Chapter 11
Chemical Bonds
The Formation of Compounds from Atoms

Electronic configurations for ions resemble those for noble gas elements. That is, they have 8 valence electrons. The octet rule also applies for covalent compounds. Shared electrons are usually present in groups of eight around each atom.

Valence electrons and Lewis Structures

• Know how to count electrons and valence electrons. Valence electrons are the electrons in the highest principal quantum level.
• Dot structures are more important for molecules than for atoms. For both, they show the number of valence electrons, and whether or not they're paired.

IONIC COMPOUNDS

• Between metals and nonmetals. Transfer of electrons to form + and - ions.
• Remember electronic configurations. Short cut: group #s on periodic table.
• Given two elements, give the name and write the formula for the ionic compound that they will form.
• Use Roman numerals for ions that can have more than one charge. Roman numeral is the charge on the cation.
• Properties of ionic compounds. (These are not molecules!)
• Lattices: account for brittleness; conductivity properties (solid, liquid, solution...)

COVALENT COMPOUNDS

• Between nonmetals and nonmetals. Shared electrons.
• Binary covalent compound nomenclature: same as for ionics, but with prefixes.
• Lewis Dot Structures are crucial to understanding shapes and properties of covalent molecules.

Rules for drawing Lewis Dot Structures:

1. Sum valence electrons from all atoms in the molecule. Add one for each negative charge, subtract one for each positive charge.
2. Draw a trial structure.
   • Put the least electronegative (most metallic) atom in the center.
   • Assume every other atom is bound to the central atom, unless told otherwise.
   • Use a dash for each bond. (1 dash = 1 bond = 2 electrons)
   • Put an octet of electrons around each atom, including the central one.
3. Count electrons in trial structure. Should equal the number of valence electrons.
   • If trial structure contains too many electrons, make a double (or triple) bond. (-4+2=-2)
   • If trial structure has too few electrons, add electron pairs to central atom. (violate octet)
4. Check octet rule again.
5. Check number of electrons again.

Additional general rules:

• General number of connected atoms. (halogens=1, chalcogens=2, Gp.V=3, Gp.IV=4, Gp.III=3, Gp.II=2, H=1)
• Usually only N, C, O, S, and P can support multiple bonds.

Polyatomic Ions

• Polyatomic ions have covalent bonds between their atoms, but they form ionic compounds as a unit.
• To determine compounds formed, just look at the charge; treat the polyatomic ion as a coherent unit.
• Learn the names of polyatomic ions. See tables in book. Find the naming patterns.

Rules for Determining Molecular Shapes

1. Draw Lewis Structure
2. Count the electron groups on the central atom. (Note: One electron group = 1 lone pair, 1 single bond, 1 double bond, 1 triple bond.)
3. Determine the electron group geometry that minimizes the electron pair repulsions.

   # of Electron Groups	Geometry of Electron Groups
   2	Linear
   3	Trigonal planar
   4	Tetrahedral
   (5)	(Trigonal bipyramidal)
   (6)	(Octahedral)

4. Place the non-bonding (lone) pairs as far apart from each other as possible.
5. The molecular geometry is the arrangement of the atoms in space. Different from electron group geometry. The electron pairs take up space, but are not included when describing the shape of the molecule.)
6. Molecular Geometry names: linear, bent, trigonal planar, trigonal pyramid, tetrahedron, (trigonal bipyramid), (T-shaped), (see-saw), (octahedral), (square pyramid), (square planar).

Electronegativity. Polar and nonpolar bonds. Polar and nonpolar molecules.

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