Chemical Bonding and Molecular Shapes

What is Organic Chemistry?

Organic chemistry is the study of molecules that are typically created and used by biological systems. Organic molecules are primarily composed of carbon and hydrogen, often containing other elements such as oxygen, nitrogen, sulfur, and halogens.

• Organic molecule: Any molecule that contains carbon.

• Hydrocarbons: Molecules containing only carbon and hydrogen.

• Example: Ethanol (C2H5OH) is an organic molecule.

Atomic Structure

Atoms are the basic units of matter, consisting of protons, neutrons, and electrons.

• Atomic number (Z): Number of protons in the nucleus.

• Mass number (A): Sum of protons and neutrons.

• Isotopes: Atoms with the same atomic number but different mass numbers.

• Ions: Atoms with unequal numbers of protons and electrons. Cations are positively charged, anions are negatively charged.

Example: Hydrogen isotopes: Protium (1H), Deuterium (2H), Tritium (3H).

Electron Configuration Principles

• Aufbau Principle: Electrons fill orbitals in order of increasing energy.

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

• Hund's Rule: Electrons fill degenerate orbitals singly before pairing.

Wave Function and Quantum Mechanics

Quantum mechanics describes electrons as both particles and waves. The probability of finding an electron in a region is given by the square of the wave function (ψ2).

• Heisenberg Uncertainty Principle: Cannot simultaneously know an electron's position and momentum.

• Atomic orbitals: Regions where electrons are likely to be found (s, p, d, f).

• Nodes: Regions where the probability of finding an electron is zero.

Molecular Orbital Theory

Atomic orbitals combine to form molecular orbitals, which can be bonding or antibonding.

• Bonding orbital: Constructive overlap, increased electron density between nuclei.

• Antibonding orbital: Destructive overlap, node between nuclei.

Example: H2 molecule forms a σ (sigma) bonding orbital.

Sigma (σ) and Pi (π) Bonds

Octet Rule

Atoms are most stable when they achieve a noble gas configuration (8 valence electrons).

• Atoms gain, lose, or share electrons to complete their octet.

• Exceptions: Hydrogen (2 electrons), Boron (6 electrons), expanded octets for elements in period 3 or higher.

Boding Preferences and Formal Charges

• Bonding preferences: Atoms form bonds to satisfy the octet rule.

• Formal charge:

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

Skeletal Structures

• Bond-line structures simplify organic molecules by omitting carbon and hydrogen labels.

• Carbons are at line ends and vertices; hydrogens on carbons are implied.

Lewis Structures

• Show all valence electrons and bonds.

• Follow octet rule and minimize formal charges.

• Multiple bonds and lone pairs are shown explicitly.

Condensed Structural Formulas

• Condensed formulas represent molecules in a compact form, showing connectivity without drawing all bonds.

Resonance Structures

• Resonance structures depict delocalization of electrons within a molecule.

• Curved arrows show electron movement.

• Resonance hybrid is the actual structure, a blend of all contributors.

Molecular Geometry (VSEPR Theory)

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Common geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Hybridization

• Atomic orbitals mix to form hybrid orbitals (sp, sp2, sp3).

• Number of electron domains determines hybridization:

Electronegativity and Bond Polarity

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

• Difference in electronegativity creates bond dipoles.

• Pauling scale is commonly used to compare electronegativities.

Bond polarity equation: where is the dipole moment, is the charge, and is the distance between charges.

Functional Groups

Functional groups are specific groups of atoms within molecules that determine chemical reactivity.

• Alkanes: Single bonds (C–C)

• Alkenes: Double bonds (C=C)

• Alkynes: Triple bonds (C≡C)

• Alcohols: –OH group

• Amines: –NH2, –NHR, –NR2

• Carbonyls: C=O group (includes aldehydes, ketones, carboxylic acids, esters, amides, etc.)

• Halides: C–X (X = F, Cl, Br, I)

• Others: Ethers, nitriles, thiols, sulfides, etc.

Example: Identify functional groups in a complex molecule by recognizing characteristic atoms and bonds.

Summary Table: Common Functional Groups

Additional info:

• Practice problems and diagrams are included throughout the notes to reinforce concepts.

• Tables and summary charts are provided for quick reference to bonding, hybridization, and functional group identification.