Chemistry chapter no 7) Modern Periodic Table

 

🧪 Chapter 7: Modern Periodic Table - Super-Set Question Bank

This chapter focuses heavily on reasoning and trends. Examiners often ask "Why?" questions rather than complex calculations.

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I. Tricky MCQs (Concept Checkers)

These questions test if you truly understand the trends, not just memorized them.

1. The element with the highest first ionization enthalpy in the periodic table is:
   a) Fluorine
   b) Helium
   c) Cesium
   d) Chlorine
   Answer: (b) Helium
   Trick: Students often pick Fluorine because it's the most electronegative, but Helium has the stable duplet and is the smallest, making it the hardest to remove an electron from.


2. Identify the species having the largest radius:
   a) Na
   b) Na⁺
   c) Mg²⁺
   d) Al³⁺
   Answer: (a) Na
   Concept: Cations are smaller than their parent atoms. Among isoelectronic ions like Na⁺, Mg²⁺, and Al³⁺, the one with the highest positive charge pulls electrons closest, making it the smallest. Na (neutral) is the largest.


3. Which of the following has the most negative electron gain enthalpy?
   a) Fluorine
   b) Chlorine
   c) Oxygen
   d) Nitrogen
   Answer: (b) Chlorine
   Classic Trap: Fluorine is very small, so incoming electrons feel repulsion. Chlorine is larger and can accommodate the extra electron more easily, releasing more energy.


4. The correct order of metallic character is:
   a) B > Al > Mg > K
   b) K > Mg > Al > B
   c) Al > Mg > B > K
   d) Mg > Al > K > B
   Answer: (b) K > Mg > Al > B
   Logic: Metallic character increases down a group and decreases across a period. Potassium (K) is furthest down-left.


5. If the valence shell electronic configuration is ns²np⁵, the element belongs to:
   a) Alkali metals
   b) Halogens
   c) Noble gases
   d) Alkaline earth metals
   Answer: (b) Halogens

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II. "Must-Do" Definitions

Memorize these exact definitions as they appear frequently in 1-mark questions.

• Modern Periodic Law: "The physical and chemical properties of elements are a periodic function of their atomic numbers."
• Ionization Enthalpy: The energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
• Electron Gain Enthalpy: The enthalpy change that takes place when an electron is added to an isolated gaseous atom in its ground state.
• Electronegativity: The ability of a covalently bonded atom to attract the shared pair of electrons toward itself.
• Isoelectronic Species: Atoms and ions that contain the same number of electrons (e.g., O^2-, F^-, Na⁺, Mg²⁺).
• Screening (Shielding) Effect: The decrease in the force of attraction exerted by the nucleus on the valence electrons due to the presence of electrons in the inner shells.

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III. Conceptual & Reasoning Questions (The "Why" Questions)

This is where most marks are lost. Be prepared to write 2-3 sentences for each.

A. Trends & Anomalies

1. Why is the first ionization enthalpy of Boron (B) smaller than Beryllium (Be)?
   Reason: Be (1s² 2s²) has a stable, fully filled 2s orbital. B (1s² 2s² 2p¹) has a single electron in the 2p orbital. Removing the 2p electron is easier than breaking the stable 2s pair.
2. Why is the first ionization enthalpy of Oxygen (O) smaller than Nitrogen (N)?
   Reason: Nitrogen (2s² 2p³) has a stable half-filled p-orbital. Oxygen (2s² 2p⁴) has one paired electron in the p-orbital; losing it relieves repulsion and leads to a stable half-filled state.
3. Why do noble gases have positive electron gain enthalpy?
   Reason: They have fully filled stable electronic configurations (ns² np⁶). They have no tendency to accept an electron; energy must be supplied to force an electron into a higher energy shell.
4. Why is the electron gain enthalpy of Fluorine less negative than Chlorine?
   Reason: Fluorine is very small. The incoming electron faces strong repulsion from the existing electrons in the small 2p subshell. Chlorine is larger (3p), so the repulsion is less.
5. Explain the diagonal relationship with an example.
   Focus: Lithium and Magnesium (or Beryllium and Aluminium) show similar properties due to similar ionic sizes and charge/size ratios.

