This chapter Class 11 Chemistry Chap 3 Classification of Element introduces students to the systematic arrangement of elements in the periodic table based on their atomic numbers. It explains how classification helps predict physical and chemical properties, and highlights key contributions from scientists like Mendeleev and Moseley.
Through concepts like periods, groups, periodic trends (atomic size, ionization enthalpy, electron gain enthalpy, electronegativity), and terms like isoelectronic species and metallic character, the chapter builds a strong foundation for understanding element behavior across the table. Solving its questions helps reinforce these principles and prepares students for further studies in chemistry.
Class 11 Chemistry Chap 3 Classification of Element Textbook Solution
Q1. What is the basic theme of organization in the periodic table?
Answer:
The basic theme of the periodic table is to arrange elements in order of increasing atomic number so that elements with similar chemical properties fall into the same vertical columns (groups). This arrangement helps to understand trends in properties such as atomic size, ionization energy, and electronegativity.
Q2. Which important property did Mendeleev use to classify the elements in his periodic table and did he stick to that?
Answer:
Mendeleev used atomic mass as the basis for classification. However, he did not strictly stick to it—he placed some elements out of order to group them with chemically similar elements, showing he prioritized chemical properties over strict order of atomic masses.
Q3. What is the basic difference in approach between the Mendeleev’s Periodic Law and the Modern Periodic Law?
Answer:
- Mendeleev’s Periodic Law: Properties of elements are a periodic function of their atomic masses.
- Modern Periodic Law: Properties of elements are a periodic function of their atomic numbers.
Q4. On the basis of quantum numbers, justify that the sixth period of the periodic table should have 32 elements.
Answer:
In the sixth period, the principal quantum number n = 6.
Orbitals filled include:
- 6s → 2 electrons
- 4f → 14 electrons
- 5d → 10 electrons
- 6p → 6 electrons
Total = 2 + 14 + 10 + 6 = 32 elements
Q5. In terms of period and group, where would you locate the element with Z = 114?
Answer:
Z = 114 → belongs to 7th period
Expected to be a part of Group 14 (like Carbon, Silicon, Lead, etc.)
Location: 7th Period, Group 14
Q6. Write the atomic number of the element present in the third period and seventeenth group of the periodic table.
Answer:
Third period: elements with atomic numbers 11 to 18
Group 17 (halogens): in third period → Chlorine
Atomic number = 17
Q7. Which element do you think would have been named by: (i) Lawrence Berkeley Laboratory (ii) Seaborg’s group?
Answer:
(i) Lawrencium (Z = 103) – named after Lawrence Berkeley Laboratory
(ii) Seaborgium (Z = 106) – named after Glenn T. Seaborg
Q8. Why do elements in the same group have similar physical and chemical properties?
Answer:
Because they have the same number of valence electrons, which determine chemical reactivity and bonding behavior.
Q9. What does atomic radius and ionic radius really mean to you?
Answer:
- Atomic radius: Distance from the nucleus to the outermost shell of a neutral atom.
- Ionic radius: Radius of an atom’s ion. Cations are smaller (loss of electrons), anions are larger (gain of electrons) than the neutral atom.
Q10. How do atomic radius and ionic radius vary across a period and down a group?
Answer:
- Across a period: Both atomic and ionic radii decrease due to increased nuclear charge pulling electrons closer.
- Down a group: Both radii increase due to addition of new electron shells.
Q11. What is the significance of the terms: – (i) Isoelectronic species, (ii) Isoelectronic series?
Answer:
(i) Isoelectronic species: Different atoms/ions with the same number of electrons.
Example: Na⁺, Mg²⁺, Ne → all have 10 electrons.
(ii) Isoelectronic series: A group of isoelectronic species arranged in order of increasing nuclear charge (Z).
Example: N³⁻ < O²⁻ < F⁻ < Ne < Na⁺ < Mg²⁺
Q12. Consider the following species: N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ and Al³⁺
Arrange them in the order of increasing ionic radii.
Answer:
All are isoelectronic (10 electrons).
As nuclear charge increases, size decreases.
Order: Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻
Q13. Explain why cations are smaller and anions are larger in radii than their parent atoms.
Answer:
- Cations lose electrons → fewer electron–electron repulsions and increased nuclear pull → smaller radius
- Anions gain electrons → more repulsion among electrons → larger radius
Q14. What is the trend in ionization enthalpy across a period and down a group?
Answer:
- Across a period: Ionization enthalpy increases due to higher nuclear charge and smaller atomic size.
- Down a group: Ionization enthalpy decreases due to increase in atomic size and shielding effect.
Q15. How would you explain the fact that the first ionization enthalpy of Na is lower than that of Mg but its second ionization enthalpy is much higher than that of Mg?
Answer:
- Na has 1 valence electron → easy to remove → low first ionization enthalpy
- After losing 1 electron, Na⁺ has stable noble gas configuration → removing another electron requires high energy → very high second ionization enthalpy
- Mg has 2 valence electrons → both first and second ionizations are comparatively easier
Q16. What is the basic difference between the terms electron gain enthalpy and electronegativity?
Answer:
- Electron gain enthalpy: Energy released or absorbed when an atom gains an electron in the gaseous state. It is a thermodynamic property.
- Electronegativity: Tendency of an atom to attract shared electrons in a chemical bond. It is a relative and dimensionless property.
Q17. How would you explain the fact that the electron gain enthalpy of fluorine is less negative than that of chlorine?
Answer:
Though fluorine is more electronegative, its small size leads to strong electron–electron repulsion in the compact 2p orbital.
Chlorine is larger, so added electron feels less repulsion, making electron gain enthalpy of Cl more negative than F.
Q18. Explain with examples: (i) Alkali metals have lowest ionization enthalpy in their periods. (ii) Noble gases have highest ionization enthalpy.
Answer:
(i) Alkali metals (e.g., Na, K) have a single loosely held valence electron. Due to large size and shielding, it’s easily removed → lowest ionization enthalpy.
(ii) Noble gases (e.g., Ne, Ar) have fully filled, stable electron configurations. High nuclear charge and symmetry resist removal of electrons → highest ionization enthalpy.
Q19. How would you explain the higher metallic character of an element in a group as compared to its period?
Answer:
- Down a group: Atomic size increases, ionization energy decreases → more metallic character
- Across a period: Ionization energy increases → metallic character decreases
So, elements down a group are more metallic than those across a period.
Q20. What are the major differences between metals and non-metals?
Answer:
| Property | Metals | Non-metals |
|---|---|---|
| 1. Valence electrons | 1–3 | 4 or more |
| 2. Ionization energy | Low | High |
| 3. Electron gain | Lose electrons to form cations | Gain electrons to form anions |
| 4. Conductivity | Good conductors of heat/electricity | Poor conductors |
| 5. Nature | Lustrous, malleable, ductile | Non-lustrous, brittle |
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Class 11 Chemistry Chap 3 Classification of Element helps students understand how elements are organized in a meaningful and logical way. By studying periodic trends such as atomic radius, ionization enthalpy, and electronegativity, learners can predict the behavior of elements across the periodic table.
The questions and solutions covered in Class 11 Chemistry Chap 3 Classification of Element strengthen conceptual clarity and problem-solving skills. Overall, mastering Class 11 Chemistry Chap 3 Classification of Element lays the groundwork for advanced topics in chemical bonding, reactivity, and periodic properties. Regular practice of Class 11 Chemistry Chap 3 Classification of Element ensures a strong command over this fundamental topic.
Class 11 Chemistry Chap 3 Classification of Element provides a scientific framework for understanding how and why elements behave the way they do. By organizing elements based on increasing atomic number and recurring chemical properties, the periodic table becomes a powerful tool for predicting reactivity, stability, and bonding patterns. Throughout Class 11 Chemistry Chap 3 Classification of Element, students explore key periodic trends such as atomic and ionic radii, ionization enthalpy, electron gain enthalpy, and electronegativity, which are essential for understanding chemical behavior.