Periodicity and Ionic Bonding

Periodicity and Ionic Bonding Study Notes

Chapter Outline

Section 9.1 Valence Electrons
Section 9.2 Atomic and Ionic Sizes
Section 9.3 Ionization Energy and Electron Affinity
Section 9.4 Ionic Bonding
Section 9.5 Lattice Energy

Section 9.1: Valence Electrons

Electron Configuration Relation
- Relate each element’s electron configuration and position in the periodic table to its number of valence electrons.
Valence and Core Electrons
- The outermost, or valence, electrons display periodicity in electron configurations.
- Group Characteristics:
- Group 1 (1A) metals: 1 electron in outer s orbital.
- Group 2 (2A) metals: 2 electrons in outer s orbital.
- All electrons not classified as valence are known as core electrons.
Main Group Elements:
- Valence electrons exist in the highest occupied energy level (principal quantum number, n).
- Valence electrons are responsible for bonding, leading to similar chemical properties among elements with the same valence electron configurations.
- The number of valence electrons correlates with the “A” group number.
Example 9.1: Determining Electron Configurations
- For O, S, Se, and Te:
- O: $1s^2 2s^2 2p^4$
- S: $1s^2 2s^2 2p^6 3s^2 3p^4$
- Se: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^4$
- Te: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^4$
- These elements share the highest energy electron configuration of $ns^2 np^4$ leading to similar chemical behaviors.
Example 9.2: Determine the number of valence electrons in Na and Cl
- Na (Sodium): Group 1A → 1 valence electron.
- Cl (Chlorine): Group 7A → 7 valence electrons.
Section 9.1 Review
- Elements in a group have similar outer electron configurations.
- Main group elements' valence electrons are in the outermost shell (n), making their numbers predictable based on group positioning.

Section 9.2: Atomic and Ionic Sizes

Trends in Atomic and Ionic Radii
- Understand periodic trends related to atomic and ionic radii through effective nuclear charge.
Determining Sizes
- Atomic and ionic sizes depend on their electronic structures and the electrostatic interactions between the electrons and nucleus.
- Electrons bear a $-1$ charge, while protons bear a $+1$ charge.
Electrostatic Principles:
1. Oppositely charged particles attract each other.
2. Like-charged particles repel each other.
3. Force intensity increases with charge magnitude.
4. Attraction/repulsion increases as charged bodies come closer.
Effective Nuclear Charge, $Z_{eff}$:
- The perceived positive charge experienced by an electron in a multi-electron atom due to core electron shielding.
- Defined by the equation: $Z_{eff} = Z - S$ where:
- $Z$ = actual nuclear charge (number of protons)
- $S$ = shielding constant, accounting for electron repulsion.
Application of Shielding to Elements
- For Sodium (Na):
- Has 1 valence electron in the 3s orbital & 10 core electrons shielding (10 × 0.85 for core electron contributions).
- For Magnesium (Mg):
- Has 2 valence electrons (2 × 0.35) and 10 core electrons (10 × 0.85).
Example 9.4: Calculate Effective Nuclear Charge for K and Ca
- Using Slater's rules, calculate shielding and effective nuclear charge:
- K: $S = 0(0.35) + 18(0.85) = 15.3$ → $Z_{eff} = 19 - 15.3 = +3.7$
- Ca: $S = 1(0.35) + 18(0.85) = 15.65$ → $Z_{eff} = 20 - 15.65 = +4.35$
Atomic Radius Trends
- Generally decreases left to right across a period due to increasing nuclear positive charge affecting Zeff.
- Atomic size increases down a group due to increasing principal quantum number (larger orbitals).
Example 9.5: Compare atomic radii of O and S
- Both in Group 16 (6A); S (3rd period) has valence electrons in n = 3, whereas O (2nd period) has n = 2 → S is larger than O.
Ionic Radius
- Cations: Removing an electron decreases negative charge, increasing attraction between remaining electrons and the nucleus → cations are smaller than neutral atoms.
- Anions: Adding an electron increases negative charge, decreasing nucleus-electron attraction and increasing repulsion among electrons → anions are larger than neutral atoms.
Example 9.6: Size comparison between Cl and Cl−
- Cl has 17 protons and electrons; Cl− has 17 protons and 18 electrons → extra electron increases electron-electron repulsion, making Cl− larger.

