Section 6-2: Covalent Bonding and Molecular Compounds
Objectives:
Define molecule and molecular formula.
Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy.
State the octet rule.
List the six basic steps used in writing Lewis structures.
Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.
Explain why scientists use resonance structures to represent some molecules.
A. Molecular Compounds
1. molecule – a. individual unit capable of existing on its own b. neutral group of atoms that are held together by covalent bonds diatomic molecule – molecule that contains only two atoms
2. molecular compound – chemical compound that has molecules as its simplest units
3. composition of a compound is given by its chemical formula
4.
5. chemical formula – each kind in a chemical compound by using atomic symbols and numerical subscripts molecular formula – formula that indicates the relative numbers of atoms of formula that shows the types and numbers of atoms combined in a single molecule of a molecular compound
B. Formation of a Covalent Bond
1. most atoms have lower potential energy when they are bonded to other atoms than they have as they are independent particles
2. change in potential energy in the formation of hydrogen-hydrogen bond;
Fig. 5, p. 179, top

a. proton of one atom and the electron of the other atom are attracted to each other which corresponds to a decrease in the total potential energy of the atoms; Fig. 6, p. 179, bottom b. two nuclei and two electrons repel each other which results in an increase in potential energy
6.2b c. two forces cancel out to from a covalent bond at a distance at which the potential energy is at a minimum
C. Characteristics of the Covalent Bond
1. bond length – average distance between two bonded atoms at their minimum potential energy; hydrogen: 75 pm
2. bond energy – energy required to break a chemical and form neutral isolated
3. overlapping orbitals; atoms; same amount of energy that is released when the bond forms
Table 1, p. 182, top; bond lengths and bond energies stable; outer electrons s
Fig. 8, p. 182, middle
and p a. covalent bond results when shared electrons form overlapping orbitals b. results in a stable noble-gas configuration of helium, 1s 2 hydrogen atom
D. The Octet Rule
1. the noble-gas atoms exist independently in nature
, for each
2. noble gases are unreactive because the electron configuration is exceptionally
subshells are completely filled by a total of eight

3. other atoms can fill their outermost
4. octet rule – chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has eight electrons in its highest energy level;
Fig. 9a, 9b; p. 183, middle s and electrons through covalent bonding p
6.2c
orbitals with electrons by sharing
5. exceptions to the octet rule include those atoms that cannot fit eight electrons, and for those that can fit more than eight electrons, in their outermost orbital a. hydrogen: forms bonds in which it is surrounded by two electrons; He b. boron: has three valence electrons therefore form bonds in which it is surrounded by six electrons c. main group elements in Period 3 and up can form bonds with valence
E. Electron-Dot Notation
1.
, involving more than eight electrons including electron-dot notation – electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol;
Fig. 10, p. 184, top left
2. element symbol represents the nucleus and all inner-shell electrons d expanded
orbitals
Sample Problem B; p. 184, middle
F. Lewis Structures
1. electron-dot notation can be used to represent molecules a. pair of dots between the symbols represents the shared electron pair of the hydrogen-hydrogen bond; H:H or H–H

6.2d b. molecule of c.
F
2 shows a pair of dots between the symbols and six additional dots, three pairs, around each symbol unshared pair, lone pair – pair of electrons that is not involved in bonding and that belongs exclusively to one atom
2. Lewis structure – formula in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons
3.
4. structural formula – single bond – covalent bond in which one pair of electrons is shared between two atoms formula that indicates kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule;
Sample Problem C; p. 185,6 bottom, top
F–F, H–Cl

6.2e
G. Multiple Covalent Bonds
1. double bond – covalent bond in which two pairs (4) of electrons are shared between two atoms a. often found in molecules containing carbon, nitrogen, or oxygen
2.
3. b. shown by two parallel dashes or two side-by-side pairs of two dots triple bond – covalent bond in which three pairs (6) of electrons are shared between two atoms multiple bond – double and triple bonds collectively a. double bonds have greater bond energy and are shorter than single bonds; Table 2, p. 187, bottom b. triple bonds are even stronger and are shorter than double bonds c. Lewis structures involving carbon, nitrogen, or oxygen often contain multiple bonds
H. Drawing Lewis Structures with Many Atoms
1. Gather information.
2. Arrange the atoms.
3. Distribute the dots.

6.2f
4. Draw the bonds.
5. Verify the structure.
Sample Problem D; p. 188, top
I. Resonance Structures
1. some molecules and ions cannot be represented adequately by a single Lewis
2. structure; resonance – an example is the structure for ozone, O
3 bonding in molecules or ions that cannot be correctly represented by a single Lewis structure