Periodic Properties of Elements

Periodic Properties of Elements

• Atomic Radii: The distance between the nucleus and the outermost shell of electrons in an atom.
  • Covalent Radii: Half the distance between the nuclei of two similar atoms joined by a covalent bond.
  • Metallic Radii: Half the distance between the nuclei of two adjacent metal atoms in a metallic crystal.
  • Van Der Waals Radii: The distance between the nuclei of two non-bonded, but adjacent atoms in a molecule or crystal.
• Ionic Radii: The distance between the nucleus of an ion and its outermost electron shell. Ions form by adding or removing electrons from neutral atoms.
  • Cations: Smaller than their parent atoms; removing electrons increases nuclear attraction.
  • Anions: Bigger than their parent atoms; adding electrons increases electron-electron repulsions.
• Ionization Energy: The minimum energy required to remove one electron from a gaseous atom in its ground state, forming a cation.
  • First Ionization Energy: Energy to remove the first electron.
  • Second Ionization Energy: Energy to remove the second electron.
  • Third Ionization Energy: Energy to remove the third electron.
  • Ionization energy generally increases across a period due to increased nuclear charge, and decreases down a group due to increased shielding and larger atomic size.
• Factors Affecting Ionization Energy:
  • Atomic Radii: Smaller radii correlate with higher ionization energies.
  • Nuclear Charge: Higher nuclear charge increases ionization energy.
  • Shielding Effect/ Screening Effect: Increased shielding reduces attraction between the nucleus and valence electrons, decreasing ionization energy.
  • Penetration Effect: The ability of an electron to penetrate inner electron shells affects shielding and effective nuclear charge. s electrons penetrate more than p, d, and f electrons.
  • Half Filled and Fully Filled Orbitals: Atoms with half-filled or fully filled orbitals have higher ionization energies due to the stability of these configurations.
• Successive Ionization Energies: Energy required to remove successive electrons, generally increases due to increasing positive charge on the ion, and significant jumps occur when removing core electrons.
• Photoionization: The process where an electron is removed from an atom or molecule by a photon. Energy from a photon ejects an electron.

Ionization Energy

• Ionization energy is the minimum energy to remove the most loosely bound electron from a gaseous atom in its ground state.
• Higher ionization energy means it's harder to remove an electron.
• Atoms with full outer shells (s² or p⁶) have high ionization energies.
• Atoms with half-filled outer shells (p³) have high ionization energies.
• Be has higher ionization energy than B because Be's complete 2s orbital makes it more stable, thus less likely to lose an electron.
• N has higher ionization energy than O because N's half-filled 2p orbital is more stable.
• Ionization energy generally increases across a period due to increasing nuclear charge, and decreases down a group due to increasing atomic size.

Electron Affinity

• Electron affinity is the change in energy when an electron is added to a gaseous atom in its ground state.
• Negative electron affinity indicates energy release; positive indicates energy absorption.
• First electron affinity is usually negative (energy released); second is usually positive (energy absorbed due to repulsions).
• Electron affinity generally increases across a period due to increased nuclear charge, and decreases down a group due to increasing atomic size.
• Cl has higher electron affinity than F due to smaller interelectronic repulsions in the larger 3p orbital of Cl.

Factors Affecting Electron Affinity

• Atomic size: Smaller atoms have higher electron affinities because of stronger attraction.
• Effective nuclear charge: Higher effective nuclear charge leads to greater attraction for electrons, thus increasing electron affinity.
• Screening effect: Larger screening effect weakens attraction for added electrons, reducing electron affinity.
• Orbital type: Electrons are more readily added to s orbitals due to greater attraction to the nucleus.

Predicting Electron Affinity Trends

• Electron affinity generally increases across a period due to increased effective nuclear charge and decreased atomic size.
• Electron affinity generally decreases down a group due to increased atomic size and weaker attraction.

Electronegativity

• Electronegativity is a measure of an atom's ability to attract electrons within a chemical bond.
• Highly electronegative atoms attract electrons and create partial negative charges.
• Electronegativity generally increases across a period due to increased nuclear charge and decreases down a group due to increasing atomic size.
• Higher electronegativity leads to stronger attraction of electrons in a covalent bond.

Factors Affecting Electronegativity

• Atomic size: Smaller atoms have higher electronegativity due to stronger nuclear pull.
• Type of ion: Positively charged ions are more electronegative than neutral atoms; negatively charged ions are less electronegative.
• Nuclear charge: Higher nuclear charge increases electronegativity.
• Hybridization: Higher s character in hybrid orbitals (sp > sp² > sp³) leads to higher electronegativity.

Predicting Electronegativity Trends

• Electronegativity increases across a period due to increasing nuclear charge and decreasing atomic size.
• Electronegativity decreases down a group due to increasing atomic size and weaker attraction.

Methods for Determining Electronegativity

• Pauling scale: Common and widely used electronegativity scale.
• Mulliken scale: Based on ionization energy and electron affinity.
• Allred-Rochow scale: Considers nuclear charge and atomic size.

Effective Nuclear Charge

• Effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in an atom.
• Z_eff is calculated by subtracting the number of core electrons from the atomic number.
• Z_eff increases across a period due to increased nuclear charge and a roughly constant number of core electrons.
• Z_eff decreases down a group due to increases in inner shells, shielding valence electrons.

Screening Effect/ Shielding Effect

• Screening effect: Reduction in the attraction between the nucleus and outer electrons due to inner electrons. Inner electrons repel outer electrons.
• Atomic number represents the total positive charge, attracting electrons.
• Screening effect reduces the effective nuclear charge experienced by outer electrons, causing weaker attraction.

Effective Nuclear Charge

• Effective nuclear charge (Z_eff) is the actual attraction felt by an outer electron from the nucleus, after considering the screening effect.
• The formula for effective nuclear charge calculation is: Z_eff = Z - S (Z = atomic number, S = screening constant)

Screening Constant (S)

• Screening constant (S) indicates the degree of shielding from inner electrons.
• S = 0.35: For the last shell electrons.
• S = 0.85: For the shell just before the outermost shell.
• S = 1: For all other inner shells.

Calculating Effective Nuclear Charge

• Step 1: Determine the screening constant (S) for the specific electron in question.
• Step 2: Apply the formula: Z_eff = Z - S
• Step 3: The result is the effective nuclear charge (Z_eff).

Example: Calculating Effective Nuclear Charge for Helium (He)

• Helium has an atomic number (Z) of 2; 2 electrons in the 1s shell.
• Screening constant (S) for the last shell electron is 0.3 (approximation)
• Z_eff = 2 - 0.3 = 1.7

Example: Calculating Effective Nuclear Charge for Carbon (C)

• Carbon has an atomic number (Z) of 6. Electronic configuration: 1s²2s²2p². Calculate Z_eff for a 2p electron.
• Screening constant (S): - Screening from 2s electrons: 2 * 0.35 = 0.70 - Screening from 1s electrons: 2 * 0.85 = 1.70
• S = 0.70 + 1.70 = 2.4
• Z_eff = 6 - 2.4 = 3.6

This quiz covers essential periodic properties of elements, focusing on atomic, covalent, metallic, van der Waals, and ionic radii. Understand the differences and implications of these measurements in chemical bonding and structure. Test your knowledge of how these radii influence the behavior of elements in the periodic table.