A Review of General Chemistry: Electrons, Bonds, and Molecular Properties

A Review of General Chemistry: Electrons, Bonds, and Molecular Properties

Introduction to Organic Chemistry

Organic chemistry is the study of carbon-containing molecules and their reactions. It is a foundational discipline in chemistry, focusing on the structure, properties, and transformations of organic compounds, which are essential to life and industry.

• Organic compounds are found in food, clothes, plastics, pharmaceuticals, and all living organisms.

Atomic Structure

Atoms are composed of a nucleus (protons and neutrons) surrounded by electrons in orbitals. The electrons in the outermost shell are called valence electrons, which are primarily involved in chemical bonding.

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles residing in orbitals outside the nucleus.

Bond Formation

Atoms are connected by bonds. A covalent bond involves two atoms sharing a pair of electrons. The optimal bond length is determined by a balance of attractive and repulsive forces:

• Attractive forces: Between positively charged nuclei and negatively charged electrons.

• Repulsive forces: Between two positively charged nuclei or two negatively charged electrons.

The Structural Theory of Matter

Substances are defined by a specific arrangement of atoms. Compounds with the same molecular formula but different connectivity are called constitutional isomers.

• Example: C2H6O can be ethanol or dimethyl ether.

Drawing Constitutional Isomers

1. Determine the valency of each atom in the molecular formula.

2. Connect atoms of highest valency first.

3. Place monovalent atoms at the periphery.

4. Consider alternative connections for the atoms.

Counting Valence Electrons

Valence electrons can be determined from the periodic table:

Atom	Valence Electrons
C	4
H	1
O	6
N	5
Br	7

Lewis Dot Structure of an Atom

Lewis structures represent valence electrons as dots around the atomic symbol. Each side can have up to two dots, representing paired electrons.

• Example: Carbon (C) has four valence electrons, so four dots are placed around the symbol.

Octet Rule and Covalent Bonding

The octet rule states that main group atoms form bonds to achieve eight valence electrons in their outer shell. Second-row elements (C, N, O, F) generally obey this rule.

• Carbon: 4 bonds, no lone pairs

• Nitrogen: 3 bonds, 1 lone pair

• Oxygen: 2 bonds, 2 lone pairs

• Hydrogen: 1 bond, no lone pairs

• Halogens: 1 bond, 3 lone pairs

Common Bonding Patterns

• Carbon: Tetravalent (4 bonds)

• Nitrogen: Trivalent (3 bonds)

• Oxygen: Divalent (2 bonds)

• Hydrogen: Monovalent (1 bond)

• Halogens: Monovalent (1 bond)

Steps for Drawing Lewis Structures

1. Write the skeletal structure (central atom is usually less electronegative; hydrogen is always terminal).

2. Count all valence electrons (adjust for charges).

3. Distribute electrons: draw single bonds, complete octets for terminal atoms, place remaining electrons on the central atom.

4. If the central atom lacks an octet, form double or triple bonds as needed.

Identifying Formal Charges

Formal charge is calculated as:

$\text{Formal Charge} = \text{Valence} - \text{Lines} - \text{Dots}$

• Example: For oxygen in CH2O, $6 - 2 - 4 = 0$

Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons. It increases across a period and decreases down a group in the periodic table.

• Fluorine (F) is the most electronegative element.

Classification of Bonds

Electronegativity Difference	Type of Bond	Bonding Electrons
Less than 0.5	Nonpolar covalent	Equally shared
0.5 to 1.7	Polar covalent	Not shared equally
Greater than 1.7	Ionic	Not shared at all

Polar Covalent Bonds

• More electronegative atom attracts electrons more strongly, gaining a partial negative charge ($\delta^-$).

• Less electronegative atom has a partial positive charge ($\delta^+$).

• Inductive effect: Withdrawal of electrons toward the more electronegative atom.

Bond-Line Structures

Bond-line structures are simplified representations where each corner or endpoint represents a carbon atom. Hydrogen atoms bonded to carbon are not shown, but those attached to heteroatoms (N, O, halogens) must be shown.

• Double bonds: Two lines

• Triple bonds: Three lines (linear)

Atomic Orbitals

Atomic orbitals are regions of space where electrons are likely to be found. The main types are s, p, d, and f orbitals. The shape of an orbital is related to the region of space containing 90–95% of the electron density.

• s orbital: Spherical

• p orbital: Dumbbell-shaped

Phases of Atomic Orbitals

• Electrons behave as both particles and waves.

• The sign of the wave function (positive or negative) is important for orbital overlap in bonding but does not relate to electrical charge.

Filling Atomic Orbitals with Electrons

• Aufbau principle: Fill lowest energy orbitals first.

• Pauli exclusion principle: Each orbital holds two electrons with opposite spins.

• Hund's rule: One electron per degenerate orbital before pairing.

Example electron configurations:

• C: $1s^2 2s^2 2p^2$

• N: $1s^2 2s^2 2p^3$

• O: $1s^2 2s^2 2p^4$

Valence Bond Theory

A bond forms when atomic orbitals overlap. Only constructive interference (overlapping waves in phase) results in a bond, and only valence electrons participate in bonding.

Single, Double, and Triple Bonds

• Single bond: One pair of electrons ($-$)

• Double bond: Two pairs of electrons ($=$)

• Triple bond: Three pairs of electrons ($\equiv$)

Sigma and Pi Bonds

Bond Type	Number of $\sigma$-bonds	Number of $\pi$-bonds
Single	1	0
Double	1	1
Triple	1	2

• Sigma ($\sigma$) bond: Direct (head-to-head) overlap of orbitals, electron density on the bond axis.

• Pi ($\pi$) bond: Side-by-side overlap of parallel p orbitals, electron density above and below the bond axis.

Hybridized Atomic Orbitals

Hybridization explains the formation of equivalent bonds in molecules like methane (CH4), ethene (C2H4), and ethyne (C2H2).

• sp3 hybridization: Tetrahedral geometry, 109.5° bond angles (e.g., CH4).

• sp2 hybridization: Trigonal planar geometry, 120° bond angles (e.g., C2H4).

• sp hybridization: Linear geometry, 180° bond angles (e.g., C2H2).

Identifying Hybridization States

Steric Number (# of things attached)	Hybridization
2	sp
3	sp2
4	sp3

Hybridization and s Character

Orbital	Percentage s character	Percentage p character
s	100%	0%
sp	50%	50%
sp2	33%	67%
sp3	25%	75%
p	0%	100%

Additional info: These foundational concepts are essential for understanding molecular structure, reactivity, and the physical properties of organic compounds. Mastery of these topics is critical for success in organic chemistry and related fields.