Class 10 Science Notes Chapter 5 (Periodic classification of elements) – Science Book

Okay, let's delve into the detailed notes for Chapter 5: Periodic Classification of Elements from the NCERT Class 10 Science textbook, focusing on aspects relevant for government exam preparation.

Chapter 5: Periodic Classification of Elements

1. Introduction: Why Classify Elements?

• Around 118 elements are known presently. Studying the properties of each element individually is extremely difficult.
• Classification involves grouping elements with similar properties together and separating elements with dissimilar properties.
• This systematic arrangement makes the study of elements easier, more organized, and helps in understanding the relationship between different elements and predicting the properties of new or undiscovered elements.

2. Early Attempts at Classification

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

  • Concept: He grouped elements with similar chemical properties into groups of three, called triads.
  • Law: When the three elements in a triad were arranged in order of increasing atomic mass, the atomic mass of the middle element was roughly the average of the atomic masses of the other two elements.
  • Example:
    • Triad 1: Lithium (Li, At. mass ≈ 7), Sodium (Na, At. mass ≈ 23), Potassium (K, At. mass ≈ 39). Average mass of Li and K = (7 + 39) / 2 = 23, which is approximately the atomic mass of Na.
    • Other examples: Calcium (Ca), Strontium (Sr), Barium (Ba); Chlorine (Cl), Bromine (Br), Iodine (I).
  • Limitation: Dobereiner could identify only a few such triads from the elements known at that time. This classification was not applicable to all known elements.

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

  • Concept: He arranged the then-known elements in order of increasing atomic masses.
  • Law: He found that every eighth element had properties similar to that of the first element, analogous to the octaves found in music (Sa, Re, Ga, Ma, Pa, Dha, Ni, Sa...).
  • Example: Starting with Lithium (Li), the eighth element is Sodium (Na). Both have similar properties. Similarly, Beryllium (Be) and Magnesium (Mg) resemble each other.
  • Limitations:
    • The law was applicable only up to Calcium (Ca). Elements beyond Calcium did not follow this pattern.
    • Newlands assumed that only 56 elements existed in nature and no more elements would be discovered.
    • To fit elements into his table, Newlands placed two elements in the same slot (e.g., Cobalt and Nickel) and also grouped some unlike elements together (e.g., Co and Ni placed with halogens like F, Cl, Br).
    • Iron (Fe), which resembles Cobalt and Nickel in properties, was placed far away from these elements.

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

• Basis: Mendeleev arranged elements based on:
  1. Increasing Atomic Mass
  2. Similarity in Chemical Properties: He primarily focused on the nature and formulae of the oxides and hydrides formed by the elements. (e.g., elements forming R₂O oxides were grouped together, elements forming RH hydrides were grouped together).
• 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.
  • There were 8 Groups (numbered I to VIII) and 6 Periods initially.
  • The formulae of oxides and hydrides were indicated at the top of each group.
• Achievements/Merits:
  1. Systematic Study: Provided a systematic way to study elements and their compounds.
  2. Prediction of New Elements: Mendeleev left gaps in his table for undiscovered elements and predicted their properties based on their position. He named them using Sanskrit numeral 'Eka' (one) prefix to the name of the preceding element in the same group.
     • Eka-boron → Scandium (Sc)
     • Eka-aluminium → Gallium (Ga)
     • Eka-silicon → Germanium (Ge)
       These elements were discovered later and their properties matched Mendeleev's predictions remarkably well.
  3. Correction of Doubtful Atomic Masses: He corrected the atomic masses of some elements like Beryllium (Be), Indium (In), and Gold (Au) based on their expected positions and properties. (e.g., Be's mass was corrected from 13.5 to 9).
  4. Accommodation of Noble Gases: When noble gases (like He, Ne, Ar) were discovered later, they could be placed in a new group (Group 0) without disturbing the existing order.
• Limitations/Demerits:
  1. Position of Isotopes: Isotopes are atoms of the same element with the same atomic number but different atomic masses (e.g., ³⁵Cl and ³⁷Cl). Since Mendeleev's table was based on increasing atomic mass, isotopes should have been given different positions, but they were not, as they have similar chemical properties. This violated his own law.
  2. Anomalous Pairs of Elements: In some cases, elements with higher atomic mass were placed before elements with lower atomic mass to maintain the similarity in properties.
     • Cobalt (Co, At. mass 58.9) was placed before Nickel (Ni, At. mass 58.7).
     • Tellurium (Te, At. mass 127.6) was placed before Iodine (I, At. mass 126.9).
  3. Position of Hydrogen: Hydrogen resembles alkali metals (Group I) as it forms positive ions (H⁺) and compounds like HCl, H₂O, H₂S. It also resembles halogens (Group VII) as it exists as a diatomic molecule (H₂) and forms covalent compounds and negative ions (H⁻, hydride ion). Mendeleev could not assign a fixed, correct position to hydrogen.

