Periodic Table: Structure, Families, Properties & Trends

Learning Objectives

• By the end of the lesson, you should be able to:

  • Know the evolution of the periodic table.

  • Explain periodic law.

  • Describe the organization of elements in the periodic table.

  • Predict elemental properties from position.

  • Identify metals, non-metals and metalloids by properties and location.

Historical Development of the Periodic Table

• Dmitri Mendeleev (1869)

  • Russian chemist/teacher.

  • Arranged elements according to increasing atomic mass & recurring similarities.

  • Left blank spaces and correctly predicted properties of then-unknown elements.

• Henry Moseley (1913)

  • Used X-ray spectroscopy.

  • Measured the nuclear charge → introduced the concept of atomic number $Z$.

  • Re-ordered table by increasing $Z$ (modern basis of periodic law).

  • Quote: “There is in the atom a fundamental quantity … the charge on the central positive nucleus.”

• Glenn Seaborg (1945)

  • Discovered & characterized transuranium elements Z>92.

  • Isolated the lanthanide & actinide series and placed them beneath the main body.

• Marie Curie (1867-1934)

  • Pioneered radioactivity research; discovered polonium & radium.

  • First woman to win a Nobel Prize; only person to win Nobel in two scientific fields (Physics & Chemistry).

  • Work led to cancer-treatment advances.

Periodic Law

• Statement: When elements are arranged in order of increasing atomic number, their physical & chemical properties repeat periodically. (also called Mendeleev’s Law in historical context).

Structural Vocabulary

Periods (Rows)

• Horizontal rows = periods/series.

• Properties change progressively across a period.

  • 1st element: very active solid.

  • Last element: inert gas.

• Row count: 7 periods in the modern table.

Groups (Columns)

• Vertical columns = groups/families.

• Elements within a group share similar—not identical—properties.

• Possess the same number of valence electrons.

• Modern table: 18 groups numbered 1–18 (or IA–VIIIA, IIIB–IIB traditional).

General Classification of Elements

Metals

• Good conductors of heat & electricity.

• Shiny (lustrous).

• Ductile: can be drawn into wire.

• Malleable: can be hammered into sheets.

• React with water → corrosion / basic oxides.

• Solid at room T except mercury (Hg).

• ≈91 of 118 known elements.

• Examples: Au, Ag, Fe, Cu.

Non-Metals

• Poor conductors (insulators).

• Non-ductile, non-malleable; solids are brittle & dull.

• Many exist as gases; Br is the only liquid; C, P, S, Se, I are solids.

• Total ≈17 recognized.

Metalloids (Semi-metals)

• Exhibit mixed properties.

• Usually brittle solids with metallic luster.

• Intermediate electrical & thermal conductivity → semiconductors.

• Form alloys like metals but often chemically behave like non-metals.

• Critical in electronics (computer chips, transistors, cell-phones, calculators).

Visual Layout

• Stair-step/zig-zag line separates metals (left) from non-metals (right); metalloids border the line.

• Two detached rows: top = lanthanides (57-71); bottom = actinides (89-103).

Seven Major Chemical Families

1. Alkali Metals (Group 1 / IA)

2. Alkaline Earth Metals (Group 2 / IIA)

3. Transition Metals (Groups 3–12 / IIIB–IIB)

4. Halogens (Group 17 / VIIA)

5. Noble Gases (Group 18 / VIIIA)

6. Lanthanides

7. Actinides

Detailed Family Profiles

Alkali Metals (Group 1)

• 1 valence electron; $\text{oxidation state} = +1$.

• Soft, shiny, clay-like; easily cut.

• Most reactive metals; react violently with water → $2\,\text{Na} + 2\,\text{H}<em>2\text{O} \to 2\,\text{NaOH} + \text{H}</em>2$ ↑ (exothermic).

• Never found free in nature; occur as salts (e.g., $\text{NaCl}$).

Alkaline Earth Metals (Group 2)

• 2 valence electrons; $\text{oxidation state} = +2$.

• Harder than alkalis; still reactive.

• Naturally found only in compounds (e.g., $\text{CaCO}_3$ chalk).

Transition Metals (Groups 3–12)

• Familiar structural metals: Fe, Cu, Ni, Zn, Ag, Au, etc.

• Good conductors; typically high melting/boiling points; all solids except Hg.

• Variable oxidation states (usually lose 1–2 outer electrons; sometimes involve $d$ electrons).

• Many colorful ionic compounds; used as pigments.

• Often form oxides (e.g., $\text{Fe}<em>2\text{O}</em>3$).

Halogens (Group 17)

• 7 valence electrons; $\text{oxidation state} = -1$ (gain 1 e⁻).

• Highly reactive non-metals.

