Periodicity and Ionic Bonding: Chapter 9 Study Notes

Periodicity and Ionic Bonding

Chapter Overview

• 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

Valence Electrons

Valence and Core Electrons

Valence electrons are the outermost electrons in an atom and are responsible for chemical bonding and periodic trends. Core electrons are all other electrons not involved in bonding.

• Group 1 (1A) metals: One electron in their outer s orbital.

• Group 2 (2A) metals: Two electrons in their outer s orbital.

• Core electrons: All electrons not in the outermost shell.

Valence Electrons in Main Group Elements

• Located in the highest occupied energy level (principal quantum number, n).

• Elements with the same valence electron configuration have similar chemical properties.

• The number of valence electrons equals the "A" group number.

Periodicity of Valence Electrons

The periodic table shows a repeating pattern in the number of valence electrons across periods and groups.

Group	Number of Valence Electrons
1A	1
2A	2
3A	3
4A	4
5A	5
6A	6
7A	7
8A	8

Example: Electron Configurations and Chemical Properties

• 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: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^4$

All have the same highest energy configuration: $ns^2 np^4$.

Example: Valence Electrons in Na and Cl

• Na: Group 1A, 1 valence electron.

• Cl: Group 7A, 7 valence electrons.

Example: Identifying Elements by Valence Electrons

• Row 3, 2 valence electrons: Mg

• Row 5, 6 valence electrons: Te

Atomic and Ionic Sizes

Sizes of Atoms and Ions

Atomic and ionic sizes are determined by electronic structure and electrostatic interactions between the nucleus and electrons.

• Electrons: -1 charge, outside nucleus.

• Protons: +1 charge, inside nucleus.

• Electrostatic principles: Opposite charges attract, like charges repel, and force increases with charge and proximity.

Effective Nuclear Charge ($Z_{eff}$)

The net positive charge experienced by an electron in a multi-electron atom.

• Formula: $Z_{eff} = Z - S$

• $Z$ = number of protons, $S$ = shielding constant

• Valence electrons experience lower $Z_{eff}$ due to shielding by core electrons.

Slater's Rules for Shielding

• Valence electrons (same n): 0.35 each

• Core electrons (lower n): 0.85 each

• Formula: $S = ext{valence electrons} imes 0.35 + ext{core electrons} imes 0.85$

Example: $Z_{eff}$ for Sodium and Magnesium

• Na: $Z = 11$, $S = 8.5$, $Z_{eff} = 2.5$

• Mg: $Z = 12$, $S = 8.85$, $Z_{eff} = 3.15$

Trends in $Z_{eff}$

• Increases across a period (left to right).

• Decreases slightly down a group.

Atomic Radius

• Decreases across a period due to increasing $Z_{eff}$.

• Increases down a group due to higher principal quantum number ($n$).

Example: Oxygen vs. Sulfur

• Sulfur (n=3) has a larger atomic radius than oxygen (n=2).

Ionic Radius

Cations

• Smaller than neutral atoms due to increased attraction to nucleus.

Anions

• Larger than neutral atoms due to increased electron-electron repulsion.

Example: Cl vs. Cl-

• Cl- is larger than Cl due to extra electron and increased repulsion.

Isoelectronic Ions

• Compare nuclear charge to determine size: more protons = smaller radius.

• Example: Cl- (17p), K+ (19p), Ca2+ (20p); Ca2+ is smallest, Cl- is largest.

Ionization Energy and Electron Affinity

Ionization Energy (IE)

Energy required to remove an electron from a gaseous atom.

• Equation: $Na(g) ightarrow Na^+(g) + e^-$

• IE decreases as atomic radius increases (down a group).

• IE increases across a period (left to right).

• Second ionization energy ($IE_2$) is always greater than first ($IE_1$).

Exceptions to IE Trends

• New subshells: Lower IE when a new subshell begins.

• Electron pairing: Paired electrons in the same orbital repel, lowering IE.

Electron Affinity (EA)

Energy change when an electron is added to a gaseous atom to form an anion.

• Negative EA: Energy released (exothermic).

• Positive EA: Energy required (endothermic).

• EA values become more negative across a period (until noble gases).

• Halogens have the most negative EA values.

Ionic Bonding

Formation of Ionic Compounds

• Metals lose electrons to form cations; nonmetals gain electrons to form anions.

• Electrostatic attraction between oppositely charged ions forms ionic bonds.

• Atoms tend to achieve noble gas electron configurations when forming ions.

Ionic Lattice

• Ionic compounds form large 3D lattices, not discrete pairs.

• The chemical formula represents the simplest ratio of ions (formula unit).

Example: CaBr2

• Ca loses 2 electrons to form Ca2+ (noble gas configuration).

• Br gains 1 electron to form Br- (noble gas configuration).

• Formula unit: CaBr2

Lattice Energy

Definition and Calculation

Lattice energy is the energy released when gaseous ions form a solid ionic compound.

• Equation: $K^+(g) + Cl^-(g) ightarrow KCl(s)$

• Calculated using the Born-Haber cycle and Hess's law.

Born-Haber Cycle Steps

1. Formation of gaseous atoms from elements.

2. Ionization of metal atom.

3. Electron affinity of nonmetal atom.

4. Formation of ionic solid from gaseous ions (lattice energy).

General Formula for Lattice Energy

• $ext{Lattice Energy} = -( ext{Enthalpy of formation} + ext{Bond energy} + ext{IE} + ext{EA})$

Trends in Lattice Energy

• Decreases as ionic radius increases (down a group).

• Increases with higher ionic charges.

• Shorter bond lengths and higher charges yield more negative (stronger) lattice energies.

Compound	Bond Length (pm)	Lattice Energy (kJ/mol)
LiF	106	1036
MgO	85	3850
NaF	136	923
CsF	238	740

Order of Lattice Energies (Least to Most Negative)

• CsF < NaF < MgCl2 < ScN

Summary of Key Trends

• Valence electrons determine chemical properties and periodicity.

• Atomic radius decreases across a period, increases down a group.

• Cations are smaller, anions are larger than their neutral atoms.

• Ionization energy increases across a period, decreases down a group.

• Electron affinity is most negative for halogens.

• Ionic compounds form 3D lattices; lattice energy depends on ionic size and charge.