CHCl3 (l); Br2 (l) and the liquid noble gases

(ii) hydrogen bonding, using ammonia and water as examples of molecules containing –NH and –OH groups

(i) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water

(j) explain the terms bond energy and bond length for covalent bonds

(k) compare the reactivities of covalent bonds in terms of bond energy, bond length and bond polarity

(l) describe, in simple terms, the lattice structure of a crystalline solid which is:

(i) ionic, as in sodium chloride and magnesium oxide

(ii) simple molecular, as in iodine

(iii) giant molecular, as in graphite and diamond

(iv) hydrogen-bonded, as in ice

(v) metallic, as in copper

[the concept of the ‘unit cell’ is not required]

(m) describe, interpret and/or predict the effect of different types of structure and bonding on the physical properties of substances

(n) suggest the type of structure and bonding present in a substance from given information