Atomic Electronic Structure

Subatomic Particles and Atomic Identity

Atoms are composed of protons, neutrons, and electrons. The arrangement and number of these particles determine the identity and properties of each element.

• Protons: Positively charged particles located in the nucleus. The number of protons defines the element (e.g., carbon has 6 protons).

• Neutrons: Neutral particles in the nucleus. Atoms of the same element with different numbers of neutrons are called isotopes (e.g., carbon-12 and carbon-13).

• Electrons: Negatively charged particles found outside the nucleus. The number of electrons determines the atom's charge.

Electron Configuration Example (Carbon):

• 6 protons, 6 electrons (neutral atom)

• Electron configuration:

• Core electrons:

• Valence electrons: (involved in bonding and chemical reactions)

Key Point: Reactions occur with valence electrons, which are the least stable and highest in energy.

Bonding and Electron Interactions

Types of Bonding

Atoms achieve stability by obtaining a full outer shell of electrons, either by transferring or sharing electrons.

• Ionic Bonds: Formed by the transfer of electrons from one atom to another.

• Covalent Bonds: Formed by the sharing of electrons between atoms. Most organic molecules are held together by covalent bonds.

Example: Hydrogen Bonding

• Hydrogen atom: 1 proton, 1 electron

• Losing an electron: (forms a proton, no valence electrons, stable)

• Gaining an electron: (forms hydride ion, full outer shell, stable)

• Sharing electrons: (each H is surrounded by 2 electrons, stable)

Key Point: Organic molecules are largely made of covalent bonds, focusing on orbital-orbital interactions.

Atomic Orbitals and Electron Waves

Wave Nature of Electrons

Electrons can be described as standing waves, leading to the concept of atomic orbitals.

• s-orbitals: Spherical, no nodes. Phases can be positive (upward displacement) or negative (downward displacement).

• p-orbitals: Dumbbell-shaped, with a node (region of zero electron density) at the nucleus.

Electron Wave Interference:

• Waves can combine constructively (additive, same phase) or destructively (cancel, opposite phase).

• Constructive interference leads to bonding; destructive interference leads to antibonding or no bond formation.

Molecular Orbitals and Bond Formation

Formation of Molecular Orbitals

When atomic orbitals from different atoms overlap, they combine to form molecular orbitals (MOs). The type of overlap determines the nature of the bond.

• Bonding Molecular Orbital (): Formed by constructive interference (same phase), resulting in increased electron density between nuclei and a stable bond.

• Antibonding Molecular Orbital (): Formed by destructive interference (opposite phase), resulting in a node between nuclei and decreased stability.

Energy Diagram for H2 Bond Formation:

• As two H atoms approach, their 1s orbitals overlap to form bonding and antibonding MOs.

• Bond length:

• Bond dissociation energy:

Conservation of Orbitals

• The number of molecular orbitals formed equals the number of atomic orbitals combined (e.g., 2 AOs → 2 MOs).

• If the antibonding MO is occupied, the bond is weakened or broken.

• Bonding MOs are lower in energy than the original atomic orbitals; antibonding MOs are higher in energy.

Types of Orbital Overlap and Bonding

Sigma (σ) Bonding

Sigma bonds are formed by end-on overlap of orbital density. For effective bonding, orbitals must be of the same phase.

• Possible combinations:

  • Two s-orbitals

  • s-orbital and p-orbital

  • Two p-orbitals (end-on)

• All single bonds in organic molecules are sigma bonds.

Pi (π) Bonding

Pi bonds are formed by side-on overlap of adjacent p-orbitals. They are present in double and triple bonds in organic molecules.

• Pi bonding involves the overlap of p-orbitals above and below the plane of the atoms.

• Pi bonds have nodal planes (regions of zero electron density) between the nuclei.

• Double bonds consist of one sigma and one pi bond; triple bonds consist of one sigma and two pi bonds.

Summary Table: Types of Bonding and Orbital Overlap

Key Concepts and Applications

• Valence electrons are responsible for chemical bonding and reactivity.

• Bond formation is driven by the stabilization achieved when atomic orbitals combine constructively.

• Molecular orbital theory explains the electronic structure and stability of molecules, especially in organic chemistry.

• Sigma and pi bonds are fundamental to understanding the structure and reactivity of organic molecules.