Comprehensive Notes on Lewis Structures and Molecular Geometry

Lewis Structures

Lewis Theory

Main-group elements strive for electron configurations resembling noble gases, adhering to the octet rule: they transfer or share electrons to achieve a complete outer shell of 8 electrons, mirroring noble gas configurations. Only valence electrons participate in bonding. Lewis Symbols represent atoms with their nucleus and core electrons symbolized, surrounded by dots indicating valence electrons. For example, Carbon (C) has the electron configuration [He]2s^22p^2, and Fluorine (F) has [He]2s^22p^5.

Drawing Lewis Structures

Count the total number of valence electrons in the structure. This is done by:
a) # valence e- = group # for main group elements
b) If the species is an anion, add e-
c) If species is a cation, subtract e-
Draw the skeletal structure.
a) Central atoms are generally the least electronegative element or the element with the highest positive oxidation state.
b) Carbon atoms are almost always central.
c) Hydrogen atoms are always terminal.
d) For oxyacids the hydrogens are always attached to the oxygens not the central atom (one hydrogen per oxygen)
Draw a bond everywhere it is needed (2 e^- per bond)
Fill the octets of terminal atoms by placing remaining electrons as lone pairs.
Place any remaining electrons on the central atom(s).
If the number of electrons available is insufficient to give each element an octet, double and triple bonds should be formed until each atom has an octet.
Calculate formal charge (check that the formal charges add up to the correct total charge).
Use multiple bonds to minimize formal charge separation if possible
Draw all resonance structures.

Lewis Structures

Bonding electrons: Electrons shared between two atoms.
Non-Bonding electrons: Electrons not shared, belonging to the atom they are on.

Formal Charge (FC)

A book-keeping method to determine the stability of a Lewis structure. The formula to determine formal charge is:

Formal Charge = # of valence e^- - # of e^- in lone pairs - # bonds

For each atom in the structure, calculate the formal charge. The sum of all FC must equal the overall charge of the molecule/ion.

Some example of drawing Lewis structure are:

He
HCl
CH_4
NH_3
H_2O
CH_3Cl
O_2
CO_2
N_2
CO
SI_2
CO_3^{2-}
H3PO4

Exceptions to the Octet Rule

Expanded octets: Elements with empty d orbitals, 3rd period or below, can have more than eight electrons (e.g., SO_2).
Odd number electron species
- Will have one unpaired electron (e.g., NO).
- Forms a Free-radical--Very reactive
Incomplete octets: (e.g., BH_3) for B, Al, Mg

Examples of exceptions include:

SO_2
BH_3
NO
PCl_5
NO_2
AlCl_3

Common Bonding Patterns

C = 4 bonds and 0 lone pairs
N = 3 bonds and 1 lone pair
O = 2 bonds and 2 lone pairs
H and halogen = 1 bond
Be = 2 bonds and 0 lone pairs
B = 3 bonds and 0 lone pairs

Structures deviating from these common patterns may have formal charges.

The most stable Lewis structure generally has:

The correct total number of valence electrons (modified by ion charge).
Each atom obeying the octet rule
- Exceptions: H, Be, B, Al & expanded octet
The least total formal charges (no formal charge larger than +1 or - 1).
The least charge separation (+ & - FC on adjacent atoms).
If FC exist, the negative FC should be on the most EN atom. The same sign FC on two adjacent atoms is extremely unlikely.

Significance of Lewis Structures

Lewis theory predicts the distribution of valence electrons in molecules, aiding in understanding bonding in many compounds. It allows prediction of molecular shapes and properties, and how they interact.

Resonance

Resonance occurs when more than one stable Lewis structure exists.

The optimal Lewis structure is an average of multiple structures, weighted most heavily towards those with the lowest energy (called the resonance hybrid). Examples include sulfur dioxide (SO2), carbonate ion (CO3^{2-}$), and nitrate ion (NO_3^{-}$).

Lewis structure Examples

CO: 10 electrons
CO_3^{2-}: 24 electrons
SO_2: 18 electrons
PBr_5
C2H2
POCl_3
NO_2^{+}
SO_4^{2-}
PO_4^{3-}
BeI_2: 16 electrons
NO_2: 18 electrons
AlI_3: 24 electrons

Molecular Geometry

Molecules are three-dimensional objects. The shape of a molecule is often described as geometric figures, which have characteristic "corners" indicating the positions of surrounding atoms around a central atom. These geometric figures also have characteristic angles called bond angles.

VSEPR Theory

VSEPR: Valence Shell Electron-Pair repulsion Theory states that Electron groups repel each other, assuming orientations about an atom to minimize repulsions.

Electron groups include bonding groups and lone pairs. any bond (single, double or triple) counts as a single group.
Electron-group geometry: the arrangement of electron groups (bonding and lone pairs) around a central atom.
Molecular geometry: the actual arrangement of atoms in space.

For example, in NO_2, there are three electron groups on N:

One lone pair
One single bond
One double bond

To determine molecular shape:

Draw the Lewis structure.
Determine the electron-group geometry.
- Each bond (single, double or triple) counts once.
- Each lone pair counts once.
Determine the molecular-group geometry.
- “Erase” lone pairs to get molecular shape.
- Lone pairs effectively decrease the bond angle below ideal.

