Molecular Structure Study Guide (page 2)
Three-dimensional structures showing shape, geometry, and valance electrons provide a model for comprehending the arrangement of atoms in a molecule. Molecular geometries play an important part in the intramolecular and intermolecular properties of a substance.
Lewis structures are formulas for compounds in which each atom exhibits an octet (eight) of valence electrons. These representations are named after Gilbert N. Lewis for his discovery that atoms in a stable molecule want to achieve a noble gas configuration of eight valance electrons. These electrons are always paired and are represented as dots for nonbonded (lone) pairs or a line for every bonded (shared) pair of electrons. The rules for writing Lewis structures are as follows:
- Sum of all the valance electrons (which should be an even number). Remember: The number of valance electrons is the group number of the element.
- Form bonds between the atoms using pairs of electrons. Usually, the least electronegative element is the central element. Hydrogen is never the central element.
- Arrange the remaining electrons as lone pairs or create double or triple bonds to satisfy the octet rule. Exceptions: Hydrogen satisfies the duet (two) rule, and boron and aluminum satisfy the six-electron rule.
Write the Lewis structure for H2O, PCl3, BF3, and CO2.
H2O: Valance electrons: 2(1) + 6 = 8
PCl3: Valance electrons: 5 + 3(7) = 26
BF3: Valance electrons: 5 + 3(7) = 26
(Boron follows the six-electron rule.)
CO2: Valance electrons: 4 + 2(6) = 16
In trying to satisfy the octet and only have 16 valence electrons, the carbon was not fulfilled with only single bonds.
A pair of electrons from each oxygen was used to form a double bond to the carbon and satisfy the octet of all three atoms.
Another way to test to see if the octet rule is met is to write all the paired electrons as dots and circle each element. Each element circled should have eight electrons (two for hydrogen and six for boron and aluminum) and clearly show the bond overlap for the bonding pairs.
3-D: Valence Shell Electron Pair Repulsion (VSEPR) Theory
The VSEPR model is based on electrostatic repulsion among electron pair orbitals. By pushing each pair as far as possible, electron pairs dictate which geometry or shape a molecule will adopt. Molecules should be written as 2-D Lewis structures, and then determine the number of bonding pairs and nonbonding pairs. A summary of the shapes and possible arrangements can be found in Figure 11.1 and Table 11.1. Double and triple bonds can be treated as one bonding pair for VSEPR theory. Such bonds count as one bonding pair.
Using VSEPR, predict the shape for the following molecules or ions: KrF2, HCN, PCl3, NO2–, NO3–.
Resonance occurs when one or more valid Lewis structures exist for a molecule or polyatomic ion. The structures that represent the substance are called resonance structures. Each resonance structure does not characterize the substance, but the average of all the resonance structures represents the molecule or polyatomic ion. Resonance structures are usually placed in brackets and separated by a double-headed arrow (↔).
Show the resonance structures for nitrate, NO3–. Although the three nitrate resonance structures are written separately, nitrate is a combination of all three structures.
Atoms in certain molecules or polyatomic ions may have a formal charge. A formal charge is the difference in the number of valance electrons in the neutral atom (group number) and the number of electrons assigned to that atom in the molecule or polyatomic ion. Mathematically, the equation is
formal charge = group number – number of bonds – number of lone electrons
Give the formal charge for each atom in the nitrate ion, NO3–, and ozone, O3 (yes, even neutral molecules can have elements with formal charges).
All formal charges must add up to the charge on the polyatomic ion or zero for a neutral molecule.
The electronegativity of an element is its strength and ability to attract paired electrons in a covalent molecule. Electronegativity increases as you move up and over to the right on the periodic table. Fluorine is the most electronegative element. (See Figure 11.2)
A dipole results in a covalent bond between two atoms of different electronegativity. A partial positive (+) and a negative charge (–) develop at both ends of the bond, creating a dipole (i.e., two poles) oriented from the positive end to the negative end. The oxygen atom is more electronegative than hydrogen in water and the result is a dipole. Dipoles are represented by a line with a perpendicular line ( ).
A dipole moment will exist in a molecule if the resulting dipoles do not cancel based on their additive vectors. If the two dipoles of water are added, water has a dipole moment in the "up" direction.
Identify the dipole for each of the following bonds: B-F, Cl-I, N-H.
Predict whether each of the following molecules has a dipole moment: SCl2, CH2Cl2.
SCl2: Sulfur dichloride has a bent structure and a dipole:
CH2Cl2: If drawn (incorrectly!) planar, the molecule shows no dipole:
However, if drawn correct as tetrahedral, the dipole is
Practice problems for these concepts can be found at - Molecular Structure Practice Questions