Theory Exercises

The Periodic Table

The periodic table is one of the most important tools in chemistry. It organizes all known elements in a way that reveals patterns and allows us to predict properties and behaviors of elements.

History and Development

Early Attempts at Organization

Before the modern periodic table, scientists tried various ways to organize elements:

  • Lavoisier (1789): First list of elements
  • Döbereiner (1829): Triads - groups of three similar elements
  • Newlands (1864): Law of Octaves - every 8th element was similar
  • Meyer (1870): Arranged by atomic volume

Mendeleev's Breakthrough (1869)

Dmitri Mendeleev created the first successful periodic table:

  • Arranged by atomic mass: Listed elements in order of increasing atomic mass
  • Periodic law: Properties repeat periodically
  • Predicted elements: Left gaps for undiscovered elements
  • Predicted properties: Accurately described missing elements
Mendeleev's amazing predictions
Mendeleev's bold move: When arranging elements by atomic mass, Mendeleev noticed some didn't fit the pattern. Instead of forcing them in, he:
  • Left blank spaces for "missing" elements
  • Predicted their properties based on neighbors
  • Sometimes switched order when properties didn't match
His predictions for "eka-silicon" (later discovered as Germanium):
  • Predicted: Atomic mass ~72, gray metal, density 5.5 g/cm³
  • Actual germanium: Atomic mass 72.6, gray metal, density 5.3 g/cm³
  • Amazing accuracy! This proved his periodic law was correct
Other successful predictions:
  • Eka-aluminum: Became gallium (1875)
  • Eka-boron: Became scandium (1879)
  • These discoveries validated Mendeleev's approach

Modern Periodic Table

Henry Moseley (1913) discovered the correct organizing principle:

  • Atomic number: Number of protons, not atomic mass
  • Fixed problems: Explained Mendeleev's "exceptions"
  • Current form: Based on electron configuration

Organization of the Periodic Table

Basic Structure

Periods (Horizontal Rows)
  • Definition: Elements with same number of electron shells
  • Number: 7 periods currently
  • Period number = number of electron shells
  • Pattern: Properties change systematically across periods
Groups/Families (Vertical Columns)
  • Definition: Elements with same number of valence electrons
  • Number: 18 groups
  • Similar properties: Elements in same group behave similarly
  • Numbering: 1-18 (modern) or 1A-8A, 1B-8B (traditional)

Element Information

Each element box typically contains:

  • Atomic number: Number of protons (top)
  • Element symbol: 1-2 letter abbreviation
  • Element name: Full name
  • Atomic mass: Average mass of atoms (bottom)

Major Groups and Families

Group 1: Alkali Metals

  • Elements: Li, Na, K, Rb, Cs, Fr
  • Valence electrons: 1
  • Properties: Soft metals, highly reactive, low density
  • Reactions: Explosive reaction with water
  • Examples: Sodium in salt, potassium in bananas

Group 2: Alkaline Earth Metals

  • Elements: Be, Mg, Ca, Sr, Ba, Ra
  • Valence electrons: 2
  • Properties: Harder than alkali metals, reactive
  • Examples: Calcium in bones, magnesium in chlorophyll

Groups 3-12: Transition Metals

  • Elements: Sc through Zn, Y through Cd, etc.
  • Properties: Hard, high melting points, good conductors
  • Special features: Variable oxidation states, colored compounds
  • Examples: Iron, copper, gold, silver

Groups 13-16: Main Group Elements

  • Group 13: Boron family (3 valence electrons)
  • Group 14: Carbon family (4 valence electrons)
  • Group 15: Nitrogen family (5 valence electrons)
  • Group 16: Oxygen family (6 valence electrons)

Group 17: Halogens

  • Elements: F, Cl, Br, I, At
  • Valence electrons: 7
  • Properties: Very reactive non-metals
  • Forms: Diatomic molecules (F₂, Cl₂, etc.)
  • Examples: Chlorine in bleach, fluorine in toothpaste

Group 18: Noble Gases

  • Elements: He, Ne, Ar, Kr, Xe, Rn
  • Valence electrons: 8 (except He with 2)
  • Properties: Unreactive, stable, colorless gases
  • Examples: Helium in balloons, neon in signs
Why noble gases are so stable
Complete outer shells:
  • Noble gases have complete outer electron shells
  • Helium: 2 electrons (full first shell)
  • Others: 8 electrons (full outermost shell)
Octet rule:
  • Atoms are most stable with 8 outer electrons
  • Noble gases naturally have this configuration
  • Other elements try to achieve this by bonding
Why they don't react:
  • No tendency to gain or lose electrons
  • Already in lowest energy state
  • Very difficult to force them to react
Recent discoveries:
  • Some noble gas compounds have been made
  • XeF₄, KrF₂ under extreme conditions
  • Still considered "unreactive" under normal conditions

