Molecular Bonding, Lewis Structures, and Molecular Geometry in GOB Chemistry

Ch. 4: Molecular (Covalent) Bonding and Shapes of Molecules

Introduction to Electronic Structure

Chemistry is fundamentally governed by the behavior of electrons, which act as a type of subatomic 'glue' that holds atoms together in molecules. Understanding charges and electron arrangements is key to predicting chemical properties and reactivity.

• Electrons are negatively charged subatomic particles involved in chemical bonding.

• Electron structure determines how atoms bond and interact.

Molecular Formulas and Structural Representations

Molecules can be represented in several ways, each providing different information about atomic connectivity and structure.

• Molecular formulas show the types and numbers of atoms in a compound, but not how they are connected. Example: CH4, H2O

• Structural formulas use lines to show how atoms are connected. Example: H–O–H for water.

• Lewis dot structures show how atoms are connected and the arrangement of valence electrons (shared and unshared pairs).

• Condensed structures do not show bonds explicitly. Example: CH3CH2OH

Valence Electrons and Lewis Dot Structures

Valence electrons are the electrons in the outermost shell of an atom and are crucial for bonding. Lewis dot structures represent these electrons as dots around the chemical symbol.

• Valence electrons are shown as dots around the element symbol.

• The number of valence electrons corresponds to the group number for main group elements.

Group	Valence Electrons
1 (IA)	1
2 (IIA)	2
13 (IIIA)	3
14 (IVA)	4
15 (VA)	5
16 (VIA)	6
17 (VIIA)	7
18 (VIIIA)	8

Additional info: This table helps predict the number of dots to place around each element in a Lewis structure.

Drawing Lewis Dot Structures

Steps for Drawing Good Lewis Structures

Lewis structures are a visual way to represent the arrangement of electrons in a molecule. Follow these steps for accurate structures:

1. Count the total number of valence electrons for all atoms in the molecule.

2. Determine connectivity: usually, the least electronegative atom is central (except hydrogen).

3. Complete the octet for each atom (except hydrogen, which only needs 2 electrons).

4. Check to see if any atoms lack an octet; create double or triple bonds as needed.

5. Replace lone pairs with bonds to form multiple bonds if necessary.

6. Minimize formal charges by adjusting electron placement.

• Single bond: 2 electrons

• Double bond: 4 electrons

• Triple bond: 6 electrons

Octet rule: Most atoms (except H, He, B, Be) aim for 8 valence electrons.

Exceptions: Hydrogen (2 electrons), Helium (2), Beryllium (4), Boron (6). Expanded octets possible for elements in period 3 and below.

Formal Charge Calculation

Formal charge helps determine the most stable Lewis structure.

• Formal charge = (valence electrons) – (non-bonding electrons) – (1/2 bonding electrons)

Examples of Lewis Structures

• H2O: Oxygen has 2 lone pairs and forms two single bonds with hydrogen.

• NH3: Nitrogen has one lone pair and forms three single bonds with hydrogen.

• CS2: Carbon forms two double bonds with sulfur atoms.

Bonding Electrons vs. Non-Bonding Electrons

Definitions

• Bonding electrons: Electrons shared between atoms to form covalent bonds.

• Non-bonding electrons (lone pairs): Electrons not involved in bonding, remain on the atom.

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the shapes of molecules based on the repulsion between electron pairs around a central atom. Electron pairs (bonding and non-bonding) arrange themselves as far apart as possible in 3D space.

• Linear: 180°, two electron groups (e.g., CO2)

• Trigonal planar: 120°, three electron groups (e.g., BF3)

• Bent: < 120°, two bonds and one or more lone pairs (e.g., H2O)

• Tetrahedral: 109.5°, four electron groups (e.g., CH4)

• Trigonal pyramidal: ~107°, three bonds and one lone pair (e.g., NH3)

Deviations from ideal geometry occur due to lone pairs, which repel more strongly than bonding pairs.

Electronegativity and Polarity

Electronegativity (χ)

Electronegativity is a measure of how strongly an atom attracts electrons in a bond. Differences in electronegativity between atoms lead to bond polarity.

• High electronegativity: F, O, N (the "Big Three")

• Electronegativity increases across a period and decreases down a group.

Polar vs. Non-Polar Bonds and Molecules

Polarity depends on both the shape of the molecule and the difference in electronegativity between atoms.

• Polar covalent bond: Electrons are shared unequally, creating partial charges ($ext{δ}^+$ and $ext{δ}^-$).

• Non-polar covalent bond: Electrons are shared equally.

• Polar molecule: Has an overall dipole moment due to shape and bond polarity (e.g., H2O).

• Non-polar molecule: No overall dipole moment (e.g., CO2).

Bond Type	Electronegativity Difference	Example
Non-polar covalent	0	Cl2
Polar covalent	Small	HCl
Ionic	Large	NaCl

Additional info: Polarity affects solubility, reactivity, and physical properties.

Nomenclature of Molecular Compounds

Rules for Naming Binary Molecular Compounds

Molecular compounds are named using prefixes to indicate the number of each type of atom.

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-
9	nona-
10	deca-

• The first element is named first; the second element is named with an -ide ending.

• Prefixes are used for both elements, except mono- is usually omitted for the first element.

Example: CO2 is carbon dioxide, N2O is dinitrogen monoxide.

Summary Table: Key Concepts

Concept	Definition	Example/Application
Lewis Dot Structure	Diagram showing valence electrons as dots	H2O, NH3
VSEPR Theory	Predicts molecular shape based on electron pair repulsion	Linear, Bent, Tetrahedral
Electronegativity	Atom's ability to attract electrons	F > O > N
Polarity	Distribution of electrical charge in a molecule	H2O (polar), CO2 (non-polar)
Nomenclature	System for naming compounds	CO2: carbon dioxide

Additional info: Mastery of these concepts is essential for understanding chemical bonding, molecular structure, and reactivity in GOB Chemistry.