Atoms and Ions: Atomic Structure, Orbitals, and Ion Formation

Atoms and Ions

Introduction

This study guide covers the foundational concepts of atomic structure, atomic orbitals, and ion formation, which are essential for understanding the chemical behavior of elements and compounds in organic chemistry. Mastery of these topics is crucial for interpreting molecular structure, reactivity, and the periodic trends that influence organic reactions.

Atomic Structure

Subatomic Particles

• Protons: Positively charged particles with a mass of approximately $1.67 \times 10^{-27}$ kg.

• Neutrons: Neutral particles with a mass similar to protons ($1.67 \times 10^{-27}$ kg).

• Electrons: Negatively charged particles with a much smaller mass ($9.11 \times 10^{-31}$ kg).

The nucleus contains protons and neutrons, while electrons occupy regions of space around the nucleus called shells or orbitals.

Atomic Number and Atomic Mass

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

• Atomic Mass (A): Sum of protons and neutrons in the nucleus.

Example: Carbon has Z = 6 and A ≈ 12.

Isotopes & Relative Atomic Mass

Isotopes

• Isotopes: Atoms of the same element (same number of protons) but different numbers of neutrons.

• Example: Carbon has isotopes $^{12}C$, $^{13}C$, and $^{14}C$.

Relative Atomic Mass

The relative atomic mass ($A_r$) is a weighted average based on the natural abundance of isotopes:

• $A_r = \sum (\text{isotope mass} \times \text{fractional abundance})$

For carbon:

• $A_r = 12 \times 0.989 + 13 \times 0.011 = 12.011$

The Chemical Elements

Definition and Examples

• Element: A pure substance that cannot be broken down by chemical means; consists of atoms with the same atomic number.

• Examples: Carbon (diamond, graphite, fullerenes), Oxygen (O2, O3).

Compounds are substances made from more than one element, e.g., CO2, H2O.

Atomic Structure: Quantum Model

Electron Shells and Energy Levels

• Electrons occupy shells defined by principal quantum numbers (n = 1, 2, 3...).

• The valence shell is the outermost shell containing electrons; it determines chemical reactivity.

Valence Shell, External Configuration & Atomic Orbitals

Valence Electrons

• Electrons in the valence shell are called valence electrons.

• Valence electrons are responsible for bonding and chemical properties.

Example: Carbon's electron configuration is $1s^2 2s^2 2p^2$; its valence shell is n = 2.

Atomic Subshells

Types and Shapes of Orbitals

• s-orbitals: Spherical shape

• p-orbitals: Dumbbell shape

• d-orbitals: Four-leaf clover shape

• f-orbitals: Complex shapes (e.g., raspberry-like)

Each shell contains one or more subshells, which are groups of equivalent orbitals.

Orbitals & Energy Levels

Filling Order and Degeneracy

• Orbitals are filled in order of increasing energy (Aufbau principle).

• Degenerate orbitals have the same energy within a subshell.

Order of filling: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d$ ...

Orbitals Energy & The Periodic Table

Periodic Table Structure

• The arrangement of elements in the periodic table reflects the order in which orbitals are filled.

• Elements are grouped into blocks (s, p, d, f) based on their valence electron orbitals.

Electron Excitation & Relaxation

Excitation

• When an electron absorbs energy, it can move to a higher energy level (excited state).

Relaxation

• When the electron returns to its ground state, it emits energy as electromagnetic radiation.

Example: Sodium emission spectrum shows characteristic lines at 589.995 nm and 589.5924 nm.

Electromagnetic Spectrum

Wavelength, Frequency, and Energy

• Electromagnetic radiation is characterized by wavelength ($\lambda$) and frequency ($\nu$).

• Energy of a photon: $E = h\nu$

• Relationship: $c = \lambda \nu$

Where $h$ is Planck's constant and $c$ is the speed of light.

Bohr's Frequency Condition

• The frequency of a photon absorbed or emitted during an electronic transition is proportional to the energy difference between states.

Ionisation Energy

Definition and Trends

• Ionisation energy: Energy required to remove an electron from a gaseous atom or ion.

• Measured in joules (J) or electron volts (eV).

• Equation: $\text{Na}(g) \rightarrow \text{Na}^+(g) + e^-$

• Trends:

  • Decreases from top to bottom in groups.

  • Increases from left to right in periods.

Ion Formation

Cations and Anions

• Cations: Positively charged ions formed by loss of electrons.

• Anions: Negatively charged ions formed by gain of electrons.

• Atoms are neutral when the number of protons equals the number of electrons.

Example: Na loses one electron to form Na+; Cl gains one electron to form Cl-.

Summary Table: Atomic Particles and Properties

Particle	Charge	Mass (kg)	Location
Proton	+1	1.67 × 10-27	Nucleus
Neutron	0	1.67 × 10-27	Nucleus
Electron	-1	9.11 × 10-31	Electron cloud

Key Equations

• $E = h\nu$

• $c = \lambda \nu$

• $A_r = \sum (\text{isotope mass} \times \text{fractional abundance})$

Additional info:

• Understanding atomic structure and ion formation is foundational for organic chemistry, as it underpins molecular bonding, reactivity, and the behavior of organic compounds.

• These concepts are directly relevant to Ch. 1 - Remembering General Chemistry, and are prerequisites for all subsequent organic chemistry topics.