Electronic Configuration and Chemical Properties

Electronic Configuration and Chemical Properties of Elements

Understanding the electronic configuration of an atom helps predict its chemical behavior and essential properties, especially when utilizing the periodic classification of elements.

2.1 Electronic Configuration of an Atom (Schrödinger Model)

Definition: The electronic configuration of an atom describes the distribution of all electrons around the nucleus in various energy levels.

How to Determine Electronic Configuration?

This involves applying specific rules or principles.

First Principle: Energy Minimization (Aufbau Principle)

• Electrons are distributed starting from the lowest energy level (i.e., according to increasing n).

Second Principle: Pauli Exclusion Principle

“Two electrons cannot have the same values for all four quantum numbers.”

• In an atomic orbital (defined by n, l, m), the only differentiating quantum number is spin ($s = +1/2$ or $s = -1/2$).
• This means a maximum of two electrons with opposite spins can occupy the same orbital.
• Example for B (Z=5): After filling 1s², 3 electrons remain. The next energy level is $n=2$.
  • For $n=2$, $l=0$ defines the 2s orbital, which can hold 2 electrons.
  • For $n=2$, $l=1$ defines three 2p orbitals (m = -1, 0, +1), which can hold 6 electrons in total.

Third Principle: Hund's Rule (Rule of Maximum Multiplicity)

“When there are several equivalent and isoenergetic atomic orbitals (i.e., identical values for n and l), the maximum number of orbitals must be occupied with parallel spins first.”

Example for C (Z=6):

• Electronic configuration: $1s^2 2s^2 2p^2$.
• The two 2p electrons will occupy separate 2p orbitals with parallel spins, rather than pairing up in one orbital.

Fourth Principle: Klechkowski's Rule (n+l Rule)

• This rule refines the Aufbau principle when energy levels converge.
• The filling of energy levels occurs according to increasing values of (n+l).
• If two orbitals have the same (n+l) value, the one with the lower n value is filled first.
• Energy Order: E_{1s} < E_{2s} < E_{2p} < E_{3s} < E_{3p} < E_{4s} < E_{3d} < E_{4p} < E_{5s} < E_{4d} \dots
• Important note: E_{4s} < E_{3d} (energy order), but r_{4s} > r_{3d} (spatial extent). As n increases, the atomic radius r generally increases.

Electrons and Valence Configuration

• Valence electrons: Those in the outermost shell ($n_{max}$) or sometimes also from the ($n_{max-1}$) non-saturated shell (e.g., transition metals).
• These electrons are primarily responsible for an atom's chemical properties.
• The presence of lone pairs, unpaired electrons, or empty orbitals in the valence shell significantly influences chemical behavior.

Examples:

• B (Z=5): $1s^2 2s^2 2p^1$. Valence electrons: 3 ($2s^2 2p^1$).
• C (Z=6): $1s^2 2s^2 2p^2$. Valence electrons: 4 ($2s^2 2p^2$).
• Cl (Z=17): $1s^2 2s^2 2p^6 3s^2 3p^5$. $n_{max}=3$. Valence electrons: 7 ($3s^2 3p^5$), distributed as 3 lone pairs and 1 unpaired electron.
• Fe (Z=26): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6$. Valence configuration: $3d^6 4s^2$. Valence electrons: 8 ($2 (4s) + 6 (3d)$).

Lewis Representation

The Lewis structure focuses on the peripheral (valence) shell electrons responsible for chemical bonding.

• H: H· ($1s^1$)
• Cl: :Cl⋅ ($3s^2 3p^5$)

Periodic Classification of Elements (Periodic Table)

• A universal reference table where elements are classified by increasing atomic number (Z), reflecting the periodicity of their chemical and physical properties.
• Organized into 18 columns (groups/families) and 7 rows (periods).
• Each column groups elements with the same number of valence electrons, leading to similar chemical properties.
• Each row (period) corresponds to the principal quantum number (n) of the valence shell.

Key Elements in Biology:

• 96% of human body mass: C, H, O, N.
• Essential 4%: Na, K, Mg, Ca, P, S, Cl.
• Trace elements: Fe, I, F, Mn, Zn, Mo, Cu, Co, Cr, Se, Ni, B.

Classification by Blocks:

• s-block: Groups 1 and 2 (e.g., Li, Na, K, Mg, Ca). Alkali metals ($ns^1$) and Alkaline earth metals ($ns^2$). Highly electropositive, form monovalent cations easily.
• p-block: Groups 13 to 18 (e.g., B, C, N, O, F, Cl, Ne, Ar). Contains non-metals, metalloids, and some metals.
• d-block: Transition metals (Groups 3-12) (e.g., Sc, Ti, Fe, Ni, Cu). Tend to form multiple cations with varying valencies.
• f-block: Lanthanides and Actinides (inner transition metals).

Periodic Properties

Atomic Radius ($r_{at}$)

• Down a group: $r_{at} \uparrow$ (increases) due to increasing n (more electron shells).
• Across a period (left to right): $r_{at} \downarrow$ (decreases) due to increasing effective nuclear charge ($Z_{eff}$) pulling electrons closer.

Ionic Radii

• Cations: r_{ionique} (cation) < r_{atomique} (atom). Losing electrons means a smaller electron cloud and increased effective nuclear charge on remaining electrons. (e.g., $Na^+$ is smaller than Na).
• Anions: r_{ionique} (anion) > r_{atomique} (atom). Gaining electrons increases electron-electron repulsion and expands the electron cloud. (e.g., $Cl^-$ is larger than Cl).

Electronegativity (EN or $\chi$)

• Definition: A physical quantity characterizing an atom's ability to attract electrons in a chemical bond.
• Up a group: EN $\uparrow$ (increases).
• Across a period (left to right): EN $\uparrow$ (increases).
• The most electronegative element is Fluorine (F) at the top right of the periodic table.
• Elements with high EN are non-metals and act as oxidizing agents (electron acceptors).
• Differences in electronegativity between bonded atoms are crucial for understanding chemical properties.

Electropositivity and Ionization Energy (Ei)

• Definition: An atom displaying electropositive character tends to lose electrons to become a cation ($X \xrightarrow{Ei > 0} X^{n+} + ne^-$).
• Ionization Energy (Ei): The energy required to remove an electron from an atom.
• Ei is directly proportional to $Z_{eff}$ and inversely proportional to the electron-nucleus distance (r).
• More electropositive elements require less Ei to lose an electron.
• These are typically large atoms located at the bottom and left of the periodic table.
• Electropositive elements are metals, good conductors, and powerful reducing agents (electron donors).

Octet Rule

“An atom tends to achieve the electron configuration of the nearest noble gas.”

• This often drives chemical behavior (e.g., Na loses 1 electron to become like Ne, Cl gains 1 electron to become like Ar).

Chemical Behavior by Family

• Alkali Metals ($ns^1$) & Alkaline Earth Metals ($ns^2$): Highly electropositive, readily lose valence electrons to form stable cations.
• Transition Metals (d-block): Tend to form multiple cations with varying valencies due to the involvement of both outer s and inner d electrons.
• Non-metals (e.g., Halogens $ns^2 np^5$): Highly electronegative, often gain electrons to form anions. Are strong oxidants (e.g., Cl, F).
• Carbon Family ($ns^2 np^2$): Central position, often prefer covalent bonding rather than forming ions.
• Noble Gases ($1s^2$ or $ns^2 np^6$): Stable, mostly unreactive due to full valence shells.