Chapter 3: Classification of Elements and Periodicity in Properties

3.1 Why Do We Need to Classify Elements?

At present, 118 elements are known. Studying each element individually would be a daunting task. Classification helps us:

• Organize elements systematically
• Predict properties of unknown elements
• Correlate properties of similar elements
• Simplify the study of chemistry

3.2 Brief History of the Development of the Periodic Table

Döbereiner’s Triads (1829)

Döbereiner observed that elements could be grouped in triads where the atomic mass of the middle element was approximately the arithmetic mean of the other two.

Limitation: Could not classify all elements known at the time.

Newlands’ Law of Octaves (1866)

When elements are arranged in order of increasing atomic mass, every eighth element has properties similar to the first — like the octave in music.

Limitation: Valid only up to calcium; failed for heavier elements.

Mendeleev’s Periodic Table (1869)

“The physical and chemical properties of elements are a periodic function of their atomic masses.”

Key Features:

• Arranged 63 elements in order of increasing atomic mass
• Left gaps for undiscovered elements (Eka-aluminium → Gallium, Eka-silicon → Germanium)
• Correctly predicted properties of undiscovered elements

Limitations:

• Position of hydrogen was ambiguous
• Isotopes could not be accommodated
• Some elements with higher atomic mass were placed before those with lower atomic mass (e.g., Co before Ni)

3.3 Modern Periodic Law and the Present Form of the Periodic Table

Modern Periodic Law (Moseley, 1913)

“The physical and chemical properties of elements are a periodic function of their atomic numbers.”

The Long Form of the Periodic Table

The modern periodic table has:

• 7 periods (horizontal rows) — corresponding to the principal quantum number \(n\)
• 18 groups (vertical columns)
• 4 blocks: s, p, d, and f

Nomenclature of Elements with Z > 100

Elements with atomic number greater than 100 are named using the IUPAC nomenclature based on their atomic number:

Example: Element 109 → Un-nil-ennium → Unnilennium (Une) (now named Meitnerium, Mt)

3.4 Periodic Trends in Properties

1. Atomic Radius

Atomic radius is the distance from the centre of the nucleus to the outermost shell of electrons.

Types:

• Covalent radius: Half the internuclear distance between two bonded atoms of the same element
• Van der Waals radius: Half the distance between adjacent atoms of the same element in non-bonded state
• Metallic radius: Half the internuclear distance between adjacent atoms in a metallic crystal

Trends:

2. Ionic Radius

• Cations are smaller than parent atoms (loss of electrons)
• Anions are larger than parent atoms (gain of electrons)
• In an isoelectronic series (same number of electrons), the ion with greater nuclear charge is smaller

Example: Isoelectronic species (all have 10 electrons):

\[ \text{O}^{2-} > \text{F}^- > \text{Ne} > \text{Na}^+ > \text{Mg}^{2+} > \text{Al}^{3+} \]

3. Ionization Enthalpy (IE)

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

\[ \text{X}(g) \rightarrow \text{X}^+(g) + e^- \quad (\Delta_i H) \]

Trends:

• Across a period: Generally increases (due to increasing nuclear charge)
• Down a group: Decreases (outermost electron is farther from nucleus)

Successive ionization enthalpies: \( IE_1 < IE_2 < IE_3 < \ldots \) (each successive electron is harder to remove)

Anomalies:

• \(IE_1\) of B < \(IE_1\) of Be (electron from 2p is easier to remove than from 2s)
• \(IE_1\) of O < \(IE_1\) of N (paired electron in O’s 2p is easier to remove)

4. Electron Gain Enthalpy (\(\Delta_{eg}H\))

The energy change when an electron is added to an isolated gaseous atom:

\[ \text{X}(g) + e^- \rightarrow \text{X}^-(g) \quad (\Delta_{eg}H) \]

• Usually negative (energy is released) for most elements
• Halogens have the most negative values (high tendency to gain electrons)
• Noble gases have positive values (stable electron configuration resists addition)

Trends:

• Across a period: Becomes more negative (more exothermic)
• Down a group: Becomes less negative (exception: Cl has more negative \(\Delta_{eg}H\) than F due to F’s small size and electron-electron repulsion)

5. Electronegativity

The tendency of an atom to attract shared electrons towards itself in a chemical bond.

Pauling Scale: F is the most electronegative element (EN = 4.0)

Trends:

• Across a period: Increases
• Down a group: Decreases

6. Valency

Valency is the combining capacity of an element.

Valency increases from 1 to 4 in Groups 1-14, then decreases from 4 to 0 in Groups 14-18.

Practice Questions

Multiple Choice Questions (MCQs)

1. The element with atomic number 35 belongs to which block of the periodic table?

 (a) s-block

 (b) p-block

 (c) d-block

 (d) f-block

2. Which of the following is the correct order of ionization enthalpy?

 (a) \(\text{B} \gt \text{Be} \gt \text{N} \gt \text{C}\)

 (b) \(\text{Be} \gt \text{B} \gt \text{C} \gt \text{N}\)

 (c) \(\text{N} \gt \text{C} \gt \text{Be} \gt \text{B}\)

 (d) \(\text{N} \gt \text{O} \gt \text{F} \gt \text{Ne}\)

3. In the isoelectronic series \(\text{Na}^+\), \(\text{Mg}^{2+}\), \(\text{Al}^{3+}\), the correct order of ionic radii is:

 (a) \(\text{Na}^+ \gt \text{Mg}^{2+} \gt \text{Al}^{3+}\)

 (b) \(\text{Al}^{3+} \gt \text{Mg}^{2+} \gt \text{Na}^+\)

 (c) \(\text{Mg}^{2+} \gt \text{Na}^+ \gt \text{Al}^{3+}\)

 (d) All have same ionic radii

4. Which of the following has the most negative electron gain enthalpy?

 (a) F

 (b) Cl

 (c) Br

 (d) I

5. The IUPAC name for element with atomic number 112 is:

 (a) Ununbium

 (b) Copernicium

 (c) Unbibium

 (d) Unnilbium

Short Answer Questions (2–3 Marks)

6. What is the modern periodic law? How does it differ from Mendeleev’s periodic law?

7. Explain why the ionization enthalpy of N is more than that of O.

8. Arrange the following in order of increasing atomic radius: C, N, Si, P.

9. What are isoelectronic species? Write four isoelectronic species with 18 electrons each.

10. How would you use the periodic table to predict the formula of a stable compound formed between potassium and sulphur?

Long Answer Questions (5 Marks)

11. (a) Discuss the trends in ionization enthalpy across a period and down a group. Explain any exceptions in Period 2.

 (b) What is electronegativity? Why does fluorine have the highest electronegativity?

12. (a) Explain the process of naming elements with atomic number greater than 100 according to IUPAC nomenclature. Give two examples.

 (b) Which of the following elements would have the highest electron gain enthalpy and why: N, O, F, S?

13. (a) Compare the first ionization enthalpies of Be, B, N, and O. Explain the observed trend with reasons.

 (b) Define ionic radius. How does it vary in an isoelectronic species?

Assertion-Reason Questions

14. Assertion (A): The first ionization enthalpy of Be is greater than that of B.

Reason (R): 2s electrons are more tightly held than 2p electrons.

15. Assertion (A): Noble gases have positive electron gain enthalpy.

Reason (R): Noble gases have stable, completely filled electronic configuration.

Answer Key