Classification of Elements and Periodicity in Properties — Revision Notes

⚡ 30-Second Revision

Modern Periodic Law: — Properties are periodic functions of atomic number.
Periods: — 7 horizontal rows, indicate valence shell $n$.
Groups: — 18 vertical columns, similar valence e- config, similar properties.
Blocks: — s, p, d, f based on differentiating electron.
Atomic Radius: — $\downarrow$ across period, $\uparrow$ down group.
Ionic Radius: — Cation < Parent, Anion > Parent. Isoelectronic: $Z_{eff} \uparrow$, radius $\downarrow$.
Ionization Enthalpy (IE): — Energy to remove e-. $\uparrow$ across period, $\downarrow$ down group. Exceptions: $IE_{Be} > IE_B$, $IE_N > IE_O$.
Electron Gain Enthalpy ($\Delta_{eg}H$): — Energy change on adding e-. More negative across period, less negative down group. Exceptions: $\Delta_{eg}H_{Cl} > \Delta_{eg}H_F$, $\Delta_{eg}H_{S} > \Delta_{eg}H_O$. Noble gases positive.
Electronegativity (EN): — Tendency to attract shared e-. $\uparrow$ across period, $\downarrow$ down group. F is highest (4.0).
Metallic Character: — Tendency to lose e-. $\downarrow$ across period, $\uparrow$ down group.
Lanthanoid Contraction: — Poor shielding of 4f e- leads to smaller 5d radii, similar to 4d.

2-Minute Revision

The Modern Periodic Table arranges elements by increasing atomic number, leading to periodic repetition of properties. It has 7 periods (valence shell number) and 18 groups (similar valence electron configuration).

Elements are categorized into s, p, d, and f blocks. Key periodic properties include atomic radius (decreases across a period, increases down a group), ionic radius (cations are smaller, anions larger than parent atoms; for isoelectronic species, size decreases with increasing nuclear charge).

Ionization enthalpy, the energy to remove an electron, generally increases across a period and decreases down a group, with important exceptions like Group 2 vs 13 and Group 15 vs 16 due to orbital stability.

Electron gain enthalpy, the energy change on adding an electron, generally becomes more negative across a period and less negative down a group. Crucial exceptions include chlorine having a more negative electron gain enthalpy than fluorine due to electron-electron repulsion.

Electronegativity, the tendency to attract shared electrons, increases across a period and decreases down a group, with fluorine being the most electronegative. Metallic character, the tendency to lose electrons, decreases across a period and increases down a group.

Remember Lanthanoid Contraction due to poor 4f shielding, causing similar radii for 4d and 5d elements.

5-Minute Revision

The foundation of inorganic chemistry lies in the Modern Periodic Table, which organizes elements based on their atomic number, leading to predictable periodic trends in their properties. This arrangement, proposed by Moseley, rectified the anomalies of Mendeleev's atomic mass-based table.

The table features 7 periods, corresponding to the principal quantum number of the valence shell, and 18 groups, where elements share similar valence electron configurations and thus similar chemical properties.

Elements are further classified into s, p, d, and f blocks based on the subshell being filled.

Understanding periodic trends is crucial. Atomic radius generally decreases across a period due to increasing effective nuclear charge and increases down a group due to the addition of new electron shells. Ionic radius follows similar trends, but cations are always smaller than their parent atoms, and anions are larger. For isoelectronic species (same electron count), the size decreases with increasing nuclear charge (e.g., $N^{3-} > O^{2-} > F^{-} > Na^{+}$).

Ionization enthalpy (IE), the energy required to remove an electron, generally increases across a period and decreases down a group. Crucial exceptions exist: $IE_1$ of Group 2 elements is higher than Group 13 (e.g., $Be > B$) due to stable $s^2$ configuration, and $IE_1$ of Group 15 is higher than Group 16 (e.g., $N > O$) due to stable half-filled $p^3$ configuration.

**Electron gain enthalpy ($\Delta_{eg}H$)**, the energy change when an electron is added, generally becomes more negative (exothermic) across a period and less negative (or more positive) down a group.