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

• Basis: Henry Moseley showed through experiments that Atomic Number (Z) is a more fundamental property of an element than its Atomic Mass. Atomic Number represents the number of protons in the nucleus of an atom.
• Modern Periodic Law: "Properties of elements are a periodic function of their atomic number."
• Structure:
  • Elements are arranged in order of increasing atomic number.
  • Consists of 18 vertical columns called Groups and 7 horizontal rows called Periods.
  • Groups (1 to 18):
    • Elements in the same group have the same number of valence electrons (electrons in the outermost shell).
    • Therefore, elements in the same group exhibit similar chemical properties.
    • Group 1: Alkali Metals
    • Group 2: Alkaline Earth Metals
    • Groups 3-12: Transition Metals
    • Group 17: Halogens
    • Group 18: Noble Gases (or Inert Gases)
  • Periods (1 to 7):
    • Elements in the same period have the same number of electron shells.
    • The period number corresponds to the principal energy level (shell number) being filled.
    • Period 1: 2 elements (filling K shell)
    • Period 2: 8 elements (filling L shell)
    • Period 3: 8 elements (filling M shell)
    • Period 4: 18 elements
    • Period 5: 18 elements
    • Period 6: 32 elements (includes Lanthanides)
    • Period 7: 32 elements (includes Actinides) - currently incomplete as per standard representation.
  • Lanthanides and Actinides: These are placed separately at the bottom of the table to avoid making the table too wide. Lanthanides belong to Period 6, and Actinides belong to Period 7.
• Position of an Element and Electronic Configuration:
  • The period number is equal to the number of shells occupied by electrons.
  • The group number is related to the number of valence electrons.
    • For Group 1 and 2 elements: Group Number = Number of valence electrons.
    • For Group 13-18 elements: Group Number = 10 + Number of valence electrons.
• Advantages over Mendeleev's Table:
  • Basis on Atomic Number: More fundamental, explains periodicity better.
  • Position of Isotopes: Resolved, as all isotopes of an element have the same atomic number, they occupy the same position.
  • Anomalous Pairs: Resolved, as the elements are arranged by increasing atomic number (e.g., Co (Z=27) comes before Ni (Z=28); Te (Z=52) comes before I (Z=53)).
  • Clear Correlation: Position of elements is clearly linked to their electronic configurations.
  • Position of Hydrogen: Still debated, but usually placed in Group 1 due to its electronic configuration (1s¹) and tendency to form H⁺. Its unique nature is acknowledged.

5. Trends in the Modern Periodic Table

Understanding trends is crucial for predicting properties.

• a) Valency:

  • Definition: Combining capacity of an element, usually equal to the number of valence electrons or (8 - number of valence electrons).
  • Across a Period (Left to Right): Valency first increases from 1 to 4 (w.r.t oxygen/hydrogen for representative elements), and then decreases from 4 to 0 (for Groups 1, 2, 13, 14, 15, 16, 17, 18 respectively: 1, 2, 3, 4, 3, 2, 1, 0).
  • Down a Group: Valency remains the same because the number of valence electrons is the same for elements in a group.

• b) Atomic Size (Atomic Radius):

  • Definition: Distance from the centre of the nucleus to the outermost electron shell.
  • Across a Period (Left to Right): Atomic size decreases.
    • Reason: Although electrons are added to the same shell, the nuclear charge (number of protons) increases. This increased positive charge pulls the electrons closer to the nucleus, reducing the size.
  • Down a Group: Atomic size increases.
    • Reason: A new electron shell is added as we move down each step in a group. The increased number of shells increases the distance between the nucleus and the outermost electrons, overriding the effect of increased nuclear charge.

• c) Metallic Character (Electropositivity):

  • Definition: Tendency of an atom to lose electrons and form positive ions (cations). Metals are electropositive.
  • Across a Period (Left to Right): Metallic character decreases.
    • Reason: Nuclear charge increases, making it harder for the atom to lose its valence electrons due to stronger attraction.
  • Down a Group: Metallic character increases.
    • Reason: Atomic size increases, the valence electrons are farther from the nucleus and experience less attraction. They can be lost more easily.
  • Location: Metals are located on the left side and centre of the periodic table.

• d) Non-metallic Character (Electronegativity - implied concept):

  • Definition: Tendency of an atom to gain electrons and form negative ions (anions). Non-metals are generally electronegative.
  • Across a Period (Left to Right): Non-metallic character increases.
    • Reason: Nuclear charge increases and atomic size decreases, making it easier for the nucleus to attract incoming electrons.
  • Down a Group: Non-metallic character decreases.
    • Reason: Atomic size increases, the nucleus is farther from the incoming electron, and the attraction force is weaker.
  • Location: Non-metals are located on the right side of the periodic table.

• e) Nature of Oxides:

  • Metallic Oxides: Generally basic in nature (e.g., Na₂O, MgO). They react with acids. Basic character increases down a group and decreases across a period.
  • Non-metallic Oxides: Generally acidic in nature (e.g., SO₂, CO₂, P₄O₁₀). They react with bases. Acidic character increases across a period and decreases down a group.
  • Amphoteric Oxides: Some oxides show both acidic and basic behaviour (e.g., Al₂O₃, ZnO). These are typically formed by elements near the metal/non-metal dividing line (metalloids).
  • Trend Across a Period: The nature of oxides changes from strongly basic → weakly basic → amphoteric → weakly acidic → strongly acidic.

• f) Metalloids (Semi-metals):

  • Elements that exhibit properties intermediate between 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).

6. Potential Exam Focus Areas:

• State Dobereiner's Law of Triads / Newlands' Law of Octaves / Mendeleev's Periodic Law / Modern Periodic Law.
• Limitations of Dobereiner's, Newlands', and Mendeleev's classifications.
• Achievements of Mendeleev's Periodic Table (especially prediction of elements).
• Basis of Mendeleev's vs. Modern Periodic Table.
• How problems of Mendeleev's table were resolved in the Modern Periodic Table.
• Relationship between electronic configuration and position (Group, Period) in the Modern Periodic Table.
• Explaining trends (Valency, Atomic Size, Metallic/Non-metallic character) across a period and down a group with reasons.
• Predicting properties (atomic size, metallic character, formula of oxide/chloride) of an element based on its position.
• Identifying elements based on given electronic configurations or positions.
• Understanding the nature of oxides across periods and down groups.
• Location of metals, non-metals, metalloids, alkali metals, halogens, noble gases.

This comprehensive overview covers the key concepts from Chapter 5, emphasizing the laws, comparisons, structure, and trends crucial for government exam preparation based on the NCERT Class 10 syllabus. Remember to correlate these concepts with electronic configurations for a deeper understanding.