• React vigorously with alkali metals to form salts (\"halogen\" = “salt-former”).

• React with H to give hydrohalic acids (HX) with bleaching properties (e.g., $\text{HCl}$).

• Exist in all three states at room T: $\text{F}<em>2, \text{Cl}</em>2$ gases; $\text{Br}<em>2$ liquid; $\text{I}</em>2$ solid.

Noble Gases (Group 18)

• Filled valence shell (8 e⁻) → extremely low reactivity (inert gases).

• Colorless, odorless.

• Uses: He in balloons (replaces flammable H₂); Ar in light-bulbs (non-reactive to filament).

Lanthanides (58–71)

• Follow lanthanum.

• Sometimes called rare-earth metals.

• 4f orbitals are internally shielded → chemistry differs from transition metals.

• Often form $+3$ ions; used in magnets, lasers, phosphors.

Actinides (89–103)

• All radioactive; many synthetic.

• 5f orbitals; wide range of oxidation states.

• U, Pu critical for nuclear energy & weapons.

Representative (Main-Group) Elements

• Groups 1A–7A (s & p blocks).

• Display a broad spectrum of properties.

• General valence relationship: Group number = number of valence electrons.

Valence / Oxidation details:

• Group 1A: 1 e⁻, $+1$.

• Group 2A: 2 e⁻, $+2$.

• Group 3A (Boron family): 3 e⁻, $+3$.

• Group 4A (Carbon family): 4 e⁻, $±4$ (commonly +4).

• Group 5A (Nitrogen family): 5 e⁻, $-3$ (can gain 3 e⁻).

• Group 6A (Oxygen family): 6 e⁻, $-2$.

• Group 7A (Halogens): 7 e⁻, $-1$.

Valence Electrons — Significance

• Outermost-shell electrons.

• Dictate electropositive/electronegative character, bond order, and chemical reactivity.

Periodic Trends

Atomic Radius

• Decreases left → right across a period (greater nuclear charge pulls electrons inward).

• Increases top → bottom within a group (additional energy levels).

Metallic Character

• Decreases left → right (elements become more non-metallic).

• Increases top → bottom inside a group.

Ionization Energy $IE$

• Energy required to remove an e⁻.

• $IE$ increases left → right (stronger nucleus–electron attraction).

• $IE$ decreases top → bottom (distance & shielding ease removal).

Electronegativity $\chi$

• Tendency to attract shared e⁻.

• $\chi$ increases left → right.

• $\chi$ decreases top → bottom.

• Range $0 \le \chi \le 4$ (Pauling scale).

• Predicts bond type (ionic, polar covalent, non-polar covalent) and molecular polarity.

Summary Diagram (verbal)

• Upper-right corner (F) → highest $IE$ & $\chi$, smallest radius.

• Lower-left corner (Fr/Cs) → largest radius, highest metallic character, lowest $IE$ & $\chi$.

Classroom Activity Prompts (Sample Questions)

• How many groups? 18.

• How many periods? 7.

• Position tasks (e.g., Group 10 / Period 5 = Pd; etc.).

• Most elements: metals.

• Non-metals location: right side.

• Alkali members (Li, Na, K, Rb, Cs, Fr) — properties: soft, reactive with H₂O, never free; naturally occur in compounds; react vigorously with water.

• Halogen members (F, Cl, Br, I, At, Ts) — name means “salt-forming”; form HX acids with H.

• Group 2 corresponds to Alkaline Earth Metals.

Real-World & Ethical Notes

• Radioactive actinides demand strict safety; enable nuclear energy with both societal benefits & risks.

• Rare-earth lanthanides underpin green tech (wind-turbines, EVs) yet raise concerns over mining impacts.

• Noble gas usage (He shortages) prompts conservation for critical medical imaging (MRI).

Key Numbers & Equations (LaTeX)

• General electron removal: $\text{M} + IE \to \text{M}^+ + e^-$.

• Halogen acid formation: $\text{X}<em>2 + \text{H}</em>2 \to 2\,\text{HX}$.

• Alkali metal–water reaction: $2\,\text{M} + 2\,\text{H}<em>2\text{O} \to 2\,\text{MOH} + \text{H}</em>2$ ↑.

Connections to Prior Knowledge

• Valence concept mirrors Lewis dot structures & octet rule.

• Periodic trends justify reactivity series (metals) & acid strength (non-metals).

• Transition metal color ties to crystal-field theory (d-orbital splitting).

Study Tips

• Memorize group valence counts; relate to oxidation states.

• When in doubt, locate element: row gives energy levels; column gives valence electrons & reactivity type.

• Visualize trends diagonally: lower-left = metallic & large; upper-right = non-metallic & small.