Periodic Properties of the Elements (General Chemistry Study Notes)

Topic 2: Periodic Properties of the Elements (Ch. 8)

Overview

This topic covers the foundational concepts of the periodic table, electron configurations, and the periodic trends that govern the chemical and physical properties of elements. Understanding these principles is essential for predicting element behavior and reactivity in chemical contexts.

• The development of the periodic table

• Electron configurations, valence electrons, and the periodic table

• Periodic trends in atomic size and effective nuclear charge

• Ionic radii

• Ionization energy

• Electron affinities and metallic character

• Examples of periodic chemical behavior

The Development of the Periodic Table

Historical Background

The periodic table is a systematic arrangement of elements based on recurring chemical properties. Dmitri Mendeleev (1871) was instrumental in its development, observing that when elements were arranged by increasing atomic mass, those with similar properties appeared in vertical columns.

• Mendeleev's Table: Allowed prediction of undiscovered elements (e.g., gallium and germanium) and their properties.

• Modern Table: Arranged by atomic number (number of protons), not atomic mass.

• Quantum Theory: Explains why the periodic table works, as electron configurations determine chemical properties.

Element	Mendeleev's Predicted Properties	Actual Properties
Gallium (eka-aluminum)	Atomic mass ~68 u, Low melting point, Density 5.9 g/cm³	Atomic mass 69.72 u, Melting point 29.8°C, Density 5.9 g/cm³
Germanium (eka-silicon)	Atomic mass ~72 u, Density 5.5 g/cm³	Atomic mass 72.6 u, Density 5.35 g/cm³

Electron Configurations, Valence Electrons, and the Periodic Table

Electron Configurations

Electron configuration describes the arrangement of electrons in an atom's orbitals. The periodic table reflects these configurations, especially for the first 18 elements.

• Elements in the same column (group) have the same number of electrons in their outermost principal energy level (valence electrons).

• Electrons in lower energy levels are called core electrons.

Element	Electron Configuration	Valence Electrons	Core Electrons
Si	1s22s22p63s23p2	4	10
Ge	1s22s22p63s23p64s23d104p2	4	28

Example: Magnesium (Mg) has the electron configuration 1s22s22p63s2. It has 2 valence electrons in the 3s orbital.

Blocks of the Periodic Table

The periodic table is divided into four blocks based on the type of atomic orbital being filled:

• s-block: Groups 1-2

• p-block: Groups 13-18

• d-block: Transition metals (Groups 3-12)

• f-block: Lanthanides and actinides

This division helps in predicting electron configurations and chemical behavior.

Periodic Trends in Atomic Size and Effective Nuclear Charge

Effective Nuclear Charge ($Z_{eff}$)

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.

• Formula: $Z_{eff} = Z - S$

• $Z$ = atomic number (number of protons)

• $S$ = shielding constant (number of core electrons)

Slater's rules provide a more accurate method for calculating $Z_{eff}$ by considering electron penetration and shell structure.

Atomic Radius

The atomic radius is a measure of the size of an atom. The covalent radius is half the distance between nuclei of two identical atoms bonded together.

• Trend across a period: Atomic radius decreases from left to right due to increasing $Z_{eff}$.

• Trend down a group: Atomic radius increases due to addition of electron shells.

Example: Order the following by increasing atomic radius: Na < Mg < K < Rb.

Ionic Radii

Definition and Trends

Ions are atoms that have gained or lost electrons. Cations (positive) are smaller than their parent atoms due to reduced electron-electron repulsion and increased $Z_{eff}$. Anions (negative) are larger due to increased repulsion and decreased $Z_{eff}$.

• Isoelectronic series: Ions with the same number of electrons and electron configuration.

• Trend: For isoelectronic ions, the more positive the charge, the smaller the radius.

Ion	Relative Size
K+	Smaller than K
Cl-	Larger than Cl

Ionization Energy

Definition and Trends

Ionization energy (IE) is the minimum energy required to remove an electron from an atom or ion in the gas phase.

• First ionization energy ($IE_1$): $\text{Na}(g) \rightarrow \text{Na}^+(g) + e^-$

• Trend across a period: $IE_1$ increases from left to right due to increasing $Z_{eff}$.

• Trend down a group: $IE_1$ decreases due to increasing atomic size and electron distance from nucleus.

• Exceptions: Subshell stability (filled or half-filled subshells) can cause deviations from the trend.

Each successive ionization energy is larger than the previous, with a significant jump after all valence electrons are removed (due to stable noble gas configuration).

Electron Affinities and Metallic Character

Electron Affinity

Electron affinity (EA) is the energy change when an electron is added to an atom or ion in the gas phase.

• Trend across a period: EA becomes more negative (greater attraction) from left to right.

• Trend down a group: EA tends to remain similar or become less negative.

• Exceptions: Some groups (e.g., noble gases) have positive or near-zero EA due to stable configurations.

Example: $\text{Cl}(g) + e^- \rightarrow \text{Cl}^-(g)$, $EA = -349 \text{ kJ/mol}$

Metallic Character

Metallic character refers to the tendency of an element to lose electrons and exhibit properties of metals (conductivity, malleability).

• Trend across a period: Metallic character decreases from left to right.

• Trend down a group: Metallic character increases.

Examples of Periodic Chemical Behaviour

Applications and Review

• Group 15 elements have valence electron configuration $ns^2np^3$.

• The element in group 15 with the smallest $IE_1$ is the one lowest in the group (e.g., bismuth).

• Order of atomic radius (smallest to largest): Ar < Si < P < Mg < Na.

• Groups with negative electron affinities: Halogens (Group 17).

• Groups likely to form +1 ions: Alkali metals (Group 1).

Additional info: These notes expand on the provided slides and images, filling in academic context and definitions for clarity and completeness.