Exercise Solutions Available
Step-by-step answers for all textbook questions in this chapter
How to Use This Page
Welcome to the exercise solutions for Chapter 3 — Periodic Classification of Elements. Here is how to get the most from this page:
1. Attempt every question yourself first before reading the solution. 2. Compare your answer with the detailed solution and explanation provided. 3. Use the Key Concepts tables for quick revision before exams. 4. Memorise the periodic trends — they are frequently tested in board exams. 5. Check the Exam Checklist at the end to confirm you have covered all important topics.
Key Concepts (Quick Recap)
Modern Periodic Law and Table
The modern periodic law states: The physical and chemical properties of elements are periodic functions of their atomic numbers. The modern periodic table has 7 periods and 18 groups.
Periodic Trends Summary
| Property | Across a Period (L → R) | Down a Group |
|---|---|---|
| Atomic radius | Decreases | Increases |
| Ionic radius | Decreases | Increases |
| Ionisation energy | Increases | Decreases |
| Electron affinity | Generally increases | Generally decreases |
| Electronegativity | Increases | Decreases |
| Metallic character | Decreases | Increases |
| Non-metallic character | Increases | Decreases |
| Reactivity of metals | Decreases | Increases |
| Reactivity of non-metals | Increases | Decreases |
Blocks of the Periodic Table
| Block | Orbitals Filled | Groups | Key Features |
|---|---|---|---|
| s-block | ns¹ to ns² | 1, 2 | Highly reactive metals, low ionisation energy |
| p-block | ns² np¹ to ns² np⁶ | 13–18 | Metals, metalloids, non-metals |
| d-block | (n-1)d¹ to (n-1)d¹⁰ | 3–12 | Transition metals, variable oxidation states |
| f-block | (n-2)f¹ to (n-2)f¹⁴ | Separate rows | Lanthanides and actinides, radioactive (most actinides) |
Types of Atomic Radius
| Type | Definition |
|---|---|
| Covalent radius | Half the distance between nuclei of two bonded atoms |
| Metallic radius | Half the distance between nuclei of two adjacent metal atoms |
| Van der Waals radius | Half the distance between nuclei of two non-bonded atoms in adjacent molecules |
Ionisation Energy — Key Points
- First ionisation energy (IE₁): Energy to remove the first electron.
- Second ionisation energy (IE₂): Energy to remove the second electron. IE₂ > IE₁ always.
- Successive ionisation energies increase because the remaining electrons are held more tightly.
- Exceptions: Be → B (drop due to 2p electron being easier to remove than 2s); N → O (drop due to 2p⁴ pairing repulsion).
Electron Affinity
- Definition: Energy released when an electron is added to a neutral gaseous atom.
- Halogens have the highest electron affinity (most negative values).
- Noble gases have positive electron affinity (unfavourable).
- Fluorine has lower electron affinity than chlorine due to its small size and high electron-electron repulsion.
