Periodic Properties of the Elements and Electron Configurations

Chapter 3: Periodic Properties of the Elements

Overview

This chapter covers the periodic trends and properties of elements, focusing on electron configurations, effective nuclear charge, atomic radius, ionization energy, and the formation of ions. Understanding these concepts is essential for predicting chemical behavior and bonding patterns.

Electron Configurations

Electron configuration describes the arrangement of electrons in an atom's orbitals. It is fundamental to understanding chemical properties and periodic trends.

• Condensed Electron Configuration: Uses the previous noble gas as a starting point, followed by the remaining electron configuration. For example, copper (Cu, Z = 29): [Ar]4s13d10.

• Valence Electrons: The electrons in the outermost (highest energy) principal energy level (n). For main-group elements, only count electrons in the highest n value. For transition elements, count electrons in the highest n value plus those in unfilled inner d sublevels.

• Example: Arsenic (As, Z = 33): [Ar]4s23d104p3. Number of valence electrons = 5 (2 from 4s and 3 from 4p).

• Transition Elements: For elements like zirconium (Zr, Z = 40): [Kr]5s24d2. Valence electrons = 2 (from 5s) + 2 (from 4d) = 4.

Summary Table: Counting Valence Electrons

Element Type	How to Count Valence Electrons
Main-group	Count electrons in highest n
Transition	Count (highest n) + (unfilled inner d)

Irregular (Anomalous) Filling Patterns

Some transition metals do not follow the expected order of orbital filling due to extra stability associated with half-filled or fully filled sublevels (Hund's Rule). For example, chromium (Cr, Z = 24) has the configuration [Ar]4s13d5 instead of [Ar]4s23d4.

• Key Point: Half-filled and fully filled d sublevels confer extra stability.

Metals and Nonmetals

• Metals: Located on the lower left and middle of the periodic table. Good conductors, malleable, ductile, shiny. Tend to lose electrons to form cations.

• Nonmetals: Located on the upper right. Poor conductors, varied properties. Tend to gain electrons to form anions.

Noble Gas Connection

Elements tend to react to achieve a noble gas electron configuration (full valence shell, usually ns2np6). This explains why metals lose electrons and nonmetals gain electrons. Noble gases are especially stable and nonreactive due to their filled valence shells.

The Formation of Ions

• Predicting Ionic Charge:

  • Nonmetals: Form negative ions (anions) with charge = Group # - 8.

  • Metals: Form positive ions (cations) with charge = Group #.

Periodic Trends

Effective Nuclear Charge ($Z_{\text{eff}}$)

The effective nuclear charge is the net positive charge experienced by valence electrons. It is less than the actual nuclear charge due to shielding by core electrons.

• Trend: $Z_{\text{eff}}$ increases across a period (left to right) and slightly increases down a group.

Atomic Radius

• Definition: The atomic radius is half the distance between the nuclei of two identical atoms bonded together.

• Trends:

  • Decreases across a period (left to right) due to increasing $Z_{\text{eff}}$.

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

Ionization Energy ($IE$)

• Definition: The energy required to remove one mole of electrons from one mole of gaseous atoms or ions.

• First Ionization Energy ($IE_1$): $\text{Atom (g)} \rightarrow \text{Ion}^+ (g) + e^-$

• Trends:

  • Increases across a period (left to right).

  • Decreases down a group.

  • Irregularities: Group 5 to 6 shows a drop due to extra stability of half-filled sublevels.

• Successive Ionization Energies: Each successive electron requires more energy to remove, especially after all valence electrons are gone.

Electron Affinity (EA)

• Definition: The energy change when one mole of electrons is added to one mole of gaseous atoms.

• Trend: Generally becomes more negative (more exothermic) across a period.

Electron Configurations of Ions

Main-Group Elements

• Atoms lose or gain electrons to achieve a filled outer level (ns2np6), becoming isoelectronic with the nearest noble gas.

• Anions (Nonmetals): Gain electrons in p orbitals until noble gas configuration is reached. Example: $\text{O} ([He]2s^22p^4) + 2e^- \rightarrow \text{O}^{2-} ([He]2s^22p^6)$

• Cations (s-block metals): Lose all electrons with the highest n value. Example: $\text{Mg} ([Ne]3s^2) \rightarrow \text{Mg}^{2+} ([Ne]) + 2e^-$

• Cations (p-block metals): Lose np electrons before ns electrons; heavier elements may form two cations. Example: $\text{Sn} ([Kr]5s^24d^{10}5p^2) \rightarrow \text{Sn}^{2+} ([Kr]5s^24d^{10}) + 2e^-$

Transition Metal Elements

• Transition metal ions rarely achieve noble gas configurations.

• They lose ns electrons before (n-1)d electrons. Example: $\text{Fe} ([Ar]4s^23d^6) \rightarrow \text{Fe}^{2+} ([Ar]3d^6) + 2e^-$

Chapter 4: Molecules and Compounds

Types of Chemical Bonds

• Ionic Bonds: Electrostatic attraction between cations and anions.

• Covalent Bonds: Sharing of electrons between two atoms.

• Metallic Bonds: Metal atoms bond to several other atoms (not covered in detail in this course).

Types of Chemical Formulas

• Empirical Formula: Shows the lowest whole-number ratio of atoms in a compound. Example: Hydrogen peroxide: HO

• Molecular Formula: Shows the actual number of atoms of each element in a molecule. Example: Hydrogen peroxide: H2O2

Molecular Models

• Structural Formulas: Show the order in which atoms are attached but not the 3D shape.

• Ball-and-Stick Models: Show the 3D shape of molecules.

The Lewis Model

• Lewis Symbols: Use dots to represent valence electrons around the element symbol.

• Steps for Drawing Lewis Symbols:

  1. Note the element's A-group number (number of valence electrons).

  2. Place one dot at a time on each of the four sides of the symbol.

  3. Pair dots until all valence electrons are used up.

• Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons.

• Metals: The number of dots equals the number of electrons lost to form a cation. Example: Mg: will lose 2 electrons to form Mg2+.

• Nonmetals: The number of unpaired dots equals the number of electrons gained to form an anion or shared in covalent bonds. Example: S: will gain 2 electrons to form S2- or share 2 electrons in covalent bonds.