Properties of Atoms & Bonds

Atomic Structure

Atoms are the fundamental units of chemical elements, consisting of a nucleus made up of protons and neutrons, surrounded by electrons in an electron cloud. Understanding atomic structure is essential for predicting chemical behavior and bonding.

• Nucleus: Contains protons (positively charged) and neutrons (neutral). The number of protons determines the atomic number and identity of the element.

• Electron Cloud: Electrons are negatively charged and occupy regions around the nucleus called orbitals. The electron cloud is much larger and more diffuse than the nucleus.

• Atomic Number (Z): Number of protons in the nucleus; also equals the number of electrons in a neutral atom.

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

• Example: Carbon (C) has Z = 6, A ≈ 12. Atomic mass is approximately the sum of protons and neutrons.

Electron Density and Orbitals

Electron density describes the probability of finding an electron at a particular location around the nucleus. Orbitals are regions of space where electrons are most likely to be found, and each orbital can hold up to two electrons.

• Electron Density: Highest near the nucleus and decreases with distance. Represented by probability distributions.

• Orbitals: Types include s, p, d, and f. The 1s orbital is spherical, while p orbitals are dumbbell-shaped and oriented along x, y, and z axes.

• Nodes: Regions of zero electron density within an orbital.

• Each orbital holds:

• Example: The 1s orbital has no nodes; the 2s orbital has one node.

Electron Shells and the Periodic Table

As elements increase in atomic number, additional electron shells and orbitals are filled. The periodic table reflects the filling order of these shells.

• Electron Shells: Principal energy levels (n = 1, 2, 3, ...), each containing one or more types of orbitals.

• Periodic Table: Each row corresponds to the filling of a new shell of orbitals.

• Example: Second period elements fill the 2s and 2p orbitals.

Atomic Orbitals: s and p Orbitals

Atomic orbitals have distinct shapes and energy levels. The s orbital is spherical, while p orbitals are oriented along three axes and have higher energy than s orbitals in the same shell.

• 1s Orbital: Spherical, closest to nucleus, no nodes.

• 2s Orbital: Spherical, one node, higher energy than 1s.

• 2p Orbitals: Three degenerate (equal energy) orbitals: 2px, 2py, 2pz. Each has a node at the nucleus.

• Example: For neon (Ne), electron configuration is .

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. Electrons fill orbitals in order of increasing energy, following specific rules.

• Aufbau Principle: Electrons occupy the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers; each orbital holds two electrons with opposite spins.

• Hund's Rule: Electrons fill degenerate orbitals singly before pairing ("bus rule").

• Example: Nitrogen (N) electron configuration:

Bonding and Molecules

Chemical bonds form when atoms transfer or share electrons to achieve filled valence shells (octet rule). Bonds hold atoms together in molecules, which are the basis of organic chemistry.

• Ionic Bonds: Electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other.

• Covalent Bonds: Electrons are shared between atoms to fill octets. Most common in organic molecules.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Example: Formation of water () involves covalent bonding between hydrogen and oxygen.

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms and the distribution of valence electrons in a molecule. They are essential for visualizing molecular structure and predicting reactivity.

• Valence Electrons: Electrons in the outermost shell, involved in bonding.

• Drawing Lewis Structures:

  1. Determine the electron configuration of each atom.

  2. Count the number of valence electrons.

  3. Arrange atoms to satisfy the octet rule.

  4. Connect atoms with bonds (lines) and show lone pairs (dots).

• Example: Lewis structure for (dimethyl ether):

  • Each carbon forms four bonds; oxygen forms two bonds and has two lone pairs.

Multiple Bonds and Bonding Patterns

Atoms can form single, double, or triple bonds depending on the number of shared electron pairs. Common bonding patterns are observed for elements in organic molecules.

• Single Bond: Two electrons shared.

• Double Bond: Four electrons shared (two pairs).

• Triple Bond: Six electrons shared (three pairs).

• Common Patterns:

  • Carbon: 4 bonds

  • Nitrogen: 3 bonds, 1 lone pair

  • Oxygen: 2 bonds, 2 lone pairs

  • Halogens: 1 bond, 3 lone pairs

• Example: Ethylene () has a double bond between two carbon atoms.

Formal Charge

Formal charge is a bookkeeping tool to determine the distribution of electrons in a molecule and predict reactivity. It is calculated for each atom in a Lewis structure.

• Formula:

• Purpose: Helps identify sites of electron density and potential reactivity.

• Example: In , the formal charge on the oxygen atom is calculated using the formula above.

Summary Table: Common Bonding Patterns in Organic Molecules

Additional info: Some context and examples have been inferred and expanded for completeness and clarity, including the summary table and stepwise instructions for Lewis structures.