Electron Group Geometry

The basic arrangements of electron groups around a central atom are:

Linear
Trigonal Planar
Tetrahedral

For molecules that exhibit resonance, the electron geometry will be the same regardless of which resonance form is used.

2 Electron Groups: Linear Electron Geometry: 2 electron groups results in the electron groups taking a linear geometry with a bond angle of 180°.3 Electron Groups: Trigonal Planar Electron Geometry: If all electron groups are bonding groups then the bond angle is 120° and the Molecular geometry / shape: trigonal planar
3 Electron Groups: Trigonal Planar Electron Geometry: If 2 electron groups are bonding groups and one is a lone pair then the bond angle is <120° and the molecular geometry / shape is bent
4 Electron Groups: Tetrahedral Electron Geometry: If all electron groups are bonding groups then the bond angle is 109.5° and the Molecular geometry = Td.
4 Electron Groups: Tetrahedral Electron Geometry:
1. If 3 electron groups are bonding groups and 1 is a lone pair then the bond angle is <109.5° the molecular geometry / shape: Trigonal pyramidal
2. If 2 electron groups are bonding groups and 2 are lone pairs then the bond angle is <109.5° the molecular geometry / shape: Bent

VSEPR Table

# e- groups	electron geometry	molecular geometry	bond angle
2	linear	2BG: linear	180^o
3	trigonal planar	3BG: trigonal planar	120^o
trigonal planar		2BG, 1LP: bent	<120^o
4	tetrahedral	4BG: Tetrahedral	109.5^o
tetrahedral		3BG, 1LP: trigonal pyramidal	<109.5^o
tetrahedral		2BG, 2LP: bent	<109.5^o

Lewis Structure Practice

Consider Nitrogen Trifluoride (NF3) and Dichlorine Monoxide (Cl2O):

For Nitrogen Trifluoride:

Number of VE: 26
Central atom: N
4 e- groups = 3 BG + 1 LP
electron geometry: Tetrahedral
Ideal bond angle: 109.5 degrees
Molecular geometry: Trigonal Pyramidal

For Dichlorine Monoxide:

Number of VE (valence electrons): 20
Central atom: O
4 Sets = 2 BA + 2 LP
Set geometry: Tetrahedral
Ideal bond angle: 109.5 degrees
Molecular geometry: Bent
Bond Angle < ideal

Polarity in Molecules

Using bond polarity and correct shapes, you can determine if a whole molecule is polar or nonpolar. Polarity of bonds and polarity of molecules are different. A molecule can have several polar bonds and still be non polar ex: CCl_4

Review: Polarity of Bonds

Boron trifluoride (BF3) is nonpolar because the polarity arrows cancel. Nitrogen trifluoride (NF3) is polar because the arrows don't cancel, resulting in an overall pull.

Notes on Polarity of Molecules

The size of the polarity arrow represents the relative difference in electronegativity. A B-F bond has a bigger polarity arrow than a B-Cl bond. When polarity arrows cancel, both size and direction are taken into account.

Some examples include:

H_2O
H_3O^{+}
CH3NH2
PCl_3
C2H4

We can also consider Formic acid, draw with correct geometry. Number of VE (valence electrons)__ Central atom__

4 e- groups = ____ BP + ____ LP
e- geometry around C__
bond angle
Molecular geometry around C__

or consider drawing the Lewis structure practice:

Number of VE (valence electrons)_
Central atom
____e- groups = ____ BP + ____ LP
e- geometry around O__
Ideal bond angle
Molecular geometry around O__

For CHCl3 and CCl4 what are the approx. Bond angles? Which molecule contains nonpolar covalent bonds? Which molecule contains polar covalent bonds? Which molecule is polar?

Draw the Lewis structure for Mg(ClO)_2 (Hint: what type of compound is this?)

Organic Molecules

Organic chemistry is the study of compounds containing C (often with H, O, N, S). Hydrocarbons only contain H and C. Alkanes are hydrocarbons with single bonds.

Examples of these are:

Methane (CH_4)
Ethane (C2H6)
Propane (C3H8)

Draw the Lewis dot structure for methanol (CH3OH) and ethanol (CH3CH_2OH) are they polar or non-polar? The order in which the atoms are written represents the order in which the atoms are bonded. H are never in the center.

Draw the lewis dot structure for acetic acid (CH_3COOH):

Number of VE (valence electrons)_
Central atom _
____e- groups = ____ BP + ____ LP
e- geometry around O__
Ideal bond angle
Molecular geometry around O__

True or False

Molecular geom. around an atom may differ from set geom.
A molecule w/linear geometry must have at least 1 dbl bond.
CO2 is linear, but SO2 is angular (bent).
A molecule w/polar covalent bonds must be polar.
H can only have 1 bond.
If the central atom bears 1 LP of e-s, the molecule is polar.
If the central atom bears 2 LP of e-s, the molecule is polar.
C never forms 5 bonds.
O can form 1, 2, or 3 bonds.
Halogens never form more than one bond.