Chapter 2 - Elements and the Periodic Table

Chapter 2: Elements and the Periodic Table

Key Knowledge

• Definitions of Elements, Isotopes, Ions

  • Elements: Substances that cannot be broken down chemically.

  • Isotopes: Atoms of the same element with different numbers of neutrons; same atomic number but different mass number.

  • Ions: Atoms that have gained or lost electrons, resulting in a charge.

    • Notation includes atomic number (Z), mass number (A), protons, neutrons, and electrons.

• Periodic Table as an Organisational Tool

  • Identifies patterns and trends among elements.

  • Displays structures and properties of elements, including:

    • Shell and subshell electronic configurations.

    • Atomic radii.

    • Electronegativity.

    • First ionisation energy.

    • Metallic and non-metallic character.

    • Reactivity.

• Critical Elements and Recycling

  • Examples: Helium, phosphorus, rare-earth elements, post-transition metals, and metalloids.

  • Importance of recycling processes for recovery of these elements.

The Atomic World

Structure of Atoms

• Nucleus Composition

  • Contains protons (positively charged) and neutrons (neutral).

Electrons

• Electron Configuration

  • Surround nucleus in negative cloud; negatively charged.

  • Electrostatic attraction between electrons and protons binds electrons to nucleus.

Discovering the Nucleus

• Ernest Rutherford's conclusions:

  • Most of an atom's mass is in the nucleus.

  • Atoms are mostly empty space.

  • Electrons orbit nucleus in circular paths.

Example: Carbon Atom

• In a neutral carbon atom, the number of protons equals the number of electrons.

Summary of Atomic Structure

• Protons: + charge in nucleus.

• Electrons: - charge around nucleus.

• Neutrons: Neutral in nucleus.

• Electrostatic forces bind electrons to protons in nucleus.

Atomic Symbol

• Each element has a unique chemical symbol (X).

• Atom type determined by number of protons (Z).

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

Analysing the Atomic Symbol

• Example: Iron atom is represented as

  • 56 Fe

  • Mass number = 56

  • Atomic number = 26

  • Protons = 26

  • Neutrons = 56 - 26 = 30

  • Electrons = 26

Isotopes

• Atoms with the same number of protons/electrons but different neutrons.

• Same atomic number (Z) and different mass number (A).

• Isotopes have similar chemical properties but different physical properties (mass, density).

• Some isotopes are radioactive.

Emission Spectra and the Bohr Model

The Bohr Model

• Electrons move in circular orbits (shells) around nucleus, corresponding to energy levels.

• Shells filled in ascending energy level order; larger radii = higher energy.

Electron Shells

• Heating may excite electrons, causing them to jump to a higher energy state.

Ground vs Excited State

• Ground State: Atom with electrons in lowest energy levels.

• Excited State: Electrons occupy higher energy levels.

Electron Configuration

• Electrons fill shells based on energy levels:

  • Rule 1: Max electrons per shell (e.g., 2 in first, 8 in second).

  • Rule 2: Lower energy shells fill first.

Valence Electrons

• Outermost shell (valence shell) contains valence electrons.

• Removal energy = ionisation energy.

• Valence electrons determine chemical properties.

The Schrodinger Model of the Atom

Revisiting the Bohr Model

• Electrons orbit rapidly in defined energy levels/shells represented by numbers (1-6) or letters (K, L, M, N, O).

Electron Configuration Examples

• Aluminium (Al): Total of 13 electrons across shells.

• Calcium (Ca): Total of 20 electrons, fills 4th shell before completely filling 3rd.

Subshell Occupancy Order

• S-block, p-block, d-block, f-block filled in order:

  1. s: max 2 electrons

  2. p: max 6 electrons

  3. d: max 10 electrons

  4. f: max 14 electrons

Exploring Valence Electrons with Examples

• Calcium: 2 valence electrons; loses to form Ca2+

• Magnesium: 2 valence electrons; becomes Mg2+

• Chlorine: 5 valence electrons; gains 1 electron to become Cl-

Ground State vs Excited State

• Filled subshells vs incorrect filling based on energy levels.

Ion vs Excited Ions

• Ions: More/less electrons than atomic number.

• Excited State: Incorrect filling order of subshells.

The Periodic Table

Features of the Periodic Table

• Arranged by increasing atomic number and periodic trends.

  • Periods: Horizontal rows (1-7).

  • Groups: Vertical columns (1-18).

• Provides information on electron configuration and valence electrons.

Groups of Elements In Periodic Table

• Valence Electrons: Involved in chemical reactions, defining chemical properties.

• Examples of Groups:

  • 1: Alkali metals.

  • 2: Alkaline earth metals.

  • 17: Halogens.

  • 18: Noble gases.

Trends in the Periodic Table

Core Charge

• Measure of attraction felt by valence electrons.

• Formula: Core Charge = Protons - Inner Shell Electrons.

• Increases across a period; constant down a group.

Atomic Radius/Size

• Decreases across a period (higher nuclear attraction).

• Increases down a group (more occupied shells).

Electronegativity

• Atom's ability to attract electrons.

• Increases across a period; decreases down a group.

Ionization Energy

• Energy needed to remove an electron.

• Increases across a period; decreases down a group.

Reactivity Trends

• Metals: Reactivity decreases across a period, increases down a group.

• Non-Metals: Reactivity increases across a period, decreases down a group.