Organic Chemistry: Structure and Bonding (Chapter 1 Study Notes)

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Structure and Bonding

Atomic Structure

Atoms are the fundamental units of matter, consisting of a dense nucleus surrounded by an electron cloud. Understanding atomic structure is essential for grasping chemical bonding and molecular properties.

Nucleus: Contains positively charged protons and uncharged neutrons.
Electron Cloud: Composed of negatively charged electrons that move around the nucleus.
Atomic Number (Z): The number of protons in the nucleus; determines the element's identity.
Mass Number (A): The sum of protons and neutrons in the nucleus.
Isotopes: Atoms of the same element with different numbers of neutrons (thus different mass numbers).
Ions: Charged species formed by gaining or losing electrons. Cations are positively charged (fewer electrons), anions are negatively charged (more electrons).

Example: Carbon-12 has 6 protons and 6 neutrons (atomic number 6, mass number 12).

The Periodic Table

The periodic table organizes elements by increasing atomic number and groups them by similar chemical properties.

Elements in the same row (period) are similar in size.
Elements in the same column (group) have similar electronic and chemical properties.
Group number indicates the number of valence electrons for main group elements.

Example: Carbon is in the second row, group 4A.

Atomic Orbitals

Electrons occupy regions of space called atomic orbitals, which have characteristic shapes and energies.

s orbital: Spherical, lower in energy within a shell.
p orbital: Dumbbell-shaped, contains a node (region of zero electron density), higher in energy than s orbitals in the same shell.

Example: The 1s orbital can hold up to two electrons.

Electron Configuration and the Periodic Table

The arrangement of electrons in atomic orbitals determines the chemical properties of elements.

First row: Only the 1s orbital is available (H and He).
Second row: Four orbitals (one 2s and three 2p) can hold up to eight electrons.
Valence electrons are those in the outermost shell and are involved in bonding.

Bonding

Chemical bonding involves the joining of atoms to achieve stable electron configurations, often resembling those of noble gases.

Ionic Bonds: Formed by the transfer of electrons from one atom (typically a metal) to another (typically a nonmetal), resulting in oppositely charged ions.
Covalent Bonds: Formed by the sharing of electrons between atoms, usually nonmetals.

Example: NaCl forms via ionic bonding; H2 forms via covalent bonding.

Lewis Structures

Lewis structures are diagrams that represent the bonding and nonbonding electrons in molecules.

Each bond is a pair of shared electrons (represented by a line).
Lone pairs (nonbonded electrons) are shown as pairs of dots.
Second-row elements obey the octet rule (maximum of 8 electrons around each atom).
Hydrogen obeys the duet rule (maximum of 2 electrons).

Steps to Draw Lewis Structures:

Arrange atoms, placing hydrogen and halogens on the periphery.
Count total valence electrons (adjust for charges).
Connect atoms with single bonds, then fill octets with lone pairs.
If octets are incomplete, form multiple bonds as needed.
Assign formal charges to atoms as necessary.

Formal Charge

Formal charge helps identify the most stable Lewis structure and is calculated as:

Atoms "own" all their lone pair electrons and half of their bonding electrons.

Isomers

Isomers are compounds with the same molecular formula but different structures.

Constitutional isomers: Differ in the connectivity of their atoms.

Example: Ethanol and dimethyl ether (C2H6O) are constitutional isomers.

Exceptions to the Octet Rule

Some elements do not follow the octet rule:

Group 2A and 3A elements (e.g., Be, B) may have fewer than 8 electrons.
Third-row elements and beyond can have expanded octets.

Resonance

Some molecules cannot be represented by a single Lewis structure. Instead, resonance structures are used to depict delocalized electrons.

Resonance structures differ only in the arrangement of electrons, not atoms.
The actual molecule is a resonance hybrid, a weighted average of all valid resonance forms.
Curved arrow notation shows the movement of electron pairs between resonance forms.

Principles of Resonance:

Resonance structures are not real, isolated forms; only the hybrid is real.
Resonance structures are not in equilibrium and are not isomers.
Valid resonance structures must have the same number of electrons and valid Lewis structures.

Molecular Geometry and VSEPR Theory

The shape of a molecule is determined by the number of groups (atoms or lone pairs) around a central atom, as described by Valence Shell Electron Pair Repulsion (VSEPR) theory.

Number of Groups	Geometry	Bond Angle	Example
2	Linear	180°	HC≡CH
3	Trigonal planar	120°	CH2CH2
4	Tetrahedral	109.5°	CH4, NH3

Lone pairs occupy more space than bonding pairs, reducing bond angles (e.g., NH3 and H2O).

Condensed and Skeletal Structures

Organic molecules can be represented in several ways:

Condensed structures: Bonds are not always shown; groups are written together (e.g., CH3CH2OH).
Skeletal structures: Carbon atoms are implied at line ends and junctions; hydrogens on carbon are omitted for simplicity.
Heteroatoms (non-carbon/non-hydrogen atoms) and their hydrogens are always shown.

Hybridization and Bonding

Atomic orbitals mix to form hybrid orbitals that explain molecular shapes and bonding patterns.

sp3 hybridization: Four groups, tetrahedral geometry (e.g., CH4).
sp2 hybridization: Three groups, trigonal planar geometry (e.g., C in ethylene, CH2=CH2).
sp hybridization: Two groups, linear geometry (e.g., C in acetylene, HC≡CH).

Number of Groups	Number of Hybrid Orbitals	Type of Hybridization
2	2	sp
3	3	sp2
4	4	sp3

Bond Types:

σ (sigma) bond: Formed by head-on overlap of orbitals; all single bonds are sigma bonds.
π (pi) bond: Formed by side-on overlap of unhybridized p orbitals; present in double and triple bonds.

Example: Ethylene (CH2=CH2) has a C=C double bond (one σ and one π bond).

Bond Length and Strength

Bond length decreases and bond strength increases with the number of shared electrons between two nuclei.

Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.

Compound	C–C Bond Length (pm)	Bond Strength (kJ/mol)
Ethane (C–C)	154	368
Ethylene (C=C)	134	614
Acetylene (C≡C)	120	839

Electronegativity and Bond Polarity

Electronegativity is a measure of an atom's ability to attract electrons in a bond.

Electronegativity increases across a period and decreases down a group.
Nonpolar bonds: Electrons are shared equally (e.g., C–C, C–H).
Polar bonds: Electrons are shared unequally due to differences in electronegativity (e.g., C–O).
Bond dipole: Partial charges (δ+ and δ–) develop at each end of a polar bond.

Depicting Polarity: An arrow points toward the more electronegative atom, with a plus sign at the less electronegative end.

Polarity of Molecules

The overall polarity of a molecule depends on both the polarity of its bonds and its geometry.

Polar molecules: Have a net dipole moment (e.g., H2O).
Nonpolar molecules: Either have no polar bonds or their bond dipoles cancel due to symmetry (e.g., CO2).

Example: Water is polar due to its bent shape; carbon dioxide is nonpolar due to its linear shape.

Summary Table: Geometry Based on Number of Groups

Number of Groups	Geometry	Bond Angle	Example
2	Linear	180°	HC≡CH
3	Trigonal planar	120°	CH2CH2
4	Tetrahedral	109.5°	CH4, NH3

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