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Revised 12/2015 Chemistry 1104 L
LEWIS STRUCTURES
The purpose of this experiment is to gain practical experience of drawing lewis structures and to use molecular models to represent the three-dimensional shapes of molecules, which will lead to a better understanding of the concepts of covalent bonding and molecular structure. With the molecular model kits provided, you will build several models of small molecules and ions.
Lewis Structures
A Lewis structure is a representation of covalent bonding where shared electron pairs are shown as lines and lone electron pairs are shown as dots.
When drawing a Lewis structure, the octet rule is followed to attain the most stable electron configuration and achieve a complete octet of electrons for each atom in the molecule. All the valence electrons of the atoms in a Lewis structure must appear in the structure.
General Guidelines
- For the A-group elements, the number of valence electrons of an atom is equal to the group number.
For example, carbon, in Group 4A, has 4 valence electrons.
- The number of unpaired electrons on an atom of Groups 4A through 8A is 8 minus the group number. For example, oxygen in Group 6A has 8 - 6 = 2 unpaired electrons and forms 2 bonds. Carbon in Group 4A has 8 - 4 = 4 unpaired electrons and forms 4 bonds. (The number of unpaired electrons is equal to the group number for Groups 1A to 3A).
- Hydrogen always forms one bond.
- Oxygen usually doesn't bond with another oxygen atom. Exceptions: peroxides, superoxides, molecular oxygen, O 2 , and ozone, O 3.
- In general, only C, N, O, and S form multiple (double and triple) bonds.
- Usually, all the electrons in a Lewis structure are paired.
Writing Lewis Structures
- Count the total number of valence electrons in the molecule or ion. For anions, add an electron for every negative charge. For cations, subtract an electron for every positive charge.
- Write the skeletal structure of the compound. Generally, the least electronegative atom occupies the central position. Hydrogen atoms are always terminal atoms. (A central atom is bonded to two or more atoms. A terminal atom is bonded to just one other atom).
- Join the atoms in the structure by single bonds. For each single bond formed, subtract two electrons from the total number of valence electrons.
- With the valence electrons remaining, complete the octets of the terminal atoms (except H). Then complete the octets of the central atom(s).
a) If pairs are still left at this point, assign them to the central atom. If the central atom is from the third or higher period, it can accommodate more than 4 electron pairs.
b) If there are "too few electrons to go around", convert single bonds to multiple bonds. A double bond compensates for a deficiency of two electrons; a triple bond compensates for a deficiency of four electrons.
Example: Draw the Lewis structure for CH 2 O
There are a total of 12 valence electrons: 2 (1 for each H) + 4 (for C) + 6 (for O) = 12
Carbon is the central atom:
The 6 remaining electrons are placed around the O atom as lone pairs.
The C atom is deficient by one pair; therefore a double bond is used.
d) For neutral molecules, the sum of the formal charges must add up to zero. For cations, the sum of the formal charges must equal the positive charge. For anions, the sum of the formal charges must equal the negative charge.
e) Avoid the same formal charge on adjacent atoms.
Resonance
Resonance structures are used when a single Lewis structure doesn't adequately describe the bonding in the molecule. Resonance structures differ only in the position of the electrons.
For example, the resonance structures for ClO 4 -^ are:
One Lewis structure doesn't accurately portray the bonding in ClO 4 -. The true structure of ClO 4 - is a "hybrid" of the four resonance structures. The central chlorine atom is identically bonded to all the terminal oxygen atoms. The bonds are intermediate between a single and three double bonds.
Resonance forms do not imply different molecules. The molecule has a single structure; it does not oscillate back and forth between the Lewis structures. The true structure has an electron distribution that is a hybrid of all possible resonance structures.
Exceptions to the Octet Rule
- Incomplete Octets: In some compounds the number of electrons surrounding the central atom is fewer than eight.
For example: BeH 2 and BF 3
H Be^ H F B^ F
F
- Expanded Octets: In some compounds there are more than eight electrons surrounding the central atom. Expanded octets occur for atoms in and beyond the third period of the periodic table.
For example: SF 6
F F
F
F
F
F S
- Odd-electron molecules: Some molecules contain an odd number of electrons. Since we need an even number of electrons for complete pairing, the octet rule cannot be satisfied for all of the atoms in any of these molecules.
For example: nitric oxide, NO
N O
TABLE 1 Molecular Geometry as a Function of Electron Group Geometry: VSEPR
Molecule type
# Electron pairs on central atom
# Bonding pairs
# Lone pairs
Arrangement of electron pairs
Molecular shape
Description
AB 2 2 2 0 linear
AB 3 3 3 0
trigonal planar
AB 2 E 3 2 1 bent
AB 4 4 4 0 tetrahedral
AB 3 E 4 3 1
trigonal pyramidal
AB 2 E 2 4 2 2 bent
AB 5 5 5 0
trigonal bipyramidal
TABLE 1 ( continued) Molecular Geometry as a Function of Electron Group Geometry: VSEPR
Molecule type
# Electron pairs on central atom
# Bonding pairs
# Lone pairs
Arrangement of electron pairs
Molecular shape
Description
AB 4 E 5 4 1 unsymmetrical tetrahedron (seesaw)
AB 3 E 2 5 3 2 T-shaped
AB 2 E 3 5 2 3 linear
AB 6 6 6 0 octahedral
AB 5 E 6 5 1 square pyramidal
AB 4 E 2 6 4 2 square planar
TABLE 2 Hybrid Orbitals and their Geometric Orientation
PURE ATOMIC ORBITALS OF THE CENTRAL ATOM
HYBRIDIZATION OF THE CENTRAL ATOM
NUMBER OF HYBRID ORBITALS (# OF ELECTRON PAIRS ON CENTRAL ATOM)
GEOMETRIC ORIENTATION OF ELECTRON PAIRS
s,p (^) sp 2 linear
s,p,p (^) sp^2 3 trigonal planar
s,p,p,p (^) sp^3 4 tetrahedral
s,p,p,p,d (^) sp^3 d 5 trigonal bipyramidal
s,p,p,p,d,d (^) sp^3 d^2 6 octahedral
Polarity
A nonpolar bond is one in which the electron pair is shared equally between atoms of identical electronegativities. (For example, H – H and Br - Br).
A polar bond is one in which there is an unequal sharing of the electron pair between atoms of different electronegativities. The electrons are attracted to the more electronegative atom and a separation of charge occurs which generates a bond dipole (+ pole and - pole). The charge separation can be represented as:
Any molecule that has a net separation of charge has a dipole moment. The dipole moment μ, is a measure of the magnitude of the separated charges and the distance between them.
μ = Q x r where Q is the charge and r is the distance between charges.
A molecule that possesses a dipole moment is a polar molecule. A molecule that does not possess a dipole moment is a nonpolar molecule. The molecular dipole moment is the vector sum of all the individual dipoles in the molecule. Depending upon the molecular shape, the bond dipoles may add together (reinforce one another) to give a polar molecule, or they may cancel one another resulting in a nonpolar molecule. Therefore, the two criteria for determining the polarity of a molecule are bond polarity and molecular geometry.
Molecules having all nonpolar bonds are nonpolar molecules, there are no dipoles present.
Polar bonds in a molecule usually cause the molecule to be polar. However, polar bonds of equal magnitude can cancel one another if they are arranged symmetrically around the central atom, resulting in no net dipole moment.
