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Periodic Table and Periodicity

Subject: Chemistry
Topic: 2
Cambridge Code: 0620 / 0971 / 5070

Organization of Periodic Table​

Periodic table - Elements arranged by atomic number and properties

Periods (Rows)​

  • Number of periods: 7
  • Period number: Highest shell occupied
  • Example: Period 1 has 1 shell, Period 3 has 3 shells

Groups (Columns)​

  • Number of groups: 18
  • Group number: Valence electrons
  • Example: Group 1 has 1 valence electron

Blocks​

s-block: Groups 1-2 (alkali, alkaline earth) p-block: Groups 13-18 (nonmetals, halogens, noble gases) d-block: Groups 3-12 (transition metals) f-block: Lanthanides and actinides

Atomic Radius​

Trend across period: Decreases (left to right)

  • Protons increase
  • Electrons in same shell
  • Increased nuclear charge pulls electrons closer

Trend down group: Increases (top to bottom)

  • New electron shells added
  • Distance from nucleus increases
  • Despite higher nuclear charge

Ionization Energy​

Ionization energy - Energy to remove electron

Trend across period: Increases (left to right)

  • Stronger nuclear attraction
  • More energy needed

Trend down group: Decreases (top to bottom)

  • Electrons farther from nucleus
  • Easier to remove
  • Inner shells shield

Exception: Noble gases have high ionization energies (full shells)

Electronegativity​

Electronegativity - Attraction for electrons in bond

Trend across period: Increases (left to right)

  • Stronger nuclear charge

Trend down group: Decreases (top to bottom)

  • Electrons farther away

Most electronegative: Fluorine (group 17, period 2)

Density​

Trend across period: Generally increases

  • More protons/neutrons
  • Metal density higher than nonmetals

Trend down group: Varies

  • Alkali metals decrease
  • Transition metals increase

Elements by Group​

Group 1: Alkali Metals​

  • Soft, silvery metals
  • Reactivity: Increases down group
  • 1 valence electron
  • Form +1 ions
  • React vigorously with water/oxygen
  • Examples: Lithium, sodium, potassium

Reactivity order: Li < Na < K < Rb

Group 2: Alkaline Earth Metals​

  • Harder than Group 1
  • Reactivity: Increases down group
  • 2 valence electrons
  • Form +2 ions
  • Less reactive than Group 1
  • Examples: Magnesium, calcium

Group 13-17: Main Group Nonmetals​

Group 13: Boron, aluminum

  • 3 valence electrons
  • Some metallic character

Group 14: Carbon, silicon

  • 4 valence electrons
  • Carbon: Life basis
  • Silicon: Semiconductors

Group 15: Nitrogen, phosphorus

  • 5 valence electrons
  • N₂ very inert (triple bond)

Group 16: Oxygen, sulfur

  • 6 valence electrons
  • O₂ essential for combustion
  • Form -2 ions (typically)

Group 17: Halogens

  • Highly reactive nonmetals
  • 7 valence electrons
  • Form -1 ions
  • Reactivity decreases down group
  • F₂ > Cl₂ > Br₂ > I₂
  • Examples: Fluorine, chlorine, bromine

Group 18: Noble Gases​

  • Inert (unreactive)
  • 8 valence electrons (full shell)
  • Complete electron configuration
  • Examples: Helium, neon, argon

Transition Metals​

Transition metals - d-block elements (Groups 3-12)

Characteristics​

  • Variable oxidation states
  • Form colored compounds
  • Often catalytic
  • Paramagnetic (attracted to magnetic field)
  • High melting/boiling points
  • Hard, dense

Examples​

Iron (Fe):

  • Forms Fe²⁺ and Fe³⁺
  • Catalyst in Haber process
  • Rusts (oxidation)

Copper (Cu):

  • Forms Cu⁺ and Cu²⁺
  • Blue solutions (Cu²⁺)
  • Excellent conductor

Manganese (Mn):

  • Variable oxidation states (+2 to +7)
  • Oxidizing agent

Metal vs Nonmetal Properties​

Metals​

  • Shiny
  • Conduct electricity
  • Conduct heat
  • Malleable (can be hammered)
  • Ductile (can be drawn into wire)
  • Sonorous (ringing sound)
  • Solid at room temp (except mercury)
  • Lose electrons (form cations)

Nonmetals​

  • Dull
  • Poor conductors (except graphite)
  • Poor heat conductors
  • Brittle
  • Not ductile
  • Gain/share electrons
  • Various states at room temp
  • Form anions or covalent compounds

Metalloids (Semimetals)​

  • Properties between metals/nonmetals
  • Examples: Silicon, arsenic, boron
  • Semiconductors

Metallic character (left to right, top to bottom):

  • Decreases across period
  • Increases down group
  • Metals left, nonmetals right

Reactivity:

  • Group 1: Increases down (more reactive)
  • Group 17: Decreases down (less reactive)
  • Transition metals: Moderate, variable

Key Points​

  1. Periodicity relates to electron configuration
  2. Trends: Atomic radius, ionization energy, electronegativity
  3. Groups have similar properties
  4. Periods show gradual property change
  5. Transition metals have variable properties
  6. Metallic character increases down/left

Practice Questions​

  1. Predict atomic radius comparison
  2. Explain ionization energy trend
  3. Compare alkali metal reactivity
  4. Describe halogen properties
  5. Explain noble gas inertness
  6. Predict element properties from position

Revision Tips​

  • Learn periodic trends clearly
  • Know group properties
  • Understand electron configuration link
  • Practice trend predictions
  • Know exceptions
  • Compare metallic/nonmetallic properties