CXC Chemistry: Atomic Structure and the Periodic Table

Introduction to Atomic Structure

The atom is the basic unit of matter that makes up all chemical elements. Understanding atomic structure is fundamental to chemistry as it explains how elements form compounds and undergo chemical reactions.

Historical Development of Atomic Theory

Dalton's Atomic Theory (1808): Proposed that all matter consists of tiny, indivisible particles called atoms.
Thomson's Plum Pudding Model (1897): Discovered electrons and proposed a model with electrons embedded in a positive sphere.
Rutherford's Nuclear Model (1911): Discovered the nucleus through the gold foil experiment, showing most of an atom's mass is concentrated in a small, dense, positively charged nucleus.
Bohr's Model (1913): Proposed that electrons orbit the nucleus in specific energy levels or shells.
Modern Quantum Mechanical Model: Describes electrons as occupying orbitals (regions of high probability) rather than definite paths.

Figure 1: Comparison of the Rutherford-Bohr Model and the Quantum Mechanical Model of the atom

Subatomic Particles

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons.

Particle	Symbol	Relative Mass	Relative Charge	Location
Proton	p or p⁺	1	+1	Nucleus
Neutron	n or n⁰	1	0	Nucleus
Electron	e or e^-	1/1836 (≈ 0)	-1	Electron cloud (orbitals)

Key Atomic Numbers

Atomic Number (Z): The number of protons in an atom's nucleus; determines the element's identity.
Mass Number (A): The total number of protons and neutrons in an atom.
Number of Neutrons: Mass number - Atomic number (A - Z)

Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. They have the same atomic number but different mass numbers.

Example: Hydrogen has three isotopes: protium (¹H), deuterium (²H), and tritium (³H).
Isotopes can be represented as: ^AX where A is the mass number and X is the element symbol.

¹H) p n Deuterium (²H) p n n Tritium (³H)

Figure 2: The three isotopes of hydrogen

Electronic Configuration

The arrangement of electrons in an atom's orbitals is called its electronic configuration.

Energy Levels and Sublevels

Electrons occupy discrete energy levels (shells) around the nucleus, labeled as n = 1, 2, 3, etc.
Each energy level contains sublevel orbitals (s, p, d, f) that can hold specific numbers of electrons:
- s orbital: maximum of 2 electrons
- p orbital: maximum of 6 electrons
- d orbital: maximum of 10 electrons
- f orbital: maximum of 14 electrons
The arrangement follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.

Writing Electronic Configurations

Electronic configurations can be written in two ways:

Full electronic configuration: Shows all orbitals (e.g., 1s² 2s² 2p⁶)
Condensed or noble gas notation: Uses the symbol of the previous noble gas in square brackets followed by the remaining configuration (e.g., [Ne] 3s¹)

Examples of Electronic Configurations

Hydrogen (H, Z=1): 1s¹
Carbon (C, Z=6): 1s² 2s² 2p² or [He] 2s² 2p²
Sodium (Na, Z=11): 1s² 2s² 2p⁶ 3s¹ or [Ne] 3s¹
Calcium (Ca, Z=20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² or [Ar] 4s²

Figure 3: Energy levels and sublevels in an atom

The Periodic Table

The periodic table organizes all known elements based on their atomic number and chemical properties.

Development of the Periodic Table

Döbereiner's Triads (1829): Grouped elements in threes with similar properties.
Newlands' Law of Octaves (1865): Noted that every eighth element had similar properties.
Mendeleev's Periodic Table (1869): Arranged elements by atomic mass and left gaps for undiscovered elements.
Modern Periodic Table: Arranged by increasing atomic number and electron configuration.

Structure of the Periodic Table

Periods: Horizontal rows (numbered 1-7), represent the principal energy level (n) of the outermost electrons.
Groups: Vertical columns (numbered 1-18), elements in the same group have similar chemical properties due to the same number of valence electrons.
Blocks: Based on the type of orbital being filled:
- s-block: Groups 1-2
- p-block: Groups 13-18
- d-block: Groups 3-12 (transition elements)
- f-block: Lanthanides and Actinides (separate rows at bottom)

Figure 4: Simplified structure of the Periodic Table showing the s, p, d, and f blocks

Classification of Elements

Main Groups and Their Properties

Group 1 (Alkali Metals): Soft, highly reactive metals with one valence electron (ns¹).
Group 2 (Alkaline Earth Metals): Less reactive than Group 1, with two valence electrons (ns²).
Group 17 (Halogens): Highly reactive non-metals with seven valence electrons (ns²np⁵).
Group 18 (Noble Gases): Unreactive gases with full outer shells (ns²np⁶, except helium: 1s²).
Transition Elements (d-block): Metals with partially filled d-orbitals, often forming colored compounds and showing variable oxidation states.

Metals, Non-metals, and Metalloids

Metals: Left and center of the periodic table. Generally have:
- Good conductors of heat and electricity
- Malleable and ductile
- Lustrous appearance
- Tend to lose electrons (form positive ions)
Non-metals: Right side of the periodic table. Generally have:
- Poor conductors of heat and electricity
- Brittle in solid state
- Dull appearance
- Tend to gain electrons (form negative ions)
Metalloids: Properties between metals and non-metals, found in a diagonal line separating metals and non-metals (B, Si, Ge, As, Sb, Te, Po).

Periodic Trends

Atomic Radius

Across a period: Decreases from left to right due to increasing nuclear charge pulling electrons closer to the nucleus.
Down a group: Increases due to additional electron shells.

Ionization Energy

The energy required to remove an electron from a gaseous atom.

Across a period: Increases from left to right due to increasing nuclear charge and decreasing atomic radius.
Down a group: Decreases due to increasing atomic radius and shielding effect.
Successive ionization energies (removing additional electrons) always increase.

Electron Affinity

The energy change when a gaseous atom accepts an electron.

Across a period: Generally becomes more exothermic (more negative) from left to right.
Down a group: Generally becomes less exothermic (less negative).
Halogens have the most negative electron affinities due to being one electron away from a noble gas configuration.

Electronegativity

The ability of an atom to attract the shared pair of electrons in a covalent bond.

Across a period: Increases from left to right.
Down a group: Decreases.
Fluorine is the most electronegative element.