Samacheer Kalvi 9th Science Periodic Classification of Elements Exercise Solutions

The modern periodic law states that "the properties of elements are a periodic function of their atomic number." It was proposed by Henry Moseley in 1913.

A period is a horizontal row in the periodic table. There are 7 periods. The period number indicates the number of electron shells in the atoms of the elements in that period.

A group is a vertical column in the periodic table. There are 18 groups. Elements in the same group have the same number of valence electrons and similar chemical properties.

The first three alkali metals (Group 1) are: Lithium (Li), Sodium (Na), Potassium (K).

The halogens (Group 17) are: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At).

Noble gases have a stable electronic configuration — a complete octet (except He which has a duplet of 2 electrons). They have no tendency to lose, gain, or share electrons, making them chemically inert.

Metalloids are elements that exhibit properties of both metals and non-metals. They lie along the zigzag line in the periodic table. Examples: Silicon (Si) and Germanium (Ge). Others include Boron (B), Arsenic (As), Antimony (Sb), and Tellurium (Te).

Dobereiner observed that elements with similar properties could be grouped in sets of three called triads. In each triad, the atomic mass of the middle element is approximately equal to the average of the atomic masses of the other two elements.

Example: The triad of lithium (Li = 7), sodium (Na = 23), and potassium (K = 39). Average of Li and K = (7 + 39) / 2 = 23, which equals the atomic mass of sodium.

Limitation: Only a limited number of triads could be identified; most elements did not fit this pattern.

Newlands' Law of Octaves: When elements are arranged in order of increasing atomic mass, the properties of every 8th element are similar to those of the 1st element, just as the 8th note in music resembles the 1st.

Limitations: 1. The law was applicable only up to calcium (Z = 20). Heavier elements did not follow this pattern. 2. It assumed that only 56 elements existed; newly discovered elements had no place. 3. It did not accommodate transition elements. 4. Some dissimilar elements were placed together (e.g., Co and Ni with halogens).

Merits: 1. Systematic classification: First successful attempt to classify all known elements systematically. 2. Prediction of new elements: Left gaps for undiscovered elements and predicted their properties accurately (eka-boron → Scandium, eka-aluminium → Gallium, eka-silicon → Germanium). 3. Correction of atomic masses: Corrected atomic masses of some elements (e.g., beryllium, indium) based on their position. 4. Grouping of similar elements: Elements with similar properties were placed in the same group.

Limitations: 1. Position of hydrogen: Resembles both alkali metals (Group 1) and halogens (Group 17); position was ambiguous. 2. Isotopes: Isotopes have different atomic masses but similar properties; could not be placed together. 3. Anomalous pairs: Co (58.9) placed before Ni (58.7), and Te (127.6) before I (126.9) — violating the atomic mass order. 4. No separate group for noble gases: Noble gases were discovered later and had no place initially. 5. Dissimilar elements grouped together: Alkali metals (Na, K) placed with coinage metals (Cu, Ag, Au).

1. Based on atomic number (not mass): Eliminates anomalies like Co/Ni and Te/I since arrangement is by atomic number (Z), not atomic mass. 2. Position of hydrogen: Hydrogen is placed in Group 1 as it has 1 valence electron, though it also shares some properties with halogens. 3. Isotopes resolved: Since isotopes have the same atomic number, they occupy the same position in the modern table. 4. Noble gases accommodated: Placed in Group 18 as a separate group. 5. Clear separation: Metals, non-metals, and metalloids are clearly separated; metalloids lie along the zigzag line. 6. Systematic arrangement: 7 periods and 18 groups accommodate all 118 known elements.

Across a period (left to right): Atomic size decreases.

• Reason: The nuclear charge increases (more protons) while the number of electron shells remains the same. The greater nuclear attraction pulls electrons closer to the nucleus.
• Example: In Period 2 — Li (152 pm) > Be > B > C > N > O > F > Ne (71 pm).

Down a group: Atomic size increases.

• Reason: A new electron shell is added in each successive period. The valence electrons are farther from the nucleus, and the shielding effect of inner electrons reduces the effective nuclear pull.
• Example: In Group 1 — Li (2 shells) < Na (3 shells) < K (4 shells) < Rb (5 shells).

Metallic character across a period (L → R): Decreases

• Reason: Increasing nuclear charge makes it harder to lose electrons. Elements become less metallic and more non-metallic.
• Example: Na (strong metal) → Mg → Al → Si (metalloid) → P → S → Cl (non-metal).

Metallic character down a group: Increases

• Reason: Atomic size increases and the outermost electrons are farther from the nucleus. The nuclear attraction on valence electrons decreases, making it easier to lose them (metallic property).
• Example: In Group 14 — C (non-metal) < Si < Ge < Sn (metal) < Pb (metal).

Non-metallic character across a period (L → R): Increases

• Reason: It becomes easier to gain electrons as nuclear charge increases and atomic size decreases.

Non-metallic character down a group: Decreases

• Reason: It becomes harder to attract electrons when the valence shell is farther from the nucleus.
• Example: F (most reactive non-metal) > Cl > Br > I.

Mendeleev's Periodic Table (1869): Mendeleev arranged the 63 known elements in order of increasing atomic mass and observed that properties of elements repeated at regular intervals. He placed elements with similar properties in the same vertical column (group). The horizontal rows were called periods.

