This module covers ionic and covalent chemical bonding. This material is covered in Chapter 8 of Brown.
Topics for lecture 7 [Brown chapter 8]
1. Ionic and covalent bonds organize atoms into compounds or molecules. Motivation: to fill all atoms’
2. Octet rule: all atoms want full valence shells; same number of valence electrons as the noble gas in their
row of the periodic table.
a. Ionic bonds form when metals transfer electrons to non-metals. The metals become cations while the
non-metals form anions.
b. Covalent bonds are a pair of shared valence electrons; each atom donates one ve-.
c. Multiple covalent bonds are stronger and shorter than single covalent bonds.
d. Bonding releases energy because the energy state of bonded atoms is lower (more stable) than the
energy state of elemental (unbounded) atoms.
e. Electronegativity determines whether covalent bonds are non-polar or polar.
f. Resonance hybrids represent covalent bonds with ‘intermediate’ electron sharing and are more stable
than structures without resonance.
i. Lewis dots represent the number of valence electrons in an atom: atom’s letter symbol is surrounded by
its valence electrons as dots. Group number = valence electron number.
ii. Atoms with 4 ve- neither gain nor lose. Atoms with less than 4 ve- lose them and become cations. Atoms
with more than 4 ve- gain more and become anions.
iii. Molecular enthalpies of formation are negative, because forming bonds to create molecules releases
iv. Ionic compounds form crystals: repeating three-dimensional arrays of alternating cations and anions.
Crystal size grows as cation/anion units are added.
v. Lattice energy is the energy required to separate one mole of ionic compound into its component
elements in their gas states. Lattice energy (enthalpy) is positive because an energy input is required to
break ionic bonds.
vi. Lattice energy increases as 1) ionic charge increases, and 2) as internuclear distance (aka both ionic
radii) decreases. Charge is the dominant factor.
vii. Covalent bond distance minimizes repulsion between the electrons and protons of bonding nuclei,
and maximizes the attractive forces between one ion’s electrons and the other’s nucleus.
viii. In Lewis structures of molecules, a line (or :) shows a covalent bond. Venn diagrams show electrons
that ‘belong’ to one atom or are shared by bonded atoms.
ix. Multiple bonds are formed by multiple pairs of two electrons. Shared electron pairs act as glue, and the
more glue between nuclei, the closer they are and the shorter their bond is.
x. Electronegativity is an atom’s ability to hold on to its own ve- and its ability to steal ve- from nearby
xi. The difference in electronegativity values of covalently bonded atoms determines whether their bonds
are non-polar (0 – 0.5) or polar (0.5 – 2.0) or ionic (> 2.0).
xii. Polarity arrows point towards the more electronegative atom and have a ‘+’ crosshatch at the less
xiii. Dipolar charges are partial charges. In a polar bond, the more electronegative atom has a dipolar
negative charge (δ-) and the less electronegative atom has a dipolar positive charge (δ+).
xiv. Guidelines for drawing Lewis dot structures: 1) count total ve-; 2) connect atomic symbols with lines;
3) add dots to fill valence shells of peripheral atoms; 4) add dots (or use multiple bonds) to fill valence
shell of central atoms.
xv. Formal charges are calculated on each atom as: FC = (atom’s ve- #) – (dots + sticks). Charges are
usually found on more electronegative atoms.
xvi. Resonance structures are Lewis dot structures (of the same molecule) that have alternate
arrangement of electrons (or alternate bonding patterns).
xvii. Reality is resonance hybrids rather than resonance structures. Evidence: resonance bonds have
identical bond length, intermediate between single or double bonds.
xviii. There are three types of exceptions to the octet rule: 1) molecules with an odd number of electrons;
2) molecules in which an atom has less than an octet; and 3) molecules in which an atom has more
than an octet. Atoms in case (3) are below row 3 of the periodic table.
Resources for students:
Links and items of interest: