Elements and Periodicity
Demarcation of the Periodic Table into s, p, d, and f-Blocks
The periodic table is systematically divided into four blocks based on the type of valence orbital where the outermost electrons are located. This classification helps chemists predict an element’s chemical behavior, bonding patterns, and physical properties.
1. s-Block Elements
The s-block consists of Groups IA (alkali metals) and IIA (alkaline earth metals), along with hydrogen. These elements have their outermost electrons in the s-orbital, following the general configuration ns¹–ns².
Important Characteristics:
- Highly reactive metals (except hydrogen, which is a nonmetal)
- Low ionization energies, making them strong reducing agents
- Form +1 (Group IA) or +2 (Group IIA) ions in compounds
- Examples:
- Sodium (Na) – [Ne] 3s¹ (Group IA, Period 3)
- Calcium (Ca) – [Ar] 4s² (Group IIA, Period 4)
These metals are soft, shiny, and excellent conductors of electricity. They react vigorously with water, with reactivity increasing down the group (e.g., potassium reacts more violently than sodium).
2. p-Block Elements
The p-block spans Groups IIIA to VIIIA, including metals, metalloids, and nonmetals. Their valence electrons occupy the p-orbital, following the configuration ns² np¹–np⁶.
Important Characteristics:
- Diverse chemical behavior:
- Metals (e.g., aluminum)
- Metalloids (e.g., silicon, used in semiconductors)
- Nonmetals (e.g., oxygen, chlorine)
- Form covalent bonds (e.g., carbon in organic chemistry)
- Halogens (Group VIIA) are highly electronegative and reactive
- Noble gases (Group VIIIA) are inert due to full valence shells
Example: Chlorine (Cl) – [Ne] 3s² 3p⁵ (Group VIIA, Period 3).
3. d-Block Elements (Transition Metals)
The d-block includes Groups IB to VIIIB (Groups 3–12), where electrons fill the d-orbitals. Their general configuration is (n-1)d¹–¹⁰ ns⁰–².
Key Characteristics:
- Variable oxidation states (e.g., Fe²⁺ and Fe³⁺)
- Form colored compounds (due to d-d electron transitions)
- High melting points and good conductivity
- Catalytic properties (e.g., platinum in car exhaust systems)
Example: Iron (Fe) – [Ar] 3d⁶ 4s² (Group VIIIB, Period 4).
4. f-Block Elements (Inner Transition Metals)
The f-block consists of lanthanides (4f) and actinides (5f), located separately at the bottom of the periodic table. Their configuration is (n-2)f¹–¹⁴ (n-1)d⁰–¹ ns².
Important Characteristics:
- Lanthanides (e.g., neodymium) are used in strong magnets
- Actinides (e.g., uranium) are radioactive and used in nuclear energy
- High density and melting points
Example: Uranium (U) – [Rn] 5f³ 6d¹ 7s² (Actinide series).
Determining Group, Period, and Block Using Electronic Configuration
Step-by-Step Guide
- Identify the Period:
- The period number corresponds to the highest principal quantum number (n) in the configuration
- Example: Magnesium (Mg) – [Ne] 3s² → Period 3
- Determine the Group:
- s-Block: Group = Number of valence electrons (e.g., Li [2s¹] → Group IA)
- p-Block: Group = 10 + (s + p electrons) (e.g., O [2s² 2p⁴] → Group VIA)
- d-Block: Group = Sum of (n-1)d + ns electrons (e.g., Cr [3d⁵ 4s¹] → Group VIB)
- Assign the Block:
- Based on the last subshell filled (s, p, d, or f)
| Element | Electronic Configuration | Period | Group | Block |
|---|---|---|---|---|
| K | [Ar] 4s¹ | 4 | IA | s |
| S | [Ne] 3s² 3p⁴ | 3 | VIA | p |
| Cu | [Ar] 3d¹⁰ 4s¹ | 4 | IB | d |
Periodicity of Physical Properties
1. Atomic Radius
The atomic radius is the distance from the nucleus to the outermost electron shell.
- Across a Period: Decreases due to the increasing nuclear charge pulling electrons closer
- Example: Na (186 pm) → Cl (99 pm)
- Down a Group: Increases because additional electron shells are added
- Example: F (71 pm) < I (140 pm)
Figure: Graph showing atomic radius trends in Period 3.
2. Ionization Energy (IE)
Ionization energy is the energy required to remove an electron from a gaseous atom.
- Across a Period: Increases (stronger nuclear attraction)
- Example: Na (496 kJ/mol) < Ar (1521 kJ/mol)
- Down a Group: Decreases (outer electrons are farther from the nucleus)
- Example: Li (520 kJ/mol) > Cs (376 kJ/mol)
Anomaly: Boron (B) has a lower IE than beryllium (Be) because removing an electron from Be’s full 2s² orbital requires more energy.
3. Electronegativity
Electronegativity measures an atom’s ability to attract shared electrons in a bond.
- Trend: Increases left to right, decreases top to bottom
- Example: F (4.0) > Cs (0.7)
4. Melting & Boiling Points
- Nonmetals: Increase down a group (stronger van der Waals forces)
- Metals: Decrease down a group (weaker metallic bonds)
- Peak at Group IVA: Due to strong covalent networks (e.g., diamond)
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Figure: Melting point trends in Groups IA, IIA, and VIIA.
5. Electron Affinity
Electron affinity measures the energy change when an atom gains an electron to form a negative ion. It indicates how readily an atom accepts electrons.
- Across a Period:
- Increases (more negative values) due to a higher effective nuclear charge
- Noble gases are exceptions (positive EA as they resist gaining electrons)
- Example: Cl (-349 kJ/mol) > Na (-53 kJ/mol)
- Down a Group:
- Decreases (less negative) because added electron shells shield the nucleus
- Example: F (-328 kJ/mol) > I (-295 kJ/mol)
- Anomalies:
- Group IIA elements have positive EA (full s-orbital resists adding electrons)
- Nitrogen has a lower EA than oxygen due to half-filled p-subshell stability
- Fluorine has a lower EA due to its smaller size and higher electron density than chlorine.
Figure: Electron affinity trends in Group VIIA elements.
6. Electrical Conductivity
Electrical conductivity depends on the availability of mobile electrons that can carry current.
- Metals (Left Side):
- Excellent conductors due to a delocalized “sea of electrons”
- Increases down groups IA/IIA (e.g., Na < K < Rb)
- Transition metals (d-block) show variable conductivity (Cu > Zn)
- Metalloids (Diagonal Band):
- Semiconductors (e.g., Si, Ge) with intermediate conductivity
- Conductivity increases with temperature (unlike metals)
- Nonmetals (Right Side):
- Poor conductors (localized electrons in covalent bonds)
- Exceptions: Graphite (conducts along planes) and doped semiconductors
- Special Cases:
- Carbon: Diamond (insulator) vs. Graphite (conductor)
- Mercury: The only liquid metal at room temperature
| Element | Type | Conductivity (S/m) | Notes |
|---|---|---|---|
| Silver (Ag) | Metal | 6.3×10⁷ | Best conductor |
| Silicon (Si) | Metalloid | 1×10⁻³ | Semiconductor |
| Sulfur (S) | Nonmetal | 5×10⁻²⁸ | Insulator |
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