Merits: 1. Systematic classification: First successful attempt to classify all known elements in a systematic manner, making study easier. 2. Prediction of undiscovered elements: Mendeleev left gaps for elements not yet discovered and predicted their properties with remarkable accuracy. His predictions for eka-boron (scandium, discovered 1879), eka-aluminium (gallium, discovered 1875), and eka-silicon (germanium, discovered 1886) proved correct. 3. Correction of atomic masses: He corrected the atomic masses of some elements (e.g., beryllium from 13.5 to 9, indium from 75 to 113) based on their expected position in the table. 4. Grouping of similar elements: Elements with similar chemical properties were placed in the same group, enabling better understanding of chemical behaviour.

Limitations: 1. Anomalous position of cobalt and nickel: Co (atomic mass 58.9) was placed before Ni (58.7), but Co has a higher atomic mass, violating the arrangement by increasing atomic mass. 2. Position of hydrogen: Hydrogen resembles both alkali metals (Group 1 — forms H⁺ ion, has 1 valence electron) and halogens (Group 17 — exists as H₂, gains 1 electron to form H⁻). Its position was ambiguous. 3. Position of isotopes: Isotopes of the same element have different atomic masses but similar properties. Mendeleev's table (based on atomic mass) could not accommodate them in the same position. 4. No place for noble gases: Noble gases were discovered later (by Ramsay) and had no place in the original table. They were later placed in a separate group. 5. Dissimilar elements in same group: Some elements with very different properties were placed in the same group (e.g., alkali metals Na, K with coinage metals Cu, Ag, Au). 6. Position of rare earth elements: Lanthanides and actinides were not properly accommodated.

The Modern Periodic Table (Long Form):

• Based on the modern periodic law: "Properties of elements are a periodic function of their atomic number."
• Proposed by Henry Moseley (1913).
• Elements arranged in order of increasing atomic number.

Periods (Horizontal Rows) — 7 in number:

The period number indicates the number of electron shells (or energy levels) in the atoms of the elements in that period.

Groups (Vertical Columns) — 18 in number:

Key groups to remember:

• Group 1: Alkali metals (Li, Na, K, Rb, Cs, Fr) — 1 valence electron, highly reactive.
• Group 2: Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) — 2 valence electrons.
• Groups 3–12: Transition metals — variable valency, form coloured compounds.
• Group 13: Boron family — 3 valence electrons.
• Group 14: Carbon family — 4 valence electrons.
• Group 15: Nitrogen family — 5 valence electrons.
• Group 16: Oxygen family (chalcogens) — 6 valence electrons.
• Group 17: Halogens (F, Cl, Br, I, At) — 7 valence electrons, highly reactive non-metals.
• Group 18: Noble gases (He, Ne, Ar, Kr, Xe, Rn) — 8 valence electrons (2 for He), chemically inert.

Elements in the same group have the same number of valence electrons and hence similar chemical properties.

Position of Elements:

• Metals: Occupy the left and central portions of the periodic table. Approximately 78% of all elements are metals. They are electropositive — lose electrons to form positive ions.
• Non-metals: Occupy the right side of the periodic table (Groups 14–18, upper portions). They are electronegative — gain or share electrons.
• Metalloids: Lie along the zigzag (staircase) line separating metals from non-metals: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), and tellurium (Te). They have properties intermediate between metals and non-metals and are useful as semiconductors.
• Noble gases (Group 18): Placed in the extreme right column. They have completely filled valence shells (ns²np⁶, except He which has 1s²). This makes them chemically inert under normal conditions. They are monoatomic, colourless, and odourless gases at room temperature.

• Across a period: Valency first increases from 1 to 4, then decreases to 0.
• Example: Na(1), Mg(2), Al(3), Si(4), P(3), S(2), Cl(1), Ar(0).
• Down a group: Valency remains the same (same number of valence electrons).
• Example: All Group 1 elements have valency 1.

• Across a period: Decreases (left to right).
• Reason: Nuclear charge increases (more protons), pulling electrons closer. Number of shells stays the same, so the effective nuclear pull is stronger.
• Example: Li (152 pm) > Be (112 pm) > B (85 pm) > ... > Ne (71 pm).
• Down a group: Increases.
• Reason: A new shell of electrons is added in each successive period, increasing the distance between the outermost electrons and the nucleus. The shielding effect of inner electrons also reduces the effective nuclear pull.
• Example: Li (2 shells) < Na (3 shells) < K (4 shells) < Rb (5 shells).

• Across a period: Decreases (left to right).
• Reason: Increasing nuclear charge makes it harder to lose electrons. Elements become less metallic and more non-metallic.
• Example: Na (strong metal) → Mg → Al → Si (metalloid) → P → S → Cl (non-metal).
• Down a group: Increases.
• Reason: Atomic size increases, the hold of the nucleus on valence electrons weakens, making it easier to lose electrons (metallic property).
• Example: C (non-metal) < Si < Ge < Sn (metal) < Pb (metal).

• Across a period: Increases (left to right).
• Reason: It becomes easier to gain electrons as nuclear charge increases and atomic size decreases.
• Down a group: Decreases.
• Reason: It becomes harder to attract electrons when the valence shell is farther from the nucleus.
• Example: F (most reactive non-metal) > Cl > Br > I.

• Across a period: Decreases.
• Down a group: Increases.
• (Same trend as metallic character)

Summary Table: