Understanding the Periodic Table

Introduction to the Periodic Table

• The periodic table is a tabular display of chemical elements
• Describes the "Map of the Building block of matter"
• Lists elements in order of increasing atomic number

Orbitals and Sublevels

• s Orbital:

  • Shape: Sphere
  • Name: Sharp
  • Orientation: One

• p Orbital:

  • Shape: Dumbbells
  • Name: Principal
  • Orientation: Three

• d Orbital:

  • Shape: Double dumbbells
  • Name: Diffuse
  • Orientation: Five

• f Orbital:

  • Shape: Complex dumbbells
  • Name: Fundamental
  • Orientation: Seven

Energies and Quantum Numbers

• Energies of subshells increase in the order: s < p < d < f
• Principle quantum number (n) determines the number of subshells
  • n = 1: 1s
  • n = 2: 2s, 2p
  • n = 3: 3s, 3p, 3d
  • n = 4: 4s, 4p, 4d, 4f

Orbitals and Electron Capacity

• s sublevel: 1 orbital
• p sublevel: 3 orbitals
• d sublevel: 5 orbitals
• f sublevel: 7 orbitals
• Each orbital can hold a maximum of 2 electrons

Blocks of the Periodic Table

• The s-block has 1 orbital
• The p-block has 3 orbitals
• The d-block has 5 orbitals
• The f-block has 7 orbitals

Principles Governing Electronic Configuration

• Pauli Exclusion Principle:

  • Each orbital can accommodate a maximum of two electrons with opposite spins
  • No more than two electrons may occupy the same orbital
  • No two electrons in one atom can have all four quantum numbers the same

• Aufbau Principle:

  • Electrons occupy the lowest possible energy levels
  • Orbitals fill so that those of the lowest energy are filled first

• Hund’s Rule:

  • For orbitals in the same sublevel, the number of electrons with the same spin is maximized
  • Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied

Electronic Configuration Notation

• 1s² 2s² 2p⁶ 3s² 3p³ is an example of notation
  • The large number (n) is the principal quantum number, which represents the principle energy level
  • The letter (l) is the angular quantum number, which represents the sublevel in the energy level
  • The superscript represents the number of electrons in the sublevel

Noble Gas Notation

• Chemical symbol of a Noble Gas in brackets becomes part of the electron configuration
• K₁₉ = [Ar]₁₈ 4s¹

Exceptional Cases in Electronic Configuration

• Chromium (Cr₂₄): Expected [Ar] 4s² 3d⁴, Actual [Ar] 4s¹ 3d⁵ (half-filled d-orbital)
• Copper (Cu₂₉): Expected [Ar] 4s² 3d⁹, Actual [Ar] 4s¹ 3d¹⁰ (full-filled d-orbital)
• Greater stability exists in half-filled (d⁵) or completely filled (d¹⁰) d subshells

Blocks on the Periodic Table and Energy Levels

• Main-group elements: s-block and p-block elements
• d-block elements:
  • Begin the transition metals
  • Begin with the 3rd energy level, but appear on the 4th row
  • d-block electrons are n - 1 from their location on the periodic table
• f-block elements:
  • Begin with the lanthanide series (#57)
  • This series is part of the sixth row, but the orbital energies put the electrons in the fourth energy level
  • The actinide series (#89) is the second row of the f-block elements
  • This series is part of the seventh row, but the orbital energies put the electrons in the fifth energy level
  • f-block electrons are n - 2 from their location on the periodic table

Periodic Trends

• 1. Size of atoms and ions
• 2. Ionization energies
• 3. Electron Affinity
• 4. Electronegativity
• 5. Polarizing power and polarizability (Fajan's Rule)
• 6. Metallic character
• 7. Diagonal Relationships in the Periodic Table

Size of Atoms and Ions

• Atomic Radius:
  • Largest atomic species are found in the SW corner of the periodic table since these atoms have the largest n.
    • The size of atoms decreases from left to right across period
    • On descending a group, the size of the atoms increases
    • Positive ions are smaller than the neutral atoms
    • Negative ions are larger than the neutral atoms

Ionization Energy

• The energy required to remove the most loosely bound electron from an isolated gaseous atom
  • There are: first, second, and third ionization energies. It depends on:
  • Size of atoms
  • The charge on the nucleus
  • The shielding effect of inner electron shells
  • The type of electrons involved
    • Ionization energy decreases on descending a group, and increases on crossing a period

Shielding Effect

• The effect where core electrons screen or shield valance electrons from the nuclear charge.
• The order of shielding: s>p>d>f
• The shielding effect is the reduction of attractive force between the nucleus (+) and its outer electrons (-) due to the blocking affect of the inner electrons
  • Across-stays the same: Electrons are added in the valence shell and the shielding electrons remain the same.
  • Down increases: Another layer of shielding is added between the valence shell and the nucleus

Electron Affinity

• The energy released when an extra electron is added to a neutral gaseous atom.
  • Depends on the size of atoms and the effective nuclear charge
  • As you go down the group, electron affinities become less

Electronegativity

• The tendency of an atom to attract electrons to itself when combined in a compound.
  • A tendency to increase from left to right across a period and decreases from up to down in a group
  • The Pauling scale is the most commonly used. Fluorine (the most electronegative element) has a value of 4.0, and values range down to cesium and francium which are the least electronegative at 0.7

Polarizing Power and Polarizability (Fajan's Rule)

• Polarizing power is the ability of the positive ion to polarize (electrically distort) by negative ions.
• Polarizability is the tendency of the negative ion to be distorted by the positive ion

Factors Affecting Polarization

• Large charge on the ions (cation and anion) Na ⁺ < Mg⁺⁺ < Al ⁺⁺⁺ and F⁻ < O ⁻² < N⁻³
• Small cation Be²⁺ > Mg²⁺ > Ca²⁺ > Sr⁺ > Ba ⁺² Their ionic radii are in the order
• Large anion: Larger anion = less strongly it can hold its outermost es

Metallic Character

• Metals are electropositive and have a tendency to lose electrons
  • The stronger this tendency is, the most electropositive and metallic an element Metallic character increases as we descend the groups and decreases from left to right.

Diagonal Relationships

• A relationship within the periodic table by which certain elements in the second period have a close chemical similarity to their diagonal neighbors in the next group of the third period.
• This similarity can be attributed to the atomic size, electronegativity, properties of compounds
• Crossing and descending the periodic table have opposing effects. Thus, on both descending a group and crossing by one element the changes cancel each other out, and elements with similar properties which have similar chemistry are often found.