A significant exception is that chlorine has a more negative $\Delta_{eg}H$ than fluorine, and sulfur more than oxygen, attributed to the high electron density and inter-electronic repulsion in the smaller 2p orbitals of F and 2p orbitals of O.

Noble gases have positive $\Delta_{eg}H$ due to their stable, filled shells.

Electronegativity, the tendency to attract shared electrons in a bond, increases across a period and decreases down a group. Fluorine is the most electronegative element. Metallic character (tendency to lose electrons) decreases across a period and increases down a group, while non-metallic character shows the opposite trend.

Finally, remember Lanthanoid Contraction, the gradual decrease in atomic/ionic radii across the 4f series due to poor shielding by 4f electrons. This leads to similar sizes and properties for 4d and 5d transition elements, a frequently tested consequence.

Prelims Revision Notes

Classification of Elements and Periodicity in Properties - NEET Revision Notes

1. Modern Periodic Law: States that physical and chemical properties of elements are periodic functions of their atomic numbers. This is the basis of the Modern Periodic Table.

2. Structure of Modern Periodic Table:

* Periods (7): Horizontal rows. Period number = principal quantum number ($n$) of valence shell. Elements in a period have different properties. * Groups (18): Vertical columns. Elements in a group have similar outer electronic configurations and thus similar chemical properties. * Blocks: s-block (Groups 1, 2), p-block (Groups 13-18), d-block (Groups 3-12), f-block (Lanthanoids, Actinoids).

3. Periodic Trends in Properties:

* Atomic Radius: * Across a period (L to R): Decreases (due to increasing $Z_{eff}$). Example: $Li > Be > B > C$. * Down a group (T to B): Increases (due to addition of new shells). Example: $F < Cl < Br < I$.

* Ionic Radius: * Cation < Parent atom ($Na^+ < Na$). Anion > Parent atom ($Cl^- > Cl$). * Isoelectronic Species: For same electron count, radius decreases with increasing nuclear charge. Example: $N^{3-} > O^{2-} > F^{-} > Na^{+} > Mg^{2+} > Al^{3+}$.

* Ionization Enthalpy (IE): Energy to remove an electron. * Across a period: Increases (due to increasing $Z_{eff}$, decreasing size). * Down a group: Decreases (due to increasing size, shielding).

* Exceptions: $IE_{Be} > IE_B$ (stable $2s^2$). $IE_N > IE_O$ (stable half-filled $2p^3$). * **Electron Gain Enthalpy ($\Delta_{eg}H$):** Energy change on adding an electron. * Across a period: Becomes more negative (more exothermic).

* Down a group: Becomes less negative (less exothermic/more positive). * Exceptions: $\Delta_{eg}H_{Cl} > \Delta_{eg}H_F$ (F's small size causes e-e repulsion). $\Delta_{eg}H_{S} > \Delta_{eg}H_O$.

Noble gases have positive $\Delta_{eg}H$. * Electronegativity (EN): Tendency to attract shared electrons. * Across a period: Increases. (F is highest, 4.0). * Down a group: Decreases. * Metallic Character: Tendency to lose electrons.

* Across a period: Decreases. * Down a group: Increases. * Non-metallic Character: Tendency to gain electrons. * Across a period: Increases. * Down a group: Decreases.

4. Lanthanoid Contraction: Gradual decrease in atomic/ionic radii from La to Lu. Caused by poor shielding of 4f electrons. Consequence: 4d and 5d transition elements in the same group have very similar radii (e.g., Zr and Hf) and properties.

Vyyuha Quick Recall

To remember the trends of atomic radius, ionization enthalpy, and electronegativity across a period and down a group, think of 'FINE':

Fluorine is the most electronegative, smallest, and highest IE element (in its period).

Increases Nuclear Effective charge across a period, so:

Atomic Radius $\downarrow$
Ionization Enthalpy $\uparrow$
Electronegativity $\uparrow$

Down a Group, Shells Add, so:

Atomic Radius $\uparrow$
Ionization Enthalpy $\downarrow$
Electronegativity $\downarrow$