IB Chemistry
Topics 4,14 – Chemical Bonding
*Opening Reflection Questions:
1. How is the process and the product different for chlorine reacting with potassium versus chlorine reacting with
phosphorus?
2. How do we explain and predict the bonding and shapes of molecules such as water, methane, ammonia, and ethane?
3. Explain the structure and properties of liquids such as C5H12(l) and CH3OH(l), and solids such as Na(s), NaCl(s), CO2(s),
and C(s) (diamond).
*Lewis Theory of Bonding
Review:
*ionic bonding  the electrostatic attraction between positive and negative ______ in the crystal lattice of a salt (salt =
___________compound)
e.g. Show the formation of an ionic bond between the following:
i) lithium and fluoride ii) potassium and oxygen
*What do they have in common?
*covalent bonding  the __________ of valence electrons between atomic nuclei within a molecule or complex ion
e.g. Illustrate the formation of covalent bonds in the following:
i) chlorine (Cl2)
ii) hydrogen chloride
*What do they have in common?
*The key ideas of the Lewis theory of bonding:
1. Atoms and ions are stable if they have a _____ gas-like electron structure; i.e. a stable octet of electrons.
2. Electrons are most stable when they are _____________.
3. Atoms form chemical bonds to achieve a stable ____________ of electrons.
4. A stable octet may be achieved by a transfer of electrons from __________ to ____________ atoms. This transfer
results in the formation of an __________ bond.
5. A stable octet of electrons may be achieved by the ____________ of electrons between _____________ atoms. This
sharing of electrons results in a _______________bond.
*Rules for Drawing Lewis Structures:
Step 1: Use the last digit of the group from the periodic table to determine the number of valence electrons for each atom.
Step 2: Place the electron on each of the four sides of an imaginary rectangle enclosing the central atom before pairing
any electrons.
Step 3: If there are more than four valence electrons, pair up the electrons as required to place all the valence electrons.
Step 4: Use the unpaired electrons to bond additional atoms with unpaired electrons to the central atom until a stable octet
is obtained.
*Activity: Draw the Lewis structures for all elements in Period 3.
*Lewis Structures and Quantum Mechanics
Element
Valence
Lewis Symbol
Nitrogen, N
3
Sulfur, S
2
Energy-level diagram
Electron
configuration
*In a Lewis symbol, the four sides of the imaginary square correspond to the four orbitals of the s and p sublevel. The
octet comes from the maximum of two electrons in the s energy-sublevel and six electrons in the p sublevel. In terms of
orbitals, this means two electrons in an s-orbital and six electrons in p-orbitals.
*Procedure for Drawing More Complicated Lewis Structures: (e.g. Polyatomic Ions / Oxyions / Oxyacids / molecules
that violate the octet rule)*coordinate covalent bonds
Step 1: Arrange atoms symmetrically around the central atom (usually the least electronegative atom and generally
appears in the formula, not usually oxygen and never hydrogen)
Step 2: Count the number of valence electrons of all atoms. For polyatomic ions, add electrons corresponding to negative
charge, and subtract electrons corresponding to the positive charge of the ion.
Step 3: Place a bonding pair of electrons between the central atom and each of the surrounding atoms.
Step 4: Complete the octets of the surrounding atoms using lone pairs of electrons (hydrogen = exception). Any
remaining electrons go on the central atom.
Step 5: If the central atom does not have an octet, move lone pairs from the surrounding atoms to form double and triple
bonds until the central atom has a complete octet.
Step 6: Draw the Lewis structure and enclose polyatomic ions with square brackets showing the ion charge.