1. Why Classify Elements?
- Elements are the basic units of matter.
- The number of known elements increased significantly over time: 31 in 1800, 63 by 1865, and currently 114, with ongoing efforts to synthesize new ones.
- It's challenging to study each element and its numerous compounds individually.
- Scientists sought a systematic way to organize knowledge, rationalize known chemical facts, and predict new ones.
2. Genesis of Periodic Classification
Johann Dobereiner (early 1800s)
First to observe trends. He noted similarities in physical and chemical properties within groups of three elements, called triads. The middle element in a triad had an atomic weight approximately halfway between the other two, and its properties were intermediate. His "law of triads" was dismissed as coincidence as it worked for only a few elements.
A.E.B. de Chancourtois (1862)
Arranged elements by increasing atomic weights on a cylindrical table, observing periodic recurrence of properties. This did not gain much attention.
John Alexander Newlands (1865) - Law of Octaves
He arranged elements in increasing order of atomic weights and noted that every eighth element had properties similar to the first, analogous to musical octaves. The law was only true for elements up to calcium. He was later awarded the Davy Medal in 1887.
Dmitri Mendeleev & Lothar Meyer (1869)
Working independently, they proposed that arranging elements by increasing atomic weights revealed similarities in physical and chemical properties at regular intervals. Lothar Meyer plotted physical properties (atomic volume, melting point, boiling point) against atomic weight, finding a periodically repeated pattern. His work was published after Mendeleev's.
Mendeleev's Periodic Law:
"The properties of the elements are a periodic function of their atomic weights".
Mendeleev's periodic law gained prominence because he:
- Used a broader range of properties and recognized the significance of periodicity.
- Sometimes ignored the strict order of atomic weights to place elements with similar properties together (e.g., Iodine before Tellurium).
- Left gaps for undiscovered elements and predicted their properties with remarkable accuracy (e.g., Eka-aluminium for Gallium, Eka-silicon for Germanium).
- His bold and successful quantitative predictions made his Periodic Table famous.
3. Modern Periodic Law and the Present Form of the Periodic Table
In 1913, Henry Moseley's work on X-ray spectra showed that the atomic number (Z), not atomic mass, is the fundamental property of an element. A plot of the square root of X-ray frequency versus atomic number yielded a straight line.
Modern Periodic Law:
"The physical and chemical properties of the elements are periodic functions of their atomic numbers."
Structure of the Modern Periodic Table:
- Periods: 7 horizontal rows. The period number corresponds to the highest principal quantum number (n).
- Groups: 18 vertical columns. Elements in a group have similar outer electronic configurations and thus similar properties.
- Lanthanoids (4f) and Actinoids (5f): Placed at the bottom to maintain the table's structure. This reconfiguration was influenced by Glenn T. Seaborg's discovery of transuranium elements. Element 106 is named Seaborgium (Sg) in his honor.
4. Nomenclature of Elements with Atomic Numbers > 100
To avoid naming controversies, IUPAC established a systematic temporary nomenclature based on numerical roots for the atomic number (e.g., un-nil-quadium for element 104). Once discovery is verified, a permanent name is assigned. Elements up to 118 have been discovered and officially named.
5. Electronic Configurations and the Periodic Table
An element's position reflects the quantum numbers of its last filled orbital.
Configurations in Periods:
Each period marks the filling of a new principal energy level (n). The number of elements in a period is twice the number of available atomic orbitals.
Groupwise Configurations:
Elements in the same group share similar valence shell electron configurations, leading to similar chemical behaviors. Exceptions: Hydrogen is placed separately, and Helium (s-block) is placed with noble gases (p-block) due to its stable, filled shell.
6. Electronic Configurations and Types of Elements: s-, p-, d-, f- Blocks
s-Block (Groups 1-2)
Alkali and alkaline earth metals. Highly reactive with low ionization enthalpies. Form +1 or +2 ions. Their compounds are predominantly ionic (except Li and Be).
p-Block (Groups 13-18)
Includes metals, non-metals, and metalloids (collectively called Representative Elements with s-block). Contains noble gases (stable) and halogens (highly reactive). Non-metallic character increases left to right; metallic character increases down a group.
d-Block (Groups 3-12) - Transition Metals
Characterized by filling inner d-orbitals. They are all metals, form colored ions, have variable oxidation states, are paramagnetic, and are often used as catalysts. They bridge the s-block and p-block elements. Zn, Cd, and Hg are exceptions and don't show most transition properties.
f-Block - Inner-Transition Metals
Lanthanoids and Actinoids. Characterized by filling f-orbitals. All are metals with similar properties within each series. Actinoids are radioactive.
7. Metals, Non-metals, and Metalloids
- Metals: >78% of elements, on the left side. Good conductors, malleable, ductile, high melting points.
- Non-metals: Top-right of the table. Poor conductors, brittle, low melting points.
- Metalloids: Border the zig-zag line. Have properties intermediate between metals and non-metals (e.g., Silicon, Germanium).
8. Periodic Trends in Physical Properties
Atomic Radius
Decreases across a period (due to increased effective nuclear charge). Increases down a group (due to addition of new energy shells and increased shielding).
Ionic Radius
Cations are smaller than their parent atoms. Anions are larger. For isoelectronic species, radius decreases as nuclear charge increases.
Ionization Enthalpy (ΔiH)
Energy to remove an electron. Increases across a period. Decreases down a group. Anomalies exist (e.g., B vs Be, O vs N) due to shielding and electron-electron repulsion in filled/half-filled orbitals.
Electron Gain Enthalpy (ΔegH)
Energy change when an electron is added. Becomes more negative across a period. Becomes less negative down a group. Anomalies exist (e.g., F vs Cl) due to electron repulsion in smaller atoms.
Electronegativity
Ability to attract shared electrons. Increases across a period. Decreases down a group. Directly related to non-metallic properties.
9. Periodic Trends in Chemical Properties
Valence/Oxidation States
For representative elements, valence is often related to the group number. Transition metals show variable oxidation states.
Anomalous Properties of Second Period Elements
The first element of each group shows unique behavior due to small size, high electronegativity, and lack of d-orbitals (limiting covalency to 4). This leads to diagonal relationships (e.g., Li resembles Mg).
Chemical Reactivity
- Across a period: Reactivity is highest at the extremes (alkali metals and halogens) and lowest in the center.
- Down a group (metals): Reactivity increases.
- Down a group (non-metals): Reactivity decreases.
- Oxides: Trend from basic (left) → amphoteric (middle) → acidic (right) across a period.