Atomic Structure and Elements

Subatomic Particles and Isotopes

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons. The number of protons defines the element, while neutrons can vary, resulting in different isotopes of the same element. Electrons occupy energy levels (shells) around the nucleus.

• Proton (p+): Positively charged, found in the nucleus.

• Neutron (n): Neutral, found in the nucleus.

• Electron (e-): Negatively charged, found in orbitals around the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Example: Chlorine has two common isotopes: and .

Atomic Number and Mass Number

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

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

• Electron Count: In a neutral atom, electrons = protons.

Quantum Mechanics and Atomic Orbitals

Quantum Numbers and Electron Configuration

Quantum mechanics describes the behavior of electrons in atoms using mathematical equations. Electrons occupy atomic orbitals, which are regions of space with a high probability of finding an electron.

• Principal Quantum Number (n): Indicates energy level (shell).

• Orbital Types: s, p, d, f (shapes and energies differ).

• Pauli Exclusion Principle: No more than two electrons can occupy the same orbital, and they must have opposite spins.

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

Example: Electron configuration of carbon:

Heisenberg Uncertainty Principle

The Heisenberg Uncertainty Principle states that it is impossible to simultaneously determine both the exact position and momentum of an electron.

• Mathematical Expression:

Octet Rule and Chemical Bonding

Octet Rule

Atoms tend to gain, lose, or share electrons to achieve a full valence shell of 8 electrons, known as the octet rule. This drives the formation of chemical bonds.

• Ionic Bonds: Electrons are transferred from one atom to another (e.g., NaCl).

• Covalent Bonds: Electrons are shared between atoms (e.g., H2O, CH4).

Example: In water (H2O), oxygen shares electrons with two hydrogens to complete its octet.

Lewis Structures and Skeletal Structures

Drawing Lewis Structures

Lewis structures represent the arrangement of electrons in molecules, showing bonds and lone pairs.

1. Determine total number of valence electrons.

2. Arrange atoms and connect with single bonds.

3. Distribute remaining electrons to complete octets.

4. Assign formal charges as needed.

Example: Lewis structure of methane (CH4):

• Carbon in center, four hydrogens bonded to carbon.

• All atoms have complete valence shells.

Skeletal Structures

Skeletal structures are simplified representations, omitting hydrogen atoms bonded to carbon and showing bonds as lines.

• Condensed Structure: CH3CH2OH

• Skeletal Structure: Lines represent carbon-carbon bonds; vertices represent carbon atoms.

Bonding in Organic Molecules & Molecular Orbitals

Atomic and Molecular Orbitals

Atomic orbitals (s, p) combine to form molecular orbitals when atoms bond. The combination can be constructive (bonding) or destructive (antibonding).

• Bonding Molecular Orbital: Lower energy, increased electron density between nuclei.

• Antibonding Molecular Orbital: Higher energy, decreased electron density between nuclei.

Types of Covalent Bonds

• σ (sigma) bond: Formed by head-on overlap of orbitals (single bonds).

• π (pi) bond: Formed by side-on overlap of p orbitals (in double/triple bonds).

Example: Ethene (C2H4) has a σ bond and a π bond between the two carbons.

Hybridization

Atomic orbitals can mix to form hybrid orbitals that explain molecular geometry.

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

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

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

Electronegativity and Polarity

Electronegativity

Electronegativity is the tendency of an atom to attract electrons in a chemical bond. Differences in electronegativity lead to bond polarity.

• Polar Covalent Bond: Unequal sharing of electrons (e.g., H–F).

• Nonpolar Covalent Bond: Equal sharing of electrons (e.g., H–H).

Dipole Moments

A dipole moment arises in molecules with polar bonds and asymmetric geometry.

• Mathematical Expression:

Valence Electrons and the Periodic Table

Counting Valence Electrons

Valence electrons are the electrons in the outermost shell, crucial for bonding and chemical reactivity.

• Group number in the periodic table often equals the number of valence electrons for main group elements.

• Example: Carbon (Group 14) has 4 valence electrons.

Periodic Trends

• Electronegativity: Increases across a period, decreases down a group.

• Atomic Radius: Decreases across a period, increases down a group.

HTML Table: Comparison of Bond Types

HTML Table: Hybridization and Geometry

Additional info:

• Some content inferred from context and standard organic chemistry curriculum, such as the explanation of molecular orbitals and hybridization.

• Visuals referenced in the notes (e.g., orbital shapes, Lewis structures) are described in text for clarity.