Chapter 4 – Atomic Theory, Periodic Table & Ion Formation

Introduction: Atoms, Elements & Matter

• Atoms compose all matter; the properties of atoms dictate the properties of the bulk substance.

• Atom = smallest identifiable unit of an element.

• Element = substance that cannot be chemically broken into simpler substances.

• Nature contains ≈ 91 distinct elements (≈ 91 kinds of naturally-occurring atoms).

• Scientists have synthesized ≈ 20 additional elements (man-made, not naturally found).

Early Atomic Ideas (460–370 BCE)

• Greek philosopher Democritus proposed matter is divisible only down to tiny, indestructible units called “atoms.”

• Atoms were thought to be indivisible, eternal, and the fundamental building blocks of everything.

• Though purely philosophical, the idea set the stage for later scientific inquiry.

Dalton’s Atomic Theory (1800s)

1. Each element is made of tiny, indestructible particles called atoms.

2. All atoms of one element have the same mass and properties that distinguish them from atoms of other elements.

3. Atoms combine in simple, whole-number ratios to form compounds (never half-atoms).

• Significance: Provided the first widely accepted, testable model linking the particulate nature of matter to measurable chemical behavior.

Discovery of the Electron — J. J. Thomson (1897)

• Used cathode-ray tube experiments.

• Findings:

  • Existence of electrons (negatively charged sub-atomic particles).

  • Electrons are much smaller & lighter than atoms.

  • Electrons are present in many substances, implying universality.

• Charge balance insight: Because atoms are overall neutral, there must be positive charge to counterbalance electron negativity.

Plum Pudding Model (Thomson’s Proposal)

• Atom pictured as a uniformly positive “pudding” with embedded negative electron “plums.”

• Predicted minimal deflection of charged particles passing through matter because charge is spread out.

Rutherford’s Gold Foil Experiment (1909–1911)

• Setup: Directed high-energy, positively charged α-particles at a thin gold foil.

• Expected (plum-pudding): Straight-through passage with tiny deviations.

• Observed:

  • Most α-particles passed straight through.

  • A few were deflected at large angles; some even bounced straight back.

• Interpretation → Nuclear Theory:

  1. Most of atom’s mass & all positive charge concentrated in a small, dense nucleus.

  2. Electrons occupy a large volume around the nucleus; most of the atom is empty space.

  3. Number of electrons (−) equals number of protons (+) → electrically neutral atom.

• Size analogy: If nucleus were a baseball, the atom would be the size of a football stadium.

Sub-Atomic Particles: Mass, Charge & Symbols

• Mass comparison: mp \approx 2000\,me.

• Neutrons add mass but no charge; crucial for isotope stability.

• Electron mass often considered negligible when computing total atomic mass.

Principles of Electrical Charge

• Fundamental property of protons (+) & electrons (−).

• Opposite charges attract; like charges repel.

• Charge neutrality occurs when total + equals total -: p^+ + e^- \rightarrow 0.

Atomic Number (Z) & Identity of Elements

• Atomic number Z = number of protons in nucleus.

• Element identity is fixed by Z; change Z ⇒ change element.

  • Hydrogen: Z=1 (always 1 proton)

  • Helium: Z=2

• In any neutral atom: \text{# electrons} = Z.

• Periodic Table arranges elements in ascending Z (today up to Z=118).

Reading the Periodic Table

• Groups / Families: vertical columns; elements share similar properties.

• Periods: horizontal rows (1–7).

• Lanthanides (Z 57–71) & Actinides (Z 89–103): often pulled out for compact display.

Element Symbols & Naming Conventions

• Symbols: 1–2 letters; first letter capital, second lower-case (e.g., Co = cobalt).

• Some derived from Latin:

  • Potassium: K (Kalium)

  • Sodium: Na (Natrium)

• Naming inspirations:

  • Properties: Argon (Greek argos = inactive)

  • Places: Polonium (Poland), Francium (France), Americium (USA)

• Correct capitalization distinguishes elements (Co) from compounds (CO = carbon monoxide).

