Periodic Classification of Elements – MCQs with Answers and Explanations
Class 10
Chemistry — Chapter 14
Board: CBSE
Format: Topic-wise MCQs with concept-clearing explanations
Quick index
Topics & MCQ ranges:
- History & Early Attempts — Q1–Q8
- Mendeleev & Modern Periodic Law — Q9–Q18
- Electronic Configuration & Valency — Q19–Q30
- Periodic Trends (size, IE, electronegativity) — Q31–Q42
- Groups & Group Properties — Q43–Q52
- Metals, Non-metals & Transition Elements — Q53–Q60
History & Early Attempts (Q1–Q8)
Q1. Döbereiner’s triads were an early attempt to classify elements. In a triad, the atomic mass of the middle element is approximately:
Answer: B. Döbereiner observed that in a triad the atomic mass of the middle element is roughly the average of the other two, e.g., Li, Na, K. This pointed to patterns in properties and masses among elements.
Q2. Newlands’ law of octaves compared chemical periodicity to musical octaves. It failed mainly because:
Answer: C. Newlands’ scheme left no gaps; it grouped dissimilar elements and failed beyond calcium. The inability to accommodate undiscovered elements and heavier elements made it inadequate.
Q3. Which of the following was a major success of Mendeleev’s periodic table?
Answer: B. Mendeleev predicted properties of undiscovered elements (e.g., eka-aluminium → gallium). His predictions were experimentally confirmed later, validating his table.
Q4. A limitation of Mendeleev’s table was that it:
Answer: B. Mendeleev’s table was based on atomic masses and could not accommodate isotopes properly; isotopes have same chemical properties but different masses. This was resolved by arranging elements by atomic number.
Q5. Newlands noticed repetition of properties every eighth element when arranged by increasing atomic mass. This pattern is known as:
Answer: C. Newlands' observation is called the law of octaves. It compared periodicity to musical octaves but failed for heavier elements.
Q6. Mendeleev left gaps in his periodic table because:
Answer: C. Mendeleev left gaps for undiscovered elements and predicted their properties. When such elements were found (gallium, germanium), it supported his table.
Q7. Which scientist’s work established atomic number as the basis for modern periodic law?
Answer: C. Moseley’s X-ray studies established that atomic number determines element properties, leading to the modern periodic law. He showed atomic number (protons) is more fundamental than atomic mass.
Q8. The discovery of noble gases led to which change in periodic classification?
Answer: A. Noble gases formed a new Group (Group 18) because of their inertness and complete valence shells. Their properties highlighted the role of electronic configuration in chemical behaviour.
Mendeleev & Modern Periodic Law (Q9–Q18)
Q9. The modern periodic law states that properties of elements are periodic functions of their:
Answer: B. Modern periodic law uses atomic number (number of protons) as the organizing principle. This arrangement resolves anomalies present in mass-based ordering.
Q10. How many groups and periods are there in the modern periodic table?
Answer: B. The modern table has 18 groups and 7 periods in the long form. Lanthanoids and actinoids are placed separately as f-block rows.
Q11. Which block contains the transition elements?
Answer: C. Transition elements occupy the d-block (groups 3–12). They characteristically have partially filled d-orbitals.
Q12. Elements with similar chemical properties occur at regular intervals when arranged by atomic number. This phenomenon is called:
Answer: B. Periodicity describes the recurring trends in properties across periods and groups. Examples include atomic radius, ionization energy trends.
Q13. Why are lanthanoids and actinoids placed separately at the bottom of the periodic table?
Answer: B. f-orbitals are filled in lanthanoids and actinoids; placing them separately keeps the table compact. They also exhibit similar chemical properties within each series.
Q14. Atomic number of an element is equal to the number of:
Answer: C. Atomic number = number of protons; in neutral atoms this equals number of electrons. Atomic number uniquely identifies an element.
Q15. Noble gases are inert because they have:
Answer: C. Noble gases have full valence shells (octet or duet) making them chemically stable and largely inert. This explains their low reactivity.
Q16. Which of the following elements is a p-block element?
Answer: C. Chlorine is a p-block element (group 17). p-block contains non-metals, metalloids and some metals in groups 13–18.
Q17. Which of these statements about the modern periodic table is false?
Answer: C. The modern periodic table has 7 periods, not 18. It has 18 groups (vertical columns).
Q18. Elements in a vertical column of the periodic table are called:
Answer: B. Vertical columns are groups; horizontal rows are periods. Groups share similar chemical properties.
Electronic Configuration & Valency (Q19–Q30)
Q19. The electronic configuration of sodium (Z=11) is:
Answer: A. Sodium: 1s²2s²2p⁶3s¹. Highest occupied shell n=3 → period 3 and s-block (Group 1).
Q20. Which of the following elements has valency 2?
Answer: B. Oxygen (group 16) commonly has valency 2 (gains two electrons to form O²⁻). Valency relates to electrons gained/lost to attain noble gas configuration.
Q21. An element has configuration [Ne]3s2 3p3. Its group and period are:
Answer: A. [Ne]3s²3p³ corresponds to phosphorus (group 15) in period 3. 3p³ → 5 valence electrons → group 15.
Q22. Which of the following elements will have the same valency as sulphur (S)?
Answer: A. Oxygen and sulphur are in the same group (16) and commonly show valency 2. Group members have similar valencies due to similar valence electron counts.
Q23. The valency of carbon is 4 because it has:
Answer: B. Carbon has 4 valence electrons (2s²2p²) and needs four more to complete octet, hence valency 4. This allows formation of four covalent bonds.
Q24. Magnesium (Z=12) readily forms Mg²⁺ because:
Answer: A. Mg: 1s²2s²2p⁶3s² — two valence electrons in 3s, which it loses to become Mg²⁺. Losing two electrons gives a noble gas configuration.
Q25. Which electronic configuration corresponds to a noble gas?
Answer: B. 1s²2s²2p⁶ is neon (complete n=2 shell). Complete valence shell = noble gas configuration.
Q26. Which of the following has electronic configuration [Ar]4s2 3d10?
Answer: A. Zinc has [Ar]4s²3d¹⁰. Filled d-orbitals result in +2 oxidation state commonly.
Q27. The period number of an element equals:
Answer: B. Period number = highest occupied principal quantum number (n). E.g., elements with outer electrons in n=3 belong to period 3.
Q28. Which species has the electronic configuration 1s2 2s2 2p6?
Answer: D. Na⁺, Ne and F⁻ all have 1s²2s²2p⁶ (neon configuration). Isoelectronic species share same electron configuration.
Q29. Which of these elements will most likely form a +1 ion?
Answer: B. Sodium (Group 1) commonly forms Na⁺ by losing one electron. Group 1 elements have valency 1.
Q30. Which of these elements is likely to form covalent bonds rather than ionic?
Answer: B. Chlorine forms covalent bonds with non-metals (and ionic with metals); as a high-electronegativity non-metal, it shares electrons in many compounds. Non-metals tend to form covalent bonds.
Periodic Trends (Atomic Size, Ionisation Energy, Electronegativity) Q31–Q42
Q31. Atomic radius generally ______ across a period from left to right and ______ down a group.
Answer: B. Atomic radius decreases across a period (stronger nuclear charge) and increases down a group (additional shells). Effective nuclear charge and shielding explain this.
Q32. Ionisation energy is the energy required to:
Answer: B. Ionisation energy = energy to remove an electron from a gaseous atom/ion. First ionisation energy removes the first electron.
Q33. Which element would have the highest first ionisation energy?
Answer: C. Fluorine (a small, highly electronegative element) has very high first ionization energy among these. Small size and high effective nuclear charge make electron removal difficult.
Q34. Electronegativity generally increases in which direction across the periodic table?
Answer: A. Electronegativity increases left to right (towards fluorine). Atoms more strongly attract shared electrons when nuclear charge is higher.
Q35. Which of the following explains why reactivity of alkali metals increases down the group?
Answer: C. Down the group valence electron is farther and more shielded, so it is easier to lose, increasing reactivity. Thus K is more reactive than Na.
Q36. Which atom has smaller atomic radius: Na or Mg?
Answer: B. Mg (to the right of Na in the same period) has a smaller radius due to greater nuclear charge pulling electrons in more strongly. Across a period radius decreases.
Q37. Which of these elements has the highest electron affinity?
Answer: A. Chlorine has high (more negative) electron affinity among these as it readily gains an electron to attain noble gas configuration. Noble gas Ar has near zero electron affinity.
Q38. The trend of metallic character across a period is:
Answer: B. Metallic character decreases across a period (elements become less willing to lose electrons). Non-metallic character increases.
Q39. Which of the following shows correct order of increasing atomic radius?
Answer: A. K (largest), Ca, Ar (smallest) in Period 4 left to right atomic radius decreases. K has lowest nuclear pull among them.
Q40. Which factor is primarily responsible for trends in atomic size across a period?
Answer: A. Increasing nuclear charge (protons) pulls electrons closer, reducing size across a period. Shielding remains nearly constant across a period.
Q41. The first ionisation energy of oxygen is slightly less than that of nitrogen because:
Answer: B. Oxygen’s 2p⁴ has paired electrons causing repulsion and making it slightly easier to remove one electron compared to half-filled 2p³ of nitrogen. Half-filled subshells are relatively stable.
Q42. Which of following has highest electronegativity?
Answer: B. Oxygen is most electronegative among these. Electronegativity increases across period and decreases down group.
Groups & Group Properties (Q43–Q52)
Q43. Which group consists of highly reactive metals that have one valence electron?
Answer: C. Group 1 are alkali metals with ns¹ configuration, highly reactive and soft. They readily lose one electron to form M⁺.
Q44. Which of the following reactions represents alkali metal with water?
Answer: A. Alkali metals react vigorously with water producing hydroxide and hydrogen gas. This demonstrates their strong reducing nature.
Q45. Which group contains halogens?
Answer: B. Halogens are in Group 17, have 7 valence electrons and form X⁻ ions readily. They are highly reactive non-metals.
Q46. Reactivity of halogens ______ down the group.
Answer: B. Reactivity decreases down Group 17 as atoms get larger and attraction for incoming electron decreases. Fluorine is the most reactive halogen.
Q47. Which of the following is true about Group 2 elements?
Answer: C. Alkaline earth metals (Group 2) commonly form +2 ions by losing two valence electrons. Examples: Mg²⁺, Ca²⁺.
Q48. Which element is most likely to displace bromine from potassium bromide solution?
Answer: C. Fluorine (most reactive halogen) will displace bromide. More reactive halogen displaces less reactive halide.
Q49. Which group contains the noble gases?
Answer: A. Noble gases are in Group 18 with complete valence shells and low reactivity. He, Ne, Ar, etc.
Q50. Which of the following properties increases down a group?
Answer: C. Atomic radius increases down a group as new shells are added. Shielding effect increases.
Q51. Which of these group trends is correct for Group 1 elements?
Answer: B. Reactivity increases down Group 1. Melting point and ionisation energy generally decrease down the group. Valence electron is further from nucleus and more easily lost.
Q52. Which of the following is true about elements in same group?
Answer: B. Elements in a group have the same number of valence electrons, leading to similar chemical properties. Other properties vary due to shell differences.
Metals, Non-metals & Transition Elements (Q53–Q60)
Q53. Which of the following is NOT a typical property of metals?
Answer: C. Metals are generally lustrous, not dull. Dull surface is more typical of non-metals.
Q54. Which of the following oxides is amphoteric?
Answer: C. Al₂O₃ is amphoteric — reacts with acids and bases. It shows both acidic and basic behaviour.
Q55. Transition elements often form coloured compounds because:
Answer: B. Partially filled d-orbitals enable d–d transitions that absorb visible light, producing colours. Example: Cu²⁺ (blue), Fe³⁺ (yellow-brown).
Q56. Which compound is ionic in nature?
Answer: C. NaCl is ionic, formed by metal (Na) and non-metal (Cl) via electron transfer. Ionic compounds form lattices of cations and anions.
Q57. Which of the following is a characteristic of non-metals?
Answer: C. Non-metals often form acidic oxides (e.g., CO₂, SO₂). They tend to gain electrons and form covalent bonds.
Q58. Which of the following elements is a transition metal?
Answer: B. Scandium (Sc) is a transition metal (d-block). It has partially filled d-orbitals in its common oxidation states.
Q59. Which of the following statements about NaCl and SiO₂ is correct?
Answer: C. NaCl is ionic (salt), while SiO₂ is a covalent network solid (high melting point, insoluble). Different bonding gives differing properties.
Q60. Which of the following elements would you predict to have the highest melting point?
Answer: C. Silicon (covalent network solid) has very high melting point compared to metals like Na and Mg and noble gas Ar. Network covalent bonding requires breaking many strong covalent bonds.