Section 9.3: Ionization Energy and Electron Affinity

Ionization Energy (IE):
- The energy needed to remove an electron from a gaseous atom.
- Example equation:
- $ext{Na(g)} ightarrow ext{Na}^+(g) + e^-$
- Periodic trends relate inversely with atomic radius and directly with Zeff.
- As atoms become larger (decreasing Zeff), the first ionization energy typically decreases.
Trends in Ionization Energy
- IE generally decreases down a group; IE increases across a period from left to right.
Example 9.8: Highest first ionization energy
- Sets to evaluate include:
- a. K, Ga, Se → Se has the highest IE
- b. O, S, Se → O has the highest IE
- c. In, As, Cl → Cl has the highest IE
Exceptions to Ionization Energy Trends
- Subshell Occupation: When a subshell starts to fill, the first IE can be lower (easier to remove) than expected; e.g., IE1( ext{Be}) > IE1( ext{B}).
- Paired Electrons: Example IE1( ext{N}) > IE1( ext{O}): The paired electrons in O's configuration increase repulsion, requiring less energy to remove.
Electron Affinity (EA):
- The energy change when an electron is added to a gaseous atom.
- Defined by the equation:
- $ext{Cl(g)} + e^- ightarrow ext{Cl}^-(g)$ ; usually results in an exothermic process (negative EA values), while noble gases show positive EA.
Trends in Electron Affinity:
- General trend of more negative EA values left to right in the periodic table.

Section 9.4: Ionic Bonding

Forming Ionic Compounds:
- Metal atoms transfer electrons to nonmetal atoms, resulting in ionic bonds through electrostatic attraction.
- Example: $ext{Na} ightarrow ext{Na}^+ + e^−$ , $ext{Cl} + e^− ightarrow ext{Cl}^−$ .
Stable Noble Gas Configurations:
- Atoms achieve stability by having noble gas electron configurations through electron transfer.
Ionic Lattice Structure:
- Ionic compounds form large 3D lattices defined by the chemical formula termed as a formula unit.
- Lattice strength arises from the collective attractive forces between the myriad oppositely charged ions.
Example 9.11: Electron configurations of Ca and Br
- Ca: $ext{Ca} ightarrow ext{Ca}^{2+} + 2e^-$
- Br: $ext{Br} + e^- ightarrow ext{Br}^-$
- Formation of ionic compound identified as $ext{CaBr}_2$ : combining 1 Ca²⁺ and 2 Br⁻ results in neutrality.

Section 9.5: Lattice Energy

Lattice Energy Concept:
- The energy change when gaseous ions combine to form a solid ionic compound.
- Highly exothermic for ionic formation. E.g., $ext{K}^+(g) + ext{Cl}^-(g) ightarrow ext{KCl}(s)$ with $ ext{ΔH}_L = −717 ext{ kJ/mol}$.
Calculating Lattice Energy
- Utilizes Born-Haber cycle, where lattice energy cannot be directly measured but inferred from several indirect energetic changes.
Example 9.12: Lattice energy for $ext{CaCl}_2$
- Chemical equation showing conversion of gaseous ions to solid lattice:
- $ext{Ca}^{2+}(g) + 2 ext{Cl}^−(g) ightarrow ext{CaCl}_2(s)$ .
Periodic Trends in Lattice Energy:
- Ionic size increases group-wise leading to decreased lattice energy due to reduced attractive forces.
- Lattice energy inversely correlates with ionic radius owing to larger ionic sizes resulting in diminished polarizing effects, leading to weake electrostatic attractions:
 $E{el} = k rac{q1q_2}{d}$ where:
- $q1$ and $q2$ = charges of ions.
- $d$ = distance between ion centers.
Example 9.14: Arranging Lattice Energies
- Given compounds such as $ext{MgCl}2, ext{NaF}, ext{CsF}, ext{ScN}$ sorted by ionic charge and size leading to the order: ( ext{CsF} < ext{NaF} < ext{MgCl}2 < ext{ScN}).
9.5 Section Review: Lattice energy calculated through the Born-Haber cycle depicts its key role in defining stability in ionic compounds, emphasizing the correlation between ionic charge, size, and the released energy in ionic bond formation.