Periodic Trends: Ionization Energy, Electron Affinity, and Chemical Bonding

Periodic Trends and Chemical Bonding

Ionization Energy

Ionization energy is a fundamental periodic property that describes the energy required to remove an electron from a gaseous atom or ion. It provides insight into the reactivity and chemical behavior of elements.

• Definition: Ionization energy (IE) is the minimum energy needed to remove an electron from an atom in the gas phase.

• First Ionization Energy (IE1): The energy required to remove the first (most weakly held) electron.

• Successive Ionization Energies: IE2, IE3, etc., refer to the energies needed to remove additional electrons after the first. These values are always higher than the previous, as removing electrons from a positively charged ion is more difficult.

• Example: Sodium (Na) has the electron configuration $1s^2 2s^2 2p^6 3s^1$. The first electron removed is from the 3s orbital. The second ionization energy is much higher because it involves removing a core electron.

Element	IE1 (kJ/mol)	IE2 (kJ/mol)
Na	496	4560

• Key Point: Successive ionization energies are not additive; each subsequent electron is harder to remove due to increased positive charge.

• Periodic Trend:

  • Down a group: Ionization energy decreases (electrons are farther from the nucleus and less tightly held).

  • Across a period: Ionization energy increases (effective nuclear charge increases, electrons are held more tightly).

• Exceptions: There are exceptions to the trend due to electron configurations (e.g., between groups 2A and 3A, and 5A and 6A).

Electron Affinity

Electron affinity measures the energy change when an electron is added to a gaseous atom, reflecting the tendency of an atom to gain electrons.

• Definition: Electron affinity (EA) is the energy released when an atom in the gas phase gains an electron.

• Trend:

  • Across a period (left to right): Electron affinity becomes more negative (more favorable for gaining electrons).

  • Down a group: No clear trend, but generally becomes less negative.

• Conventions: A more negative value indicates a greater tendency to accept an electron.

• Exceptions: Some elements (e.g., Be, N, noble gases) have low or positive electron affinities due to stable electron configurations.

Metallic and Nonmetallic Character

The periodic table can be divided into metals and nonmetals, each with characteristic properties related to ionization energy and electron affinity.

• Metals:

  • Low ionization energies (easily lose electrons to form cations).

  • Low electron affinities (do not readily gain electrons).

• Nonmetals:

  • High ionization energies (difficult to lose electrons).

  • High (more negative) electron affinities (readily gain electrons to form anions).

• Trend: Metallic character decreases across a period and increases down a group.

Basic Concepts of Chemical Bonding

Lewis Dot Symbols

Lewis dot symbols are a simple way to represent the valence electrons of main group elements, which are crucial for understanding chemical bonding.

• Definition: The number of dots equals the number of valence electrons (group number for main group elements).

• Example: Sulfur (Group 6A) has 6 valence electrons, so its Lewis symbol has 6 dots.

• Noble Gases: Have 8 valence electrons (except He, which has 2), corresponding to a filled shell.

The Octet Rule

Atoms tend to gain, lose, or share electrons to achieve a noble gas electron configuration, typically 8 electrons in their valence shell (the octet rule).

• Application: Explains the formation of most chemical bonds.

• Exceptions: Hydrogen (duet rule), Be, B, and elements in period 3 and beyond (expanded octets).

Ionic Bonding

Ionic bonds form through the transfer of electrons from metals to nonmetals, resulting in the formation of oppositely charged ions that attract each other.

• Formation: Metal atoms lose electrons to become cations; nonmetals gain electrons to become anions.

• Example: NaCl forms from Na+ and Cl- ions.

• Lattice Energy: The energy released when ions form a crystalline lattice. It is governed by Coulomb's law: $E = k \frac{Q_1 Q_2}{r}$ where $Q_1$ and $Q_2$ are the charges, $r$ is the distance between ions, and $k$ is a proportionality constant.

• Properties: Ionic compounds have high melting points due to strong electrostatic forces.

Compound	Lattice Energy (kJ/mol)
LiF	1036
LiCl	853
LiBr	807
LiI	757

Covalent Bonding and Lewis Structures

Covalent bonds involve the sharing of electron pairs between nonmetal atoms, resulting in molecules with specific shapes and properties.

• Single, Double, and Triple Bonds: Atoms can share one, two, or three pairs of electrons, respectively. Bond strength and bond length follow:

  • Bond strength: triple > double > single

  • Bond length: single > double > triple

• Octet Rule: Most atoms (except H, Be, B, and expanded octet elements) follow the octet rule in covalent compounds.

• Lewis Structures: Visual representations showing how valence electrons are arranged among atoms in a molecule.

• Steps for Drawing Lewis Structures:

  1. Sum valence electrons for all atoms (add for anions, subtract for cations).

  2. Determine the skeletal structure (least electronegative atom is central, H and F are never central).

  3. Connect atoms with single bonds.

  4. Distribute remaining electrons to satisfy octets (duet for H).

  5. Use multiple bonds if necessary to complete octets.

Resonance and Formal Charge

Some molecules or ions can be represented by more than one valid Lewis structure, called resonance structures. Formal charge helps determine the most stable structure.

• Resonance: Occurs when two or more Lewis structures can be drawn for a molecule, differing only in the placement of electrons.

• Example: Ozone (O3) and nitrite ion (NO2-) have resonance structures.

• Bond Order: Resonance leads to bond lengths that are intermediate between single and double bonds.

• Formal Charge: Calculated as: Formal charge = (valence electrons) – (lone pair electrons) – ½(bonding electrons) Structures with minimized formal charges and negative charges on the most electronegative atoms are favored.

Exceptions to the Octet Rule

Some molecules do not follow the octet rule, especially those with central atoms from period 2 (incomplete octet) or period 3 and beyond (expanded octet).

• Incomplete Octet: Be, B, and Al can have fewer than 8 electrons.

• Expanded Octet: Elements in period 3 or higher can have more than 8 electrons due to available d orbitals (e.g., PCl5, ICl4-).

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a chemical bond. Differences in electronegativity determine bond polarity.

• Trend: Increases across a period (left to right), decreases down a group. Fluorine is the most electronegative element.

• Bond Types by Electronegativity Difference (ΔEN):

  • ΔEN < 0.5: Nonpolar covalent

  • 0.5 ≤ ΔEN < 2.0: Polar covalent

  • ΔEN ≥ 2.0: Ionic

• Dipole Moment: A measure of bond polarity, with both magnitude and direction.

• Example: HCl is polar covalent; NaCl is ionic.

Summary Table: Bond Type Classification by Electronegativity Difference

ΔEN	Bond Type
< 0.5	Nonpolar covalent
0.5 – 1.9	Polar covalent
≥ 2.0	Ionic

Additional info: These notes synthesize and expand upon the provided lecture content, including definitions, trends, and examples for clarity and completeness.