Covalent Bonding I: Basic Concepts – Study Notes

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Covalent Bonding I: Basic Concepts

Rationale for Molecule and Compound Formation

Atoms combine to achieve a more stable electron configuration, as described by Gilbert Lewis. Maximum stability is achieved when an atom is isoelectronic with a noble gas, meaning it has the same electron configuration as a noble gas.

Isoelectronic: Having the same number of electrons as another species, typically a noble gas.
Atoms form molecules and compounds to reach this stable state.

Valence Electrons & Bonding

Only valence electrons (electrons in the outermost shell) are involved in chemical bonding. When atoms interact, valence electrons may be transferred or shared, and atoms tend to obey the Lewis Octet Rule.

Lewis Octet Rule: Atoms gain, lose, or share electrons to achieve eight electrons in their outer shell (an octet), similar to noble gases.
Exceptions to the octet rule exist (see below).

Lewis Structures of Atoms

Lewis structures use the element symbol and dots to represent valence electrons. This notation helps visualize bonding and electron arrangement.

Example: Na (1 dot), S (6 dots), C (4 dots), Mg (2 dots).

Lewis Dot Symbols Table

Element	Valence Electrons	Lewis Dot Symbol
Na	1	Na•
S	6	S••••••
C	4	C••••
Mg	2	Mg••
Cl	7	Cl•••••••
O	6	O••••••

The Ionic Bond

Definition and Formation

An ionic bond is the electrostatic force that holds ions together in an ionic compound. It forms when there is a large difference in electronegativities between atoms, resulting in electron transfer.

Examples: LiF and Na2O are ionic compounds.
Ionic compounds consist of formula units and are typically high-melting solids.

Representing Ion Formation

Electron configurations: Show how electrons are distributed before and after ion formation.
Orbital diagrams: Visualize electron transfer between atoms.
Lewis electron-dot symbols: Indicate the resulting ions and their electron arrangements.

Sample Problem: Depicting Ion Formation

Use orbital diagrams and Lewis symbols to show the formation of Na+ and O2− ions, then determine the formula of the compound (Na2O).

Other Examples

K and Br: K → K+ + e−; Br + e− → Br−; K+ + Br− → KBr
Ca and N: 3Ca → 3Ca2+ + 6e−; 2N + 6e− → 2N3−; 3Ca2+ + 2N3− → Ca3N2
Na and N: 3Na → 3Na+ + 3e−; N + 3e− → N3−; 3Na+ + N3− → Na3N

The Covalent Bond

Definition and Types

A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Covalent bonding occurs when the difference in electronegativities is too small for electron transfer.

Single covalent bond: Two atoms share one pair of electrons (e.g., H2O, F2).
Double bond: Two atoms share two pairs of electrons (e.g., O2, CO2).
Triple bond: Two atoms share three pairs of electrons (e.g., N2).

Bond Lengths and Strengths

Bond length is the distance between the nuclei of two covalently bonded atoms. Multiple bonds are shorter and stronger than single bonds.

Bond Type	Bond Length (pm)
C–C (single)	154
C=C (double)	133
C≡C (triple)	120
C–N (single)	143
C≡N (triple)	116

Bond order: Triple bond < Double bond < Single bond (in terms of length).
Bond strength increases as bond order increases.

Sample Problem: Comparing Bond Length and Strength

Given C=O, C–O, C≡O, rank by decreasing bond length and strength:

Bond length: C–O > C=O > C≡O
Bond strength: C≡O > C=O > C–O

Electronegativity

Definition and Trends

Electronegativity is the ability of an atom to attract electrons in a chemical bond. Fluorine (F) is the most electronegative element.

Electronegativity increases across a period and decreases down a group.

Electronegativities of Common Elements Table

Element	Electronegativity (Trend)
F	Highest
O	High
N	High
Cl	High
Cs	Lowest

Polar Covalent Bonds

A polar covalent bond has greater electron density around one atom, resulting in unequal sharing of electrons. This creates partial charges (δ+ and δ−).

Example: H–F bond, where F is electron-rich and H is electron-poor.

Classification of Bonds by Electronegativity Difference

Difference	Bond Type
< 0.4	Nonpolar Covalent
≥ 2.0	Ionic
0.4 < Difference < 2.0	Polar Covalent

Nonpolar covalent: Electrons shared equally.
Polar covalent: Electrons shared unequally.
Ionic: Electrons transferred.

Writing Lewis Structures

Steps for Drawing Lewis Structures

Lewis structures show how valence electrons are arranged among atoms in a molecule.

Count all valence electrons.
Determine the central atom (usually the least electronegative).
Draw single bonds to the central atom.
Place remaining electrons as lone pairs.
Convert lone pairs to double or triple bonds as needed to satisfy the octet rule.

Hydrogen only wants 2 electrons (duet rule).
Shared electrons are counted as owned by both atoms.

Bond Rules Table

Element	Typical Number of Covalent Bonds
Hydrogen (and Halogens)	1
Oxygen (and Sulfur)	2
Nitrogen (and Phosphorus)	3
Carbon (and Silicon)	4

Examples

NF3: Nitrogen forms three single bonds with fluorine atoms.
NH4+: Nitrogen forms four single bonds with hydrogen atoms (ammonium ion).
ClO4−: Chlorine forms four single bonds with oxygen atoms (perchlorate ion).

Procedures for Lewis Structures

Determine electrons needed (N): 8 for most atoms, 2 for H.
Determine electrons available (A): Count all valence electrons.
Electrons shared (S) = N – A.
Number of bonds = S / 2.
Use remaining electrons to complete octets.

Formal Charge and Lewis Structure

Definition and Calculation

Formal charge is the difference between the number of valence electrons in an isolated atom and the number assigned to that atom in a Lewis structure.

Formula:
The sum of formal charges in a molecule or ion must equal the overall charge.

Guidelines for Formal Charges

Structures with no formal charges are preferred for neutral molecules.
Structures with large formal charges are less plausible.
Negative formal charges should be placed on more electronegative atoms.

Example: CH2O (Formaldehyde)

Calculate formal charges for each atom using the formula above.
Choose the structure with formal charges closest to zero.

The Concept of Resonance

Definition and Examples

Resonance structures are two or more valid Lewis structures for a molecule that cannot be represented accurately by only one structure. The true structure is a hybrid of all resonance forms.

Example: Carbonate ion (CO32−) has three resonance structures.
Example: N2O has multiple resonance forms with different formal charges.

Exceptions to the Octet Rule

Types of Exceptions

Incomplete Octet: Some atoms (e.g., B, Be) can have fewer than 8 electrons (e.g., BeH2, BF3).
Odd-Electron Molecules: Molecules with an odd number of electrons (e.g., NO).
Expanded Octet: Atoms in period 3 or higher (e.g., S, P, Xe) can have more than 8 electrons (e.g., SF6, PF5, XeF4).

Examples Table

Type	Example
Incomplete Octet	BeH2, BF3
Odd-Electron	NO
Expanded Octet	SF6, PF5, XeF4

Summary: Covalent bonding involves the sharing of electrons to achieve stable electron configurations, often following the octet rule. Lewis structures, formal charges, resonance, and exceptions to the octet rule are essential concepts for understanding molecular structure and stability in general chemistry.