Class 10 Science Notes Chapter 5 (Chapter 5) – Examplar Problems (English) Book

Detailed Notes with MCQs of Chapter 5, 'Periodic Classification of Elements', from your NCERT Exemplar. This is a crucial chapter, not just for your board exams, but also forms the foundation for chemistry questions in various government exams. We'll cover the key concepts thoroughly.

Chapter 5: Periodic Classification of Elements - Detailed Notes

1. Need for Classification:

• Early chemists faced the challenge of studying a growing number of discovered elements individually.
• Classification was needed to organize elements based on their similarities and differences, making their study systematic and easier.

2. Early Attempts at Classification:

• (a) Dobereiner's Triads (Johann Dobereiner, 1817):

  • Concept: Arranged chemically similar elements in groups of three (triads). The atomic mass of the middle element was approximately the arithmetic mean (average) of the atomic masses of the other two elements.
  • Examples:
    • Lithium (Li, 6.9), Sodium (Na, 23.0), Potassium (K, 39.0). Average = (6.9 + 39.0) / 2 = 22.95 ≈ 23.0
    • Calcium (Ca, 40.1), Strontium (Sr, 87.6), Barium (Ba, 137.3). Average = (40.1 + 137.3) / 2 = 88.7 ≈ 87.6
  • Limitations: Dobereiner could identify only a few such triads from the elements known at that time. This classification was not applicable to all elements.

• (b) Newlands' Law of Octaves (John Newlands, 1866):

  • Concept: Arranged the known elements in order of increasing atomic masses. He observed that the properties of every eighth element were similar to those of the first element, like the notes in a musical octave (Sa, Re, Ga, Ma, Pa, Dha, Ni, Sa).
  • Example: Properties of Lithium (1st) were found similar to Sodium (8th); Beryllium (2nd) similar to Magnesium (9th), and so on.
  • Limitations:
    • Applicable only up to Calcium (Ca). Elements beyond Ca did not follow the law.
    • Assumed only 56 elements existed and no new elements would be discovered. The discovery of Noble gases later disrupted the pattern.
    • To fit elements into his table, Newlands placed two elements in the same slot (e.g., Co and Ni) and put some unlike elements under the same note (e.g., Co, Ni, Pd, Pt, Ir under the note 'Do' along with F, Cl, Br which have very different properties).

3. Mendeleev's Periodic Table (Dmitri Mendeleev, 1869):

• Basis: Increasing order of atomic mass and similarity in chemical properties. He focused on the compounds formed by elements with Oxygen (Oxides) and Hydrogen (Hydrides) as these are very reactive and formed compounds with most elements.
• Mendeleev's Periodic Law: "The properties of elements are the periodic function of their atomic masses."
• Structure: Contained vertical columns called 'Groups' and horizontal rows called 'Periods'.
• Achievements (Merits):
  • Systematic Study: Grouped elements with similar properties, making study organized.
  • Correction of Atomic Masses: Corrected the atomic masses of some elements like Beryllium (from 13.5 to 9), Indium, Gold based on their expected position and properties.
  • Prediction of New Elements: Left gaps for undiscovered elements and predicted their properties based on their position. He named them using Sanskrit numeral 'Eka' (one) as a prefix to the name of the preceding element in the same group.
    • Eka-boron (predicted) -> Scandium (discovered later)
    • Eka-aluminium (predicted) -> Gallium (discovered later)
    • Eka-silicon (predicted) -> Germanium (discovered later)
    • The properties of the discovered elements matched remarkably well with Mendeleev's predictions.
• Limitations (Demerits):
  • Position of Isotopes: Isotopes (atoms of the same element with same atomic number but different mass numbers) have different atomic masses but similar chemical properties. Mendeleev's table based on atomic mass couldn't assign them separate places or justify placing them together.
  • Anomalous Pairs (Wrong Order of Atomic Masses): In some cases, elements with higher atomic mass were placed before elements with lower atomic mass to maintain similarity in properties (e.g., Cobalt (Co, 58.9) placed before Nickel (Ni, 58.7); Tellurium (Te, 127.6) before Iodine (I, 126.9)).
  • Position of Hydrogen: Hydrogen resembles alkali metals (Group 1) in forming positive ions (H+) and halides, oxides, sulphides (e.g., HCl, H₂O, H₂S like NaCl, Na₂O, Na₂S). It also resembles halogens (Group 17) in existing as a diatomic molecule (H₂) and forming covalent compounds. Its position was not definitively fixed.