I. Choose the Best Answer (MCQs)
II. Fill in the Blanks
1. The modern periodic law states that properties of elements are periodic functions of their atomic number.
2. The number of periods in the modern periodic table is 7.
3. The element with the lowest ionisation energy in Group 1 is Cs (Caesium).
4. The d-block elements are also called transition elements.
5. Lanthanides and actinides are placed in the f-block of the periodic table.
6. The atomic radius of sodium is greater than that of magnesium (Na > Mg across a period).
7. The most electronegative element is Fluorine with a value of 4.0 on the Pauling scale.
8. The electron affinity of noble gases is positive (unfavourable to add an electron).
9. Ionisation energy generally increases across a period from left to right.
10. The metallic character decreases across a period and increases down a group.
11. The elements of Group 17 are called halogens.
12. The first period contains 2 elements.
13. The number of elements in the 6th period is 32.
14. The covalent radius is half the internuclear distance between two bonded identical atoms.
15. The element with the highest electron affinity in the periodic table is Chlorine.
III. True or False (with Corrections)
1. Statement: Atomic radius increases across a period from left to right.
2. Statement: Noble gases have the highest ionisation energy in their respective periods.
3. Statement: Electronegativity and metallic character show the same trend in the periodic table.
4. Statement: All transition elements are metals.
5. Statement: Electron affinity becomes less negative going down a group.
IV. Match the Following
| Column A | Column B |
|---|---|
| 1. Alkali metals | (a) Group 17 elements |
| 2. Halogens | (b) Group 1 elements |
| 3. Noble gases | (c) Group 2 elements |
| 4. Alkaline earth metals | (d) Group 18 elements |
| 5. Lanthanides | (e) 4f block elements |
| 6. Transition metals | (f) d-block, Groups 3–12 |
1 — (b): Alkali metals are Group 1 elements (Li, Na, K, Rb, Cs, Fr) 2 — (a): Halogens are Group 17 elements (F, Cl, Br, I, At) 3 — (d): Noble gases are Group 18 elements (He, Ne, Ar, Kr, Xe, Rn) 4 — (c): Alkaline earth metals are Group 2 elements (Be, Mg, Ca, Sr, Ba, Ra) 5 — (e): Lanthanides are the 4f block elements (Ce to Lu) 6 — (f): Transition metals are d-block elements in Groups 3–12
V. Answer Briefly (2-mark)
VI. Answer in Detail (5-mark)
Atomic Radius
Across a period (decreases):
- Nuclear charge increases by one unit for each successive element.
- Electrons are added to the same principal energy level.
- Increased nuclear charge pulls the electron cloud closer to the nucleus.
- Shielding effect remains nearly constant within the same shell.
Down a group (increases):
- New electron shells are added for each successive element.
- The principal quantum number (n) increases.
- Outermost electrons are farther from the nucleus.
- Shielding effect increases, reducing effective nuclear charge on outer electrons.
Ionisation Energy
Across a period (increases):
- Nuclear charge increases while atomic radius decreases.
- Outermost electrons are held more tightly.
- More energy is needed to remove an electron.
Exceptions:
- Be (Z=4) → B (Z=5): IE drops because B's 2p electron is easier to remove than Be's paired 2s electrons.
- N (Z=7) → O (Z=8): IE drops because O has a paired 2p electron (2p⁴) which experiences repulsion.
Down a group (decreases):
- Atomic radius increases significantly.
- Outermost electrons are farther from the nucleus.
- Shielding effect increases.
- Less energy is needed to remove an electron.
Electronegativity
Across a period (increases):
- Increasing nuclear charge and decreasing atomic radius.
- Atoms become more effective at attracting bonding electrons.
Down a group (decreases):
- Increasing atomic radius.
- Bonding electrons are farther from the nucleus.
- Shielding effect reduces the effective nuclear charge.
Summary Table:
s-Block Elements (Groups 1 and 2)
- Last electron enters the s-orbital.
- Group 1 (Alkali metals): Li, Na, K, Rb, Cs, Fr
- Group 2 (Alkaline earth metals): Be, Mg, Ca, Sr, Ba, Ra
Properties:
- Highly reactive metals with low ionisation energies.
- Form ionic compounds by losing electrons (M → M⁺ or M → M²⁺).
- Strong reducing agents.
- React vigorously with water (alkali metals explode).
- Oxides are basic.
- Flame colours: Na (yellow), K (violet), Ca (brick red), Sr (crimson), Ba (green).
p-Block Elements (Groups 13–18)
- Last electron enters the p-orbital.
- Includes metals (Al, Sn, Pb, Tl), metalloids (B, Si, Ge, As, Sb, Te), and non-metals (C, N, O, F, Cl, etc.).
- Noble gases (Group 18) have complete octets.
Properties:
- Wide range of properties — from metals to non-metals.
- Form covalent compounds.
- Oxides can be acidic, basic, or amphoteric.