Periodic Trends

The periodic table reveals predictable patterns in element properties:

1. Atomic Radius

Across a Period (Left to Right)
  • Trend: Atomic radius decreases
  • Reason: More protons pull electrons closer
  • Example: Na > Mg > Al > Si > P > S > Cl
Down a Group
  • Trend: Atomic radius increases
  • Reason: More electron shells added
  • Example: Li < Na < K < Rb < Cs

2. Ionization Energy

Energy needed to remove an electron

Across a Period
  • Trend: Ionization energy increases
  • Reason: Electrons held more tightly
  • Highest: Noble gases (complete shells)
Down a Group
  • Trend: Ionization energy decreases
  • Reason: Outer electrons farther from nucleus
  • Lowest: Alkali metals (easy to lose 1 electron)

3. Electronegativity

Ability to attract electrons in bonds

Patterns
  • Highest: Fluorine (top right, excluding noble gases)
  • Lowest: Francium (bottom left)
  • Increases: Left to right across periods
  • Decreases: Top to bottom in groups

4. Metallic Character

Distribution
  • Metals: Left side and center of table
  • Non-metals: Upper right
  • Metalloids: Diagonal line between metals and non-metals
Trends
  • Decreases: Left to right across periods
  • Increases: Top to bottom in groups
  • Most metallic: Bottom left (Fr, Cs)
  • Most non-metallic: Top right (F, O, N)
Understanding periodic trends with examples
Atomic radius trend example - Period 3:
  • Na (11 protons): Large atom, 1 valence electron
  • Mg (12 protons): Smaller, more nuclear charge
  • Al (13 protons): Even smaller
  • Cl (17 protons): Much smaller, strong nuclear pull
Ionization energy example - Group 1:
  • Li: 520 kJ/mol (small atom, electron close to nucleus)
  • Na: 496 kJ/mol (larger atom, electron farther out)
  • K: 419 kJ/mol (even larger, easier to remove electron)
  • Cs: 376 kJ/mol (largest, easiest to ionize)
Reactivity patterns:
  • Alkali metals: More reactive down the group (easier to lose electron)
  • Halogens: More reactive up the group (better at gaining electrons)
  • Most reactive combination: Cs + F → CsF

Electron Configuration and the Periodic Table

Relationship to Periodic Structure

The periodic table's structure reflects how electrons fill atomic orbitals:

Blocks of the Periodic Table
  • s-block: Groups 1-2 (s orbitals filling)
  • p-block: Groups 13-18 (p orbitals filling)
  • d-block: Groups 3-12 (d orbitals filling)
  • f-block: Lanthanides and actinides (f orbitals filling)

Valence Electrons

The number of valence electrons determines chemical behavior:

Main Group Elements
  • Group number = valence electrons (for groups 1-2, 13-18)
  • Group 1: 1 valence electron
  • Group 14: 4 valence electrons
  • Group 17: 7 valence electrons
  • Group 18: 8 valence electrons (except He with 2)

Predicting Properties

Using Position to Predict Behavior

Chemical Reactivity
  • Group 1: Lose 1 electron easily, form +1 ions
  • Group 2: Lose 2 electrons, form +2 ions
  • Group 17: Gain 1 electron, form -1 ions
  • Group 18: Don't react under normal conditions
Bonding Patterns
  • Metals + Non-metals: Ionic bonds
  • Non-metals + Non-metals: Covalent bonds
  • Metals + Metals: Metallic bonds

Physical Properties

State at Room Temperature
  • Metals: Mostly solids (except Hg)
  • Non-metals: Gases or brittle solids
  • Noble gases: All gases
  • Halogens: F₂, Cl₂ (gases), Br₂ (liquid), I₂ (solid)

Applications of Periodic Trends

Material Science

  • Semiconductor design: Silicon and germanium (Group 14)
  • Superconductors: Transition metal compounds
  • Catalysts: Transition metals with variable oxidation states

Medicine and Biology

  • Essential elements: Ca for bones, Fe for blood
  • Toxic elements: Heavy metals like Pb, Hg
  • Radioisotopes: Medical imaging and treatment

Environmental Chemistry

  • Pollution monitoring: Detecting toxic elements
  • Water treatment: Using chemical properties for purification
  • Atmospheric chemistry: Ozone depletion, greenhouse gases

Special Features

Lanthanides and Actinides

The f-block elements are usually shown separately:

Lanthanides (Elements 57-71)
  • Also called: Rare earth elements
  • Properties: Similar chemical properties
  • Uses: Electronics, magnets, catalysts
  • Examples: Neodymium (magnets), europium (TV phosphors)
Actinides (Elements 89-103)
  • Properties: All radioactive
  • Natural: Only Th, Pa, U occur naturally
  • Synthetic: Others made in laboratories
  • Uses: Nuclear power, weapons, research