Lewis Structures and Molecular Shapes: Study Guide

Lewis Structures and Molecular Shapes

Introduction to Lewis Structures

Lewis structures, also known as electron dot formulas, are diagrammatic representations of the valence electrons in molecules and ions. They help predict the three-dimensional arrangement of atoms and electron groups using the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• Lewis Structure: Shows bonding and non-bonding (lone pair) electrons around atoms.

• VSEPR Theory: Predicts molecular shapes based on repulsion between electron groups.

Steps for Constructing Lewis Structures

Follow these steps to construct Lewis structures for molecules and ions:

1. Count total valence electrons: Add one electron for each negative charge, subtract one for each positive charge.

2. Choose the central atom: Place the least electronegative atom (except hydrogen) in the center.

3. Arrange outer atoms: Place all other atoms around the central atom.

4. Connect atoms: Use single covalent bonds to connect the central atom to each outer atom.

5. Complete octets: Fill the octet for outer atoms (hydrogen needs 2, beryllium 4, boron 6).

6. Place remaining electrons: Put any leftover electrons on the central atom.

7. Form multiple bonds: If octets are incomplete, create double or triple bonds as needed.

8. Resonance structures: If multiple possibilities exist, calculate formal charges to select the best structure.

9. Exceptions to the octet rule:

   • Hydrogen, beryllium, and boron do not complete octets.

   • Some nitrogen compounds may have odd numbers of electrons.

   • Elements in period 3 and below can expand their octet due to vacant d-orbitals.

Formal Charge Calculation

Formal charge helps determine the most stable Lewis structure:

• Formula: $\text{Formal charge} = \text{Valence electrons} - (\text{Non-bonding electrons} + \frac{1}{2} \text{Bonding electrons})$

• Structures with formal charges closest to zero are preferred.

• If charges are equal, negative charges should be on more electronegative atoms.

AXE Notation and Molecular Classification

AXE notation classifies molecules based on the number of bonding and non-bonding electron groups around the central atom:

• A: Central atom

• X: Bonding electron groups (atoms bonded to A)

• E: Non-bonding electron groups (lone pairs on A)

Example: For NH3, the notation is AX3E1 (3 bonding groups, 1 lone pair).

Table: Molecular Types, Geometries, and Bond Angles

The following table summarizes the relationship between AXE types, electron group geometry, molecular geometry, and bond angles:

Molecule/Ion Type	Total Electron Groups	Bonding Groups	Non-bonding Groups	Electron Group Geometry	Molecular Geometry	Bond Angle
AX2	2	2	0	Linear	Linear	180°
AX3	3	3	0	Trigonal planar	Trigonal planar	120°
AX2E1	3	2	1	Trigonal planar	Bent	~120°
AX4	4	4	0	Tetrahedral	Tetrahedral	109.5°
AX3E1	4	3	1	Tetrahedral	Trigonal pyramid	~107°
AX2E2	4	2	2	Tetrahedral	Bent	~104.5°
AX5	5	5	0	Trigonal bipyramidal	Trigonal bipyramidal	180°, 120°
AX4E1	5	4	1	Trigonal bipyramidal	See-saw	~90°, ~120°
AX3E2	5	3	2	Trigonal bipyramidal	T-shape	~90°
AX2E3	5	2	3	Trigonal bipyramidal	Linear	180°
AX6	6	6	0	Octahedral	Octahedral	90°
AX5E1	6	5	1	Octahedral	Square pyramid	~90°
AX4E2	6	4	2	Octahedral	Square planar	90°

Note: Lone pairs and double bonds compress bond angles compared to ideal geometries.

Examples of Lewis Structures and Molecular Shapes

Below are examples of molecules and ions for which Lewis structures, AXE types, and molecular geometries can be determined:

• HCN: Linear molecule (AX2), bond angle 180°, polar.

• BeF2: Linear (AX2), bond angle 180°, nonpolar.

• SO2: Bent (AX2E1), bond angle ~120°, polar.

• O3: Bent (AX2E1), bond angle ~117°, polar.

• BF3: Trigonal planar (AX3), bond angle 120°, nonpolar.

• CO32−: Trigonal planar (AX3), bond angle 120°, nonpolar.

• CH4: Tetrahedral (AX4), bond angle 109.5°, nonpolar.

• ClO4−: Tetrahedral (AX4), bond angle 109.5°, nonpolar.

• PF3: Trigonal pyramid (AX3E1), bond angle ~107°, polar.

• H2O: Bent (AX2E2), bond angle ~104.5°, polar.

• PF5: Trigonal bipyramidal (AX5), bond angles 90°, 120°, 180°, nonpolar.

• SF4: See-saw (AX4E1), bond angles ~90°, ~120°, polar.

• BrF3: T-shape (AX3E2), bond angle ~90°, polar.

• XeF2: Linear (AX2E3), bond angle 180°, nonpolar.

• IOF5: Octahedral (AX6), bond angle 90°, polar.

• TeF5−: Square pyramid (AX5E1), bond angle ~90°, polar.

• XeF4: Square planar (AX4E2), bond angle 90°, nonpolar.

• XeO2F2: Square planar (AX4E2), bond angle 90°, polar.

• SO42−: Tetrahedral (AX4), bond angle 109.5°, nonpolar.

Example: For H2O:

• Lewis structure: Oxygen in center, two hydrogens bonded, two lone pairs on oxygen.

• AXE type: AX2E2

• Electron group geometry: Tetrahedral

• Molecular geometry: Bent

• Bond angle: ~104.5°

• Polarity: Polar

Experimental Design and Model Construction

To visualize molecular shapes, use a molecular model kit to construct three-dimensional models based on Lewis structures and AXE types. This helps understand spatial arrangement and bond angles.

1. Draw the Lewis structure for each molecule or ion.

2. Assign the AXE type.

3. Determine electron group and molecular geometry using the table above.

4. Construct the model using a kit.

5. Indicate bond angles and polarity for each structure.

Summary Table: Properties of Selected Molecules and Ions

Formula	Lewis Structure	AXE Type	Electron Group Geometry	Molecular Geometry	Bond Angle	Polarity
HCN	H–C≡N	AX2	Linear	Linear	180°	Polar
BeF2	F–Be–F	AX2	Linear	Linear	180°	Nonpolar
SO2	O=S=O (with lone pair on S)	AX2E1	Trigonal planar	Bent	~120°	Polar
O3	O–O=O (with lone pairs)	AX2E1	Trigonal planar	Bent	~117°	Polar
BF3	F–B–F (three F atoms)	AX3	Trigonal planar	Trigonal planar	120°	Nonpolar
CH4	H–C–H (four H atoms)	AX4	Tetrahedral	Tetrahedral	109.5°	Nonpolar
H2O	H–O–H (two lone pairs on O)	AX2E2	Tetrahedral	Bent	~104.5°	Polar
PF5	P with five F atoms	AX5	Trigonal bipyramidal	Trigonal bipyramidal	90°, 120°, 180°	Nonpolar
SF4	S with four F atoms, one lone pair	AX4E1	Trigonal bipyramidal	See-saw	~90°, ~120°	Polar
XeF4	Xe with four F atoms, two lone pairs	AX4E2	Octahedral	Square planar	90°	Nonpolar
SO42−	S with four O atoms	AX4	Tetrahedral	Tetrahedral	109.5°	Nonpolar

Additional info: Bond angles may be compressed in structures with lone pairs or double bonds. The polarity of a molecule depends on both its shape and the distribution of electronegative atoms.