Atomic Structure, Bonding, and Electronegativity in Organic Chemistry

Atomic Structure and Important Elements

Key Elements in Organic Chemistry

Organic chemistry primarily involves a select group of elements, often referred to as the "big 4": carbon (C), hydrogen (H), oxygen (O), and nitrogen (N). Other important elements include halogens (group 7: F, Cl, Br, I) and two common drivers of reactivity: phosphorus (P) and sulfur (S).

• Carbon: Forms four bonds (tetravalent), allowing for diverse molecular structures.

• Hydrogen: Forms one bond.

• Oxygen: Typically forms two bonds.

• Nitrogen: Typically forms three bonds.

• Halogens: Usually form one bond.

Periodic Trends: Elements are organized by groups and periods, with group number often corresponding to the number of valence electrons.

Atomic Structure

Subatomic Particles

• Protons: Positively charged particles located in the nucleus; determine atomic number and identity of the element.

• Neutrons: Neutral particles in the nucleus; contribute to atomic mass but not charge.

• Electrons: Negatively charged particles in orbitals around the nucleus; involved in chemical bonding and reactions.

The number of protons equals the number of electrons in a neutral atom. Electrons in the outermost shell (valence electrons) participate in chemical bonding.

Electronic Structure

Orbitals and Electron Configuration

Electrons occupy regions of space called orbitals, which are grouped into shells at increasing distances from the nucleus. Each shell is identified by a principal quantum number (n = 1, 2, 3, ...).

• Each orbital can hold up to 2 electrons.

• Organic elements are found in the first two rows of the periodic table (n = 1 and n = 2).

• Electrons fill the lowest energy orbitals first (Aufbau Principle).

• Orbitals of equal energy are singly occupied before pairing (Hund's Rule).

• No two electrons in an atom can have the same set of quantum numbers (Pauli Exclusion Principle).

Example Electron Configuration:

• Carbon (C, atomic number 6): $1s^2\ 2s^2\ 2p^2$

• Nitrogen (N, atomic number 7): $1s^2\ 2s^2\ 2p^3$

Bond Formation and the Octet Rule

Octet Rule

Atoms tend to form bonds to achieve a stable configuration with 8 electrons in their valence shell (the octet rule), similar to noble gases.

• Atoms will transfer or share electrons to satisfy the octet rule.

• Exceptions: Hydrogen (2 electrons), Beryllium (4 electrons), Boron (6 electrons), and elements in period 3 or higher (can have expanded octets).

Lewis Structures

Drawing Lewis Structures

Lewis structures represent the bonding in covalent molecules, showing shared electron pairs as lines (bonds) and lone pairs as dots.

• Bonding electrons: Shared between atoms (lines).

• Nonbonding electrons: Lone pairs (dots).

• Count valence electrons for each atom and distribute to satisfy the octet rule (except for hydrogen).

Example: For water (H2O):

• Oxygen: 6 valence electrons

• Each hydrogen: 1 valence electron

• Lewis structure: H–O–H with two lone pairs on O

Charges: For every negative charge, add one electron; for every positive charge, subtract one electron.

Valence, Bonds, and Lone Pairs

Typical Valence and Bonding Patterns

Element	Valence	Bonds	Lone Pairs
C	4	4	0
N	5	3	1
O	6	2	2
F	7	1	3

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is an atom's ability to attract electrons in a chemical bond. The Pauling scale is commonly used to compare electronegativities.

• Increases from left to right across a period.

• Decreases down a group.

• Most electronegative element: Fluorine (F).

Bond Polarity:

• Nonpolar covalent bond: Electrons shared equally (e.g., H–H).

• Polar covalent bond: Electrons shared unequally, creating dipole moments (e.g., H–Cl).

• Ionic bond: Electrons transferred (e.g., Na+Cl–).

Dipole moment arises from differences in electronegativity, resulting in partial charges on atoms.

Formal Charges

Calculating Formal Charge

Formal charge helps keep track of electron distribution in molecules.

• Formula: Formal charge = (valence electrons) – (bonds) – (nonbonding electrons)

• Sum of formal charges in a molecule equals the overall charge.

Example: For the cyanide ion (CN–):

• Carbon: 4 valence electrons – 3 bonds – 2 nonbonding electrons = –1

• Nitrogen: 5 valence electrons – 3 bonds – 2 nonbonding electrons = 0

Summary Table: Electronegativity Values (Pauling Scale)

Element	Electronegativity
H	2.1
C	2.5
N	3.0
O	3.5
F	4.0
Cl	3.0
Br	2.8
I	2.5

Additional info: These notes provide foundational concepts for understanding atomic structure, bonding, and molecular properties in organic chemistry. Mastery of these topics is essential for predicting molecular behavior and reactivity.