4. The Modern Periodic Table (Henry Moseley, 1913):

• Basis: Increasing order of atomic number (Z). Moseley showed through experiments that atomic number (number of protons in the nucleus) is a more fundamental property of an element than its atomic mass.
• Modern Periodic Law: "Properties of elements are a periodic function of their atomic number."
• Significance: Arranging elements by atomic number automatically resolved most of Mendeleev's anomalies:
  • Isotopes: All isotopes of an element have the same atomic number, so they occupy the same position.
  • Anomalous Pairs: The order Co (Z=27) before Ni (Z=28) and Te (Z=52) before I (Z=53) is correct according to atomic number.
• Structure: Consists of 18 vertical columns called Groups and 7 horizontal rows called Periods.

5. Position of Elements in the Modern Periodic Table:

• Periods:
  • There are 7 periods. The period number corresponds to the principal electron shell (valence shell) being filled.
  • Number of elements in a period: 1st (2), 2nd (8), 3rd (8), 4th (18), 5th (18), 6th (32), 7th (incomplete, currently 32). These numbers correspond to 2n², where n is the shell number (with modifications for higher periods due to energy level overlaps).
• Groups:
  • There are 18 groups. Elements in the same group have the same number of valence electrons and hence, similar chemical properties.
  • Group 1: Alkali Metals (Valence electrons = 1)
  • Group 2: Alkaline Earth Metals (Valence electrons = 2)
  • Groups 3-12: Transition Metals
  • Group 13: Boron Family (Valence electrons = 3)
  • Group 14: Carbon Family (Valence electrons = 4)
  • Group 15: Nitrogen Family (Pnictogens) (Valence electrons = 5)
  • Group 16: Oxygen Family (Chalcogens) (Valence electrons = 6)
  • Group 17: Halogens (Valence electrons = 7)
  • Group 18: Noble Gases (Inert Gases) (Valence electrons = 8, except He which has 2)
• Electronic Configuration and Position: The electronic configuration of an element determines its position. The number of valence electrons determines the group number (for main groups), and the principal quantum number (n) of the valence shell determines the period number.
  • Example: Sodium (Na), Z=11. Electronic Configuration: 2, 8, 1. Valence shell = 3rd shell (Period 3). Valence electrons = 1 (Group 1).

6. Trends in the Modern Periodic Table:

• (a) Valency:

  • Across a Period (Left to Right): For representative elements (Groups 1, 2, 13-18), valency with respect to hydrogen/electropositive elements first increases from 1 to 4, then decreases from 4 to 0. Valency with respect to oxygen increases from 1 to 7. Alternatively, it's often considered equal to the number of valence electrons (Groups 1, 2, 13, 14) or (8 - number of valence electrons) (Groups 15, 16, 17, 18).
  • Down a Group: Valency remains the same because the number of valence electrons is the same.

• (b) Atomic Size (Atomic Radius):

  • Across a Period (Left to Right): Atomic size decreases. Reason: The atomic number increases, meaning the positive charge on the nucleus (nuclear charge) increases. This increased nuclear charge pulls the electrons (in the same shell) closer to the nucleus, reducing the size.
  • Down a Group: Atomic size increases. Reason: A new electron shell is added at each step down the group. The increased number of shells increases the distance between the outermost electrons and the nucleus, despite the increase in nuclear charge. The inner shells also exert a shielding effect.

• (c) Metallic Character (Electropositivity): Tendency of an atom to lose electrons.

  • Across a Period (Left to Right): Metallic character decreases. Reason: Due to increasing nuclear charge and decreasing atomic size, the valence electrons are held more tightly, making it harder to lose them. Elements transition from metals (left) -> metalloids -> non-metals (right).
  • Down a Group: Metallic character increases. Reason: Atomic size increases, and the valence electrons are farther from the nucleus. The nuclear pull on these electrons weakens, making it easier to lose them.

• (d) Non-metallic Character (Electronegativity): Tendency of an atom to gain electrons.