- Include the most electronegative elements (F, O, Cl, N).
d-Block Elements (Groups 3–12, Transition Metals)
- Last electron enters the (n-1)d orbital.
- Three series: 3d (Sc–Zn), 4d (Y–Cd), 5d (La–Hg).
Properties:
- Hard, dense metals with high melting and boiling points.
- Variable oxidation states due to participation of (n-1)d electrons.
- Form coloured compounds and ions.
- Act as catalysts (Fe, Ni, Pt, V₂O₅).
- Form complex compounds (coordination compounds).
- Show paramagnetism due to unpaired d electrons.
f-Block Elements (Lanthanides and Actinides)
- Last electron enters the (n-2)f orbital.
- Lanthanides (4f): Ce (Z=58) to Lu (Z=71) — placed in the 6th period.
- Actinides (5f): Th (Z=90) to Lr (Z=103) — placed in the 7th period.
Properties:
- Lanthanides are also called rare earth elements.
- All actinides after uranium (Z=92) are transuranic — artificially prepared and radioactive.
- Show limited variable oxidation states (mostly +3).
- Lanthanide contraction: steady decrease in atomic size across the 4f series.
Definition
Electron affinity is the energy released when an electron is added to a neutral gaseous atom to form a negative ion:
It is expressed in kJ/mol. A more negative (or more exothermic) value means greater tendency to accept an electron.
Trend Across a Period
Electron affinity generally becomes more negative (more energy released) across a period from left to right because:
- Nuclear charge increases, so the atom attracts the added electron more strongly.
- Atomic radius decreases, so the electron is closer to the nucleus.
Trend Down a Group
Electron affinity generally becomes less negative (less energy released) down a group because:
- Atomic radius increases, so the added electron is farther from the nucleus.
- Shielding effect increases, reducing effective nuclear charge on the added electron.
Anomalies
Comparative Table
Common Mistakes to Avoid
1. Confusing ionisation energy and electron affinity: Ionisation energy is about removing an electron (energy needed); electron affinity is about adding an electron (energy released). They are opposite processes.
2. Assuming fluorine has the highest electron affinity: Fluorine has the highest electronegativity, but chlorine has a higher (more negative) electron affinity due to fluorine's small size causing electron repulsion.
3. Forgetting exceptions in ionisation energy trends: Be → B and N → O drops in ionisation energy are important exceptions that are frequently tested.
4. Mixing up metallic and non-metallic character trends: Metallic character increases down and decreases across. Non-metallic character does the opposite. Do not reverse these.
5. Treating all d-block elements as typical transition metals: Zn, Cd, and Hg have a complete d¹⁰ configuration and do not show typical transition metal properties like variable oxidation states.
6. Ignoring shielding effect when explaining trends: The shielding effect is crucial for understanding why ionisation energy decreases down a group despite increasing nuclear charge.
7. Confusing isoelectronic species comparisons: For isoelectronic species, the one with the higher nuclear charge has the smaller radius and higher ionisation energy. Always compare nuclear charge, not just the number of electrons.
8. Forgetting that electron affinity can be positive: Noble gases and some other elements have positive electron affinity, meaning energy must be supplied to add an electron. This is not an anomaly — it reflects the stability of their configuration.
Exam Checklist
- [ ] Can state the modern periodic law and explain the arrangement of elements
- [ ] Can explain trends in atomic radius across a period and down a group
- [ ] Can explain trends in ionisation energy with exceptions (Be→B, N→O)
- [ ] Can compare electron affinity of fluorine and chlorine and explain the anomaly
- [ ] Can define electronegativity and state the most electronegative element
- [ ] Can describe s, p, d, f blocks with examples and characteristic properties
- [ ] Can compare metallic and non-metallic character trends
- [ ] Can explain isoelectronic species and their radius comparison
- [ ] Can define and differentiate covalent, metallic, and Van der Waals radii
- [ ] Can explain why noble gases have the highest ionisation energy in each period