B. Structure & Location

6. Write the general outer electronic configuration of:
   • s-block: ns^1-2
   • p-block: ns² np^1-6
   • d-block: (n-1)d^1-10 ns^1-2
   • f-block: (n-2)f^1-14 (n-1)d^0-1 ns²
7. Locate the group and period of an element with Z=17.
   Method: Write configuration (1s² 2s² 2p⁶ 3s² 3p⁵). Principal quantum number n=3 (Period 3). Valence electrons 2+5=7. Since it's p-block, Group = 10 + 7 = 17.

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IV. Short Notes & Distinguish

• Distinguish between Metals and Non-metals based on ionization enthalpy and electronegativity.
• Write a short note on the characteristics of Transition Elements (d-block). (Focus on variable valency, colored compounds, and catalytic properties).
• Write a short note on Lanthanides and Actinides. (Mention they are f-block, inner transition elements).

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🎓 Exam Strategy for This Chapter

• Don't just memorize trends (Left-to-Right, Top-to-Bottom). Memorize the exceptions (Be vs B, N vs O, F vs Cl). The exam papers usually ask about the exceptions.
• Know the blocks: If they give you an atomic number (e.g., Z=24), you must be able to instantly say "Chromium, Period 4, Group 6, d-block".

🧪 CHAPTER 7 — MODERN PERIODIC TABLE

🔹 Section A — MCQ (1 mark each) — (Each question: 2 points)

1. 1. If the valence shell electronic configuration is ns² np⁵, the element belongs to
   a) Alkali metals b) Halogens c) Alkaline earth metals d) Actinides
   • 1.The configuration ns² np⁵ corresponds to seven valence electrons (group 17).
   • 2.Elements with ns² np⁵ are halogens (highly reactive non-metals).
2. 2. According to modern periodic law, the properties of elements are the periodic function of
   a) Atomic mass b) Atomic number c) Neutron number d) Mass number
   • 1.Modern periodic law: properties vary periodically with atomic number (Z).
   • 2.Atomic number equals the number of protons and determines electronic structure and properties.
3. 3. In modern periodic table, period number indicates
   a) No. of valence electrons b) Principal quantum number c) Valence shell configuration d) Azimuthal quantum number
   • 1.The period number equals the principal quantum number (n) of the outermost shell.
   • 2.It indicates the highest energy level occupied by electrons in the ground state.
4. 4. General outer electronic configuration of d-block elements is
   a) (n – 1)d¹⁻¹⁰ ns⁰⁻² b) ns² np⁶ c) ns² d) ns² np³
   • 1.D-block elements have (n−1)d electrons and up to two ns electrons: (n−1)d¹–¹⁰ ns⁰–².
   • 2.They are called transition elements because d-orbitals are being filled.
5. 5. The number of groups and periods in the modern periodic table are
   a) 7 groups & 18 periods b) 8 groups & 7 periods c) 18 groups & 7 periods d) 18 groups & 8 periods
   • 1.Modern (long-form) periodic table has 18 vertical groups and 7 horizontal periods (for known ground-state elements).
   • 2.Groups show similar valence electron arrangements; periods correspond to principal quantum numbers.
6. 6. Which of the following is not a representative element?
   a) s-block b) p-block c) d-block d) Noble gas
   • 1.Representative (or main-group) elements are s- and p-block elements (Groups 1,2 and 13–18).
   • 2.d-block elements (transition metals) are not representative elements.

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🔹 Section B — Short Answer Questions (1–2 Marks) — (Each question: 4 points)

1. 1. State the modern periodic law.
   • 1.Properties of elements are a periodic function of their atomic numbers (Z).
   • 2.Atomic number determines electronic configuration, which explains periodicity of chemical and physical properties.
   • 3.This law corrects anomalies in earlier mass-based arrangements.
   • 4.It forms the basis of the modern long-form periodic table arranged by Z.
2. 2. Define group and period.
   • 1.Group: a vertical column in the periodic table; elements in a group have similar valence electron configurations and chemical properties.
   • 2.Period: a horizontal row in the periodic table; elements in a period have the same highest principal quantum number (n) but different properties.
   • 3.Group number (for main-group) often indicates number of valence electrons.
   • 4.Period number equals the principal quantum number of the outermost shell.
3. 3. Mention the total number of groups and periods in the modern periodic table.
   • 1.The long-form periodic table contains 18 groups (vertical columns).
   • 2.It has 7 periods (horizontal rows) for the known ground-state elements.
   • 3.Groups are numbered 1–18 in the IUPAC system.
   • 4.Periods correspond to principal quantum numbers n = 1 to 7.
4. 4. Which elements are called d-block elements? Why are they called so?
   • 1.d-block elements are those in which the (n−1)d orbitals are being filled (transition elements).
   • 2.Their general outer configuration is (n−1)d¹–¹⁰ ns⁰–².
   • 3.They are placed between s- and p-block in the long-form table (groups 3–12).
   • 4.They are called d-block because their characteristic electrons occupy d-orbitals which determine many of their properties (variable oxidation states, colored ions, magnetism).
5. 5. What are f-block elements? Where are they placed in the periodic table?
   • 1.f-block elements are those where 4f or 5f orbitals are being filled (lanthanoids and actinoids).
   • 2.Their general configuration involves (n−2)f electrons along with (n−1)d and ns electrons.
   • 3.They are placed separately as two rows (inner transition series) below the main table for compactness.
   • 4.Lanthanoids follow La (Ce–Lu) and actinoids follow Ac (Th–Lr) in the inner transition series.
6. 6. Why are noble gases placed in Group 18?
   • 1.Noble gases have complete valence shells (ns² np⁶, except He which is 1s²), giving them very stable electronic configurations.
   • 2.Complete shells make them chemically inert with very low reactivity.
   • 3.Group 18 groups elements with similar full-valence configurations and similar low chemical reactivity.
   • 4.Their physical and chemical properties (low boiling points, monatomic gases) are consistent across the group.
7. 7. What do the period number and group number of an element indicate?
   • 1.Period number indicates the principal quantum number (n) of the outermost shell (the highest energy level occupied).
   • 2.Group number (especially for main-group elements) indicates the number of valence electrons and thus typical valency.
   • 3.Group position determines similarity of chemical properties among elements in the same column.
   • 4.Period position reflects gradual changes in properties across a row due to increasing nuclear charge.
8. 8. State two features of the long form of periodic table.
   • 1.It arranges elements in 18 groups and 7 periods according to increasing atomic number, showing clear periodicity of properties.
   • 2.It separates elements into s-, p-, d- and f-blocks, making electronic structure and trends (like ionization energy, atomic radius) easy to visualize.
   • 3.It places transition elements (d-block) and inner transition elements (f-block) appropriately.
   • 4.It allows prediction of properties and placement of undiscovered elements (historically useful).
9. 9. Define core electrons and valence electrons.
   • 1.Valence electrons are the electrons in the outermost shell of an atom that participate in bonding (determine chemical properties).
   • 2.Core electrons are all the inner-shell electrons (not in the outermost shell) that do not usually participate in bonding.
   • 3.Valence electrons determine group behaviour; core electrons shield valence electrons from the nucleus.
   • 4.Core electrons remain essentially unchanged during most chemical reactions; valence electrons are involved in bond formation.
10. 10. Write the general electronic configuration of s, p, d, and f-block elements.
    • 1.s-block: outer configuration ns¹–² (Groups 1–2).
    • 2.p-block: outer configuration ns² np¹–⁶ (Groups 13–18).
    • 3.d-block: outer configuration (n−1)d¹–¹⁰ ns⁰–² (Groups 3–12).
    • 4.f-block: outer/inner configuration (n−2)f¹–¹⁴ (inner transition series, lanthanoids/actinoids).