  • Across a Period (Left to Right): Non-metallic character increases. Reason: Increasing nuclear charge and decreasing atomic size make it easier for the atom to attract incoming electrons.
  • Down a Group: Non-metallic character decreases. Reason: Increasing atomic size means the nucleus is less effective at attracting incoming electrons due to increased distance and shielding effect.

• (e) Nature of Oxides:

  • Across a Period (Left to Right): The nature of oxides changes from basic to amphoteric to acidic.
    • Metals form basic oxides (e.g., Na₂O, MgO).
    • Metalloids or some metals form amphoteric oxides (react with both acids and bases, e.g., Al₂O₃, ZnO).
    • Non-metals form acidic oxides (e.g., SO₂, Cl₂O₇).
  • Down a Group:
    • For metals, the basic character of oxides increases (e.g., BeO < MgO < CaO).
    • For non-metals, the acidic character of oxides generally decreases.

7. Metalloids:

• Elements that exhibit properties intermediate between those of metals and non-metals.
• Located along the zig-zag line separating metals and non-metals in the periodic table.
• Examples: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Polonium (Po).

Multiple Choice Questions (MCQs):

1. Which of the following correctly represents a Dobereiner's Triad?
   (a) Li, Be, B
   (b) H, Li, Na
   (c) Cl, Br, I
   (d) Fe, Co, Ni

2. According to Newlands' Law of Octaves, the properties of Sodium are similar to those of:
   (a) Beryllium
   (b) Lithium
   (c) Magnesium
   (d) Potassium

3. Mendeleev's Periodic Law states that the properties of elements are a periodic function of their:
   (a) Atomic Number
   (b) Number of Neutrons
   (c) Atomic Mass
   (d) Number of Valence Electrons

4. The element 'Eka-aluminium' predicted by Mendeleev was later discovered as:
   (a) Silicon
   (b) Germanium
   (c) Scandium
   (d) Gallium

5. An element has electronic configuration 2, 8, 5. In the Modern Periodic Table, it belongs to:
   (a) Period 3, Group 5
   (b) Period 5, Group 3
   (c) Period 3, Group 15
   (d) Period 5, Group 15

6. Which of the following elements has the largest atomic radius?
   (a) Na
   (b) Mg
   (c) K
   (d) Ca

7. As we move from left to right in Period 3 of the Modern Periodic Table, the metallic character of elements:
   (a) Increases
   (b) Decreases
   (c) Remains the same
   (d) First increases then decreases

8. Which of the following sets of elements is written in order of increasing non-metallic character?
   (a) F, Cl, Br, I
   (b) C, N, O, F
   (c) Na, Mg, Al, Si
   (d) Be, Mg, Ca, Sr

9. An element X forms an oxide X₂O₃ which is amphoteric in nature. Element X is likely to be:
   (a) Sodium (Na)
   (b) Aluminium (Al)
   (c) Sulphur (S)
   (d) Magnesium (Mg)

10. Which of the following was NOT a limitation of Mendeleev's Periodic Table?
    (a) Position of Isotopes
    (b) Prediction of new elements
    (c) Anomalous pair of Tellurium and Iodine
    (d) Position of Hydrogen

Answers to MCQs:

1. (c) Cl, Br, I
2. (b) Lithium
3. (c) Atomic Mass
4. (d) Gallium
5. (c) Period 3, Group 15 (Period = 3rd shell; Valence electrons = 5, so Group = 10+5 = 15)
6. (c) K (Atomic size increases down a group (K>Na, Ca>Mg) and decreases across a period (Na>Mg, K>Ca). Comparing K and Ca, K is larger. Comparing K and Na, K is larger. K is the largest among the options).
7. (b) Decreases
8. (b) C, N, O, F (Non-metallic character increases across a period). Option (a) shows decreasing non-metallic character down a group. Option (c) shows decreasing metallic character. Option (d) shows increasing metallic character down a group.
9. (b) Aluminium (Al) (Al₂O₃ is a common example of an amphoteric oxide).
10. (b) Prediction of new elements (This was a major achievement/merit, not a limitation).

Make sure you understand the reasons behind the trends and the limitations of the earlier models. This chapter requires understanding the link between electronic configuration, position in the table, and properties. Good luck with your preparation!