Chemistry: Class 11 — Chapter 3: Classification of Elements and Periodicity in Properties

• Dobereiner's Triads (1829): Grouped elements into sets of three (Triads) where the atomic weight of the middle element was approximately the arithmetic mean of the other two (e.g., $Li$, $Na$, $K$).
• Newlands' Law of Octaves (1865): Arranged elements in increasing order of atomic weights, noting that every eighth element exhibited similar properties, analogous to musical octaves. This law was valid only up to Calcium.
• Lothar Meyer (1868): Plotted physical properties (atomic volume, melting point, boiling point) against atomic weight, obtaining a periodically repeating pattern.

Dmitri Mendeleev proposed that the properties of elements are a periodic function of their atomic weights. He arranged elements in horizontal rows (periods) and vertical columns (groups).

• Key Achievements: Left gaps for undiscovered elements (e.g., Eka-Aluminium and Eka-Silicon, later discovered as Gallium and Germanium) and predicted their properties with remarkable accuracy.
• Anomalies: Occasionally ignored strict atomic weight order to group elements with similar properties together (e.g., Iodine placed after Tellurium).

In 1913, Henry Moseley showed that atomic number ($Z$) is a more fundamental property than atomic mass, based on regularities in characteristic X-ray spectra ($\sqrt{\nu}$ vs $Z$). This led to the formulation of the Modern Periodic Law:

"The physical and chemical properties of the elements are periodic functions of their atomic numbers."

• Periods: 7 horizontal rows. The period number corresponds to the highest principal quantum number ($n$) of the elements in that period.
• Groups: 18 vertical columns numbered 1 to 18 (replacing older IA-VIIB notations). Elements in a group share similar outer electronic configurations.

Comprises Group 1 (alkali metals, $ns^1$) and Group 2 (alkaline earth metals, $ns^2$). They are highly reactive metals with low ionization enthalpies, forming $1+$ and $2+$ cations respectively.

Comprises Groups 13 to 18 ($ns^2 np^{1-6}$). Together with the s-block, they are known as Representative Elements. Group 17 elements are halogens, Group 16 are chalcogens, and Group 18 are noble gases.

Comprises Groups 3 to 12, characterized by the filling of inner d-orbitals with a general configuration of $(n-1)d^{1-10} ns^{0-2}$. They are all metals, typically forming colored ions, exhibiting variable oxidation states, and showing paramagnetism.

Consists of two series at the bottom of the table: Lanthanoids ($Z=58-71$) and Actinoids ($Z=90-103$), with outer configuration $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$. They are all metals; actinoids are radioactive.

• Metals: Comprise >78% of elements, located on the left. They are malleable, ductile, and good conductors.
• Non-metals: Located on the top right, brittle, poor conductors of heat and electricity.
• Metalloids (Semi-metals): Elements bordering the diagonal divide (e.g., $Si$, $Ge$, $As$, $Sb$) showing mixed properties.

• Atomic Radius: Decreases across a period due to increasing effective nuclear charge pulling electrons closer. Increases down a group due to the addition of new principal energy shells which shield outer electrons.
• Covalent vs. Metallic Radius: Covalent radius is half the distance between two single-bonded atoms; metallic radius is half the internuclear distance between adjacent metal cores in a crystal.
• Ionic Radius: Cations are smaller than their parent atoms due to loss of electrons and increased effective nuclear charge. Anions are larger due to increased electron-electron repulsion.
• Isoelectronic Species: Atoms/ions with the same number of electrons (e.g., $O^{2-}$, $F^-$, $Na^+$, $Mg^{2+}$). Size decreases as nuclear charge increases.

The energy required to remove an electron from an isolated gaseous atom in its ground state:

$$X(g) \rightarrow X^+(g) + e^-$$

• Trends: Generally increases across a period (increasing nuclear charge) and decreases down a group (increasing atomic size and shielding).
• Exceptions: Boron has lower $\Delta_i H$ than Beryllium because removing a $2p$ electron is easier than a stable $2s$ electron. Oxygen has lower $\Delta_i H$ than Nitrogen due to inter-electronic repulsion in the doubly occupied $2p$ orbital of oxygen.

The enthalpy change when an electron is added to a neutral gaseous atom:

$$X(g) + e^- \rightarrow X^-(g)$$

• Trends: Becomes more negative across a period and less negative down a group.
• Exceptions: Halogens have highly negative values. Noble gases have large positive values. Chlorine has a more negative $\Delta_{eg} H$ than Fluorine because the compact $2p$ shell of Fluorine experiences significant inter-electronic repulsion.

The qualitative measure of an atom's ability to attract shared electrons in a chemical compound. It increases across a period and decreases down a group. Fluorine is the most electronegative element (assigned 4.0 on the Pauling scale).

Valence is determined by the number of valence electrons. For representative elements, it is usually equal to the number of valence electrons or $8$ minus that number. Oxidation state represents the charge acquired by an atom based on electronegativity differences in a molecule.

The first element of each group (Lithium, Beryllium, Boron to Fluorine) differs significantly from heavier congeners due to:

• Small atomic/ionic size
• High charge-to-size ratio
• High electronegativity
• Absence of d-orbitals (limiting maximum covalency to 4)
• Ability to form $p\pi-p\pi$ multiple bonds

Diagonal Relationship: Lithium resembles Magnesium, and Beryllium resembles Aluminium due to similar ionic sizes and charge/radius ratios.

• Reactivity is highest at the extremes of a period (highly electropositive metals on the left, highly electronegative non-metals on the right) and lowest in the center.
• Oxide Trends: Leftmost elements form strongly basic oxides (e.g., $Na_2O$), rightmost elements form strongly acidic oxides (e.g., $Cl_2O_7$), and central elements form amphoteric (e.g., $Al_2O_3$) or neutral oxides (e.g., $CO$, $NO$, $N_2O$).