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Molecular Shapes, Lewis Dot Structures, and Their Impact

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Shapes of Molecules

Introduction

Not all molecules are shaped the same way, and these differences in shape cause significant impacts in their chemical and physical properties. Understanding molecular shape is essential for predicting reactivity, polarity, and interactions with other molecules. Before discussing the impacts of shape, it is necessary to determine how to find the shape of molecules.

Lewis Dot Structures

Lewis Dot Structures Theory

A Lewis Dot Structure is a diagram that represents the valence electrons of atoms within a molecule. Valence electrons are the electrons in the outermost shell of an atom and are responsible for chemical bonding. In Lewis structures:

  • Bonding pairs (shared electrons) are shown as two dots (:) or a single line (-) between atoms.

  • Lone pairs (nonbonding electrons) are placed on the outside of atoms.

Lewis structures are not typically used for ionic compounds, but are essential for covalent compounds.

Valence Electrons and the Periodic Table

The number of valence electrons for each element can be determined from its group number on the Periodic Table:

Group Number

Number of Valence Electrons

1A (1)

1

2A (2)

2

3A (13)

3

4A (14)

4

5A (15)

5

6A (16)

6

7A (17)

7

8A (18)

8

These valence electrons are used to determine the Lewis Dot Structure of a molecule.

Drawing Lewis Dot Structures for Atoms

Each atom's valence electrons are represented as dots around the chemical symbol. For example:

  • Group 1A: 1 dot

  • Group 2A: 2 dots

  • Group 3A: 3 dots

  • Group 4A: 4 dots

  • Group 5A: 5 dots

  • Group 6A: 6 dots

  • Group 7A: 7 dots

  • Group 8A: 8 dots

For diatomic elements (e.g., F2), each atom shares electrons to complete its octet.

Bonding Patterns of Nonmetals

The number of bonds an atom can make is equal to the number of electrons needed to complete its octet. For example, nitrogen (Group 5A) has 5 valence electrons and needs 3 more to complete its octet, so it forms 3 bonds.

Element

Valence Electrons

Typical Number of Bonds

H

1

1

C

4

4

N

5

3

O

6

2

F, Cl, Br, I

7

1

Additional info: Hydrogen only requires two electrons to be complete (duet rule).

Calculating Valence Electrons in Molecules

To determine the total number of valence electrons in a molecule, add the valence electrons for each atom present:

  • Example: CO2 Carbon (Group 4A): 4 valence electrons Oxygen (Group 6A): 6 valence electrons × 2 = 12 Total: 4 + 12 = 16 valence electrons

  • Example: H2S Hydrogen (Group 1A): 1 valence electron × 2 = 2 Sulfur (Group 6A): 6 valence electrons Total: 2 + 6 = 8 valence electrons

Practice: Valence Electron Counting

Practice problems may ask for the total number of valence electrons in molecules such as N2O4, SF6, BF3, CCl4, C2H6. Add up the valence electrons for each atom to find the total.

Single, Double, and Triple Bonds

Each single bond contains two electrons, each double bond contains four electrons, and each triple bond contains six electrons.

  • Example: O2 Two oxygen atoms (Group 6A): 6 × 2 = 12 valence electrons Each oxygen needs 2 more electrons to complete its octet, so they share four electrons (double bond).

  • Example: N2 Two nitrogen atoms (Group 5A): 5 × 2 = 10 valence electrons Each nitrogen needs 3 more electrons, so they share six electrons (triple bond).

Steps for Drawing Lewis Dot Structures for Compounds

  1. Determine the arrangement of atoms (central atom is usually the least electronegative).

  2. Determine the total number of valence electrons.

  3. Attach each bonded atom to the central atom with a pair of electrons (single bond).

  4. Place the remaining electrons to complete octets (except for hydrogen, which only needs two electrons).

  • Example: CH4 Carbon in the center, four hydrogens attached. Total valence electrons: 4 (C) + 4 × 1 (H) = 8 Structure: H–C–H (with all hydrogens bonded to carbon)

  • Example: CO2 Carbon in the center, two oxygens attached. Total valence electrons: 4 (C) + 2 × 6 (O) = 16 Structure: O=C=O (with double bonds between C and each O)

Summary Table: Lewis Dot Structure Steps

Step

Description

1

Determine arrangement of atoms (central atom is least electronegative)

2

Calculate total number of valence electrons

3

Attach bonded atoms to central atom with single bonds

4

Distribute remaining electrons to complete octets (or duets for H)

Key Terms

  • Valence electrons: Electrons in the outermost shell of an atom, involved in bonding.

  • Lewis Dot Structure: Diagram showing valence electrons as dots around the chemical symbol.

  • Bonding pair: Shared pair of electrons between atoms (covalent bond).

  • Lone pair: Nonbonding pair of electrons on an atom.

  • Octet rule: Atoms tend to form bonds until they are surrounded by eight valence electrons (except hydrogen, which follows the duet rule).

Examples and Applications

  • Water (H2O): Oxygen in the center, two hydrogens attached, two lone pairs on oxygen.

  • Ammonia (NH3): Nitrogen in the center, three hydrogens attached, one lone pair on nitrogen.

  • Methane (CH4): Carbon in the center, four hydrogens attached, no lone pairs on carbon.

Additional info: These foundational concepts are essential for understanding molecular geometry, polarity, and chemical reactivity in later chapters.

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