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🔹 Section C — 3 Mark Questions — (Each question: 6 points)

1. 1. Explain the classification of elements into s, p, d, and f blocks with suitable examples.
   • 1.s-block: Elements with outer configuration ns¹–² (Groups 1–2). Example: Na (3s¹), Ca (4s²).
   • 2.p-block: Elements with outer configuration ns² np¹–⁶ (Groups 13–18). Example: C (2s²2p²), Cl (3s²3p⁵).
   • 3.d-block: Transition elements with (n−1)d filling and ns electrons. Example: Fe ([Ar]4s²3d⁶), Cu ([Ar]4s¹3d¹⁰).
   • 4.f-block: Inner transition elements filling 4f or 5f orbitals (lanthanoids and actinoids). Example: Ce (4f¹5d¹6s²), U (5f³6d¹7s²).
   • 5.This block classification corresponds to the orbital type that receives the last electron and explains common chemical behavior in each block.
   • 6.Blocks are shown in the long-form periodic table: s on left, p on right, d in centre, f as separate rows below.
2. 2. Explain the significance of group and period number with respect to position of an element.
   • 1.Group number reflects the number of valence electrons (for main-group elements) and therefore predicts typical valency and chemical reactivity.
   • 2.Elements in the same group show similar chemical properties due to similar valence configurations (e.g., alkali metals behave similarly).
   • 3.Period number gives the principal quantum number (n) of the outermost shell, indicating the size and energy level of valence shell.
   • 4.Elements across a period show gradual change in properties (atomic radius, ionization energy) due to increasing nuclear charge.
   • 5.Position (group & period) allows prediction of oxidation states, types of compounds formed, and periodic trends.
   • 6.Group & period together place an element precisely in the table, determining block (s/p/d/f) and hence many physical/chemical characteristics.
3. 3. Write electronic configurations and identify the block, group, and period: a) 1s² 2s² 2p⁵ b) [Ne] 3s² 3p¹ c) [Ar] 4s² 3d⁶
   • a)1s²2s²2p⁵ → Element: Fluorine (F); Block: p-block; Group: 17; Period: 2.
   • b)[Ne]3s²3p¹ → Element: Aluminium (Al); Block: p-block; Group: 13; Period: 3.
   • c)[Ar]4s²3d⁶ → Element: Iron (Fe); Block: d-block; Group: 8 (transition metal); Period: 4.
   • (Each identification follows from highest occupied orbitals and electron count.)
   • Block determination: look at orbital receiving last electron (s/p/d/f).
   • Group/period: group from valence electrons (main-group) or d-block numbering; period from highest n.
4. 4. Explain the advantages of long form periodic table (any six).
   • 1.Clear arrangement by increasing atomic number (Z) removes anomalies of mass-based tables.
   • 2.Separation into s, p, d, f blocks shows orbital filling and explains chemical behavior and trends.
   • 3.Helps predict properties and valencies of elements, including transition and inner-transition series.
   • 4.Provides consistent group numbering (1–18) for international use and comparison.
   • 5.Allows placement of lanthanoids and actinoids (f-block) separately for compactness while preserving periodicity.
   • 6.Helps in understanding periodic trends (atomic radius, ionization energy, electronegativity) systematically across periods and groups.
5. 5. State the limitations of Mendeleev periodic table and explain how modern periodic table overcomes them.
   • 1.Limitation: Mendeleev arranged elements by atomic mass which caused anomalies in order (some elements didn't fit property-wise by mass).
   • 2.Overcome: Modern table uses atomic number (Z) which correlates directly with electronic structure and resolves such anomalies.
   • 3.Limitation: Mendeleev left gaps for undiscovered elements but some placements were ambiguous for isotopes and hydrogen.
   • 4.Overcome: Modern periodic law explains placement by proton number and electronic configuration; isotopes have same Z so do not disrupt order.
   • 5.Limitation: No clear position for hydrogen and placement of some transition elements was problematic in mass-based table.
   • 6.Overcome: Using atomic number and blocks (s,p,d,f) gives clear positions for hydrogen, transition and inner-transition elements and explains their properties.