Periodic Law & Mendeleev’s Arrangement (1869)

• Periodic Law: When elements are arranged by increasing mass (now atomic number), periodic recurrence of properties emerges.

• Mendeleev grouped similar properties in columns, successfully predicted missing elements & their properties.

• Underlying explanation later provided by quantum theory (Chapter 9).

Broad Classifications on the Periodic Table

• Metals (left & center; majority):

  • Good conductors (heat, electricity)

  • Malleable (flatten), ductile (draw into wire)

  • Lustrous (shiny)

  • Tend to lose electrons → form cations (positive ions)

• Non-metals (upper right + H):

  • Varied states (solids, liquids, gases)

  • Poor conductors

  • Tend to gain electrons → form anions (negative ions)

• Metalloids / Semimetals (along zig-zag line: B, Si, Ge, As, Sb, Te, etc.):

  • Mixed properties

  • Semiconductors; conductivity tunable (basis of modern electronics)

Named Groups of the Main Table

• Group 1: Alkali metals (very reactive, +1 ions)

• Group 2: Alkaline earth metals (+2 ions)

• Group 17: Halogens (−1 ions; highly reactive non-metals)

• Group 18: Noble gases (inert; filled valence shells)

• Transition metals (center block): properties less predictable from position.

Ion Formation: Gaining & Losing Electrons

• Ion = atom with net charge.

• Cation: positive, formed by electron loss (e.g., \text{Mg}^{2+}).

• Anion: negative, formed by electron gain (e.g., \text{O}^{2-}).

• Charge notation: magnitude first, sign second (e.g., 2+ not +2).

Logic & Examples

1. Oxygen ion \text{O}^{2-}

   • Z=8 → 8 protons.

   • 2- ⇒ gained 2 electrons: 8+2=10 electrons.

   • Charge check: +8 + (−10) = −2.

2. Chloride ion \text{Cl}^{-}

   • Z=17.

   • −1 ⇒ gained 1 e⁻ → 18 e⁻.

   • Net charge: +17 + (−18) = −1.

3. Magnesium ion \text{Mg}^{2+}

   • Z=12.

   • +2 ⇒ lost 2 electrons: 12−2 = 10 e⁻.

   • Net charge: +12 + (−10) = +2.

Practical / Real-World Connections

• Electronics: Semiconductor behavior of metalloids underpins modern computing & solar cells.

• Ion formation explains salt formation, battery chemistry, nerve impulses.

• Nuclear density insight influences understanding of nuclear energy & stellar processes.

Conceptual & Philosophical Notes

• Scientific laws (e.g., Periodic Law) summarize patterns; theories (e.g., Quantum theory) provide explanations.

• Historical evolution shows science as iterative: philosophical ideas → experimental evidence → refined models.

• Ethical implication: Synthesis of new elements expands knowledge but raises questions on radioactive waste & resource allocation.

Key Numbers & Equations (All in SI or Atomic Units)

• Proton mass: m_p = 1.67262\times10^{-27}\;\text{kg}.

• Neutron mass: m_n = 1.67493\times10^{-27}\;\text{kg}.

• Electron mass: m_e = 9.10938\times10^{-31}\;\text{kg}.

• Atomic mass unit definition: 1\,\text{u} = \tfrac{1}{12}\,m(^{12}\text{C}).

• Charge neutrality: Z = #e^-\;\text{(neutral atom)}.

• Ion electron count: e^- = Z − \text{charge} (charge positive for cations, negative for anions).

Summary Checklist

• [ ] Atoms: definition & ubiquity.

• [ ] Historical milestones: Democritus → Dalton → Thomson → Rutherford.

• [ ] Sub-atomic particles: masses, charges, roles.

• [ ] Nuclear model vs. Plum pudding.

• [ ] Atomic number & periodic arrangement.

• [ ] Metal, non-metal, metalloid properties.

• [ ] Group names & typical ion charges.

• [ ] Ion formation rules & examples.

Review this sheet before tackling exercises on atomic structure, periodic trends, and ion notation.