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🔹 Section D — Long Answer Questions (4 Marks) — (Each question: 8 points)

1. 1. Discuss periodic trends across a period and down a group for the following properties: a) Atomic radius b) Metallic / non-metallic character c) Valency d) Ionization enthalpy
   • a1)Atomic radius across a period: decreases left to right due to increasing nuclear charge pulling electrons closer.
   • a2)Atomic radius down a group: increases due to addition of new electron shells (higher n), despite increase in nuclear charge.
   • b1)Metallic character across a period: decreases (elements become less metallic and more non-metallic) because ionization energy increases.
   • b2)Metallic character down a group: increases as atoms become larger and lose electrons more easily (lower IE), so metals become more metallic.
   • c1)Valency across a period (for representative elements): generally increases from 1 to 4 (left to right) then decreases from 4 to 0 (toward noble gases) — reflects filling of valence electrons.
   • c2)Valency down a group: usually remains similar for elements of the same group (same number of valence electrons), though higher elements may show variable oxidation states.
   • d1)Ionization enthalpy across a period: increases left to right due to stronger nuclear attraction and smaller atomic radius.
   • d2)Ionization enthalpy down a group: decreases because valence electrons are farther from nucleus and better shielded, so less energy is needed to remove them.
2. 2. Explain Dobereiner triads, Newland’s law of octaves, and Mendeleev’s periodic law with two examples each.
   • 1.Dobereiner triads: Early classification where elements in groups of three (triads) showed middle element having average properties (atomic mass) of extremes. Example triads: Ca–Sr–Ba (atomic mass of Sr ≈ average of Ca and Ba); Cl–Br–I (properties follow pattern).
   • 2.Dobereiner significance: showed periodicity but applied only to few element groups.
   • 3.Newland’s law of octaves: Arranged elements by increasing atomic mass; every 8th element had similar properties (like musical octaves). Example: Li, Na, K show similarity; Be and Mg as another pair in the pattern.
   • 4.Newland’s limitation: worked well for lighter elements but failed for heavier ones and did not allow gaps for undiscovered elements.
   • 5.Mendeleev’s periodic law: Properties of elements are periodic functions of their atomic masses (as Mendeleev phrased it) and he arranged elements in table leaving gaps to predict undiscovered elements. Examples: Predicted eka-aluminium → gallium (properties matched prediction); predicted eka-silicon → germanium (properties as predicted).
   • 6.Mendeleev’s approach demonstrated predictive power and corrected ordering by grouping similar properties, paving way for modern table (later improved by atomic number ordering).
3. 3. Given the electronic configurations, identify i) Block ii) Group iii) Period iv) Category of element (metal / non-metal / metalloid) a) 1s² b) [Ne] 3s² 3p⁴ c) [Ar] 4s² 3d¹⁰ 4p³ d) [Xe] 4f¹⁴ 5d⁶ 6s²
   • a)1s² → Element: Helium (He). Block: s-block. Group: 18 (commonly placed in Group 18 due to noble-gas behaviour). Period: 1. Category: Noble gas (non-metal).
   • b)[Ne]3s²3p⁴ → Element: Sulfur (S). Block: p-block. Group: 16. Period: 3. Category: Non-metal.
   • c)[Ar]4s²3d¹⁰4p³ → Element: Arsenic (As). Block: p-block. Group: 15. Period: 4. Category: Metalloid.
   • d)[Xe]4f¹⁴5d⁶6s² → Element with Z = 76 (Osmium, Os). Block: d-block. Group: 8. Period: 6. Category: Metal (transition metal).
   • (Identification uses highest occupied orbital for block, electron count for group, highest n for period.)
   • (Category determined from position and typical chemical behavior — metals on left/centre, non-metals on right, metalloids in between.)