Periodic Classification of Elements – Short Answer Type Questions
Class 10
Chemistry — Chapter 14
Board: CBSE
Chapter Standard: NCERT Class 10
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History & Early Attempts (Q1–Q10)
Q1. What is a triad according to Döbereiner?
A triad is a group of three elements with similar chemical properties where the atomic mass of the middle element is approximately the average of the other two (example: Li, Na, K).
Q2. State Newlands' law of octaves.
Newlands' law of octaves arranged elements by increasing atomic mass and noticed that every eighth element showed similar properties, like musical octaves. It worked well for lighter elements but failed later.
Q3. Give one major limitation of Newlands' law.
It failed for elements beyond calcium and did not leave spaces for undiscovered elements or account for variable properties; it grouped dissimilar elements together.
Q4. What were the key features of Mendeleev's periodic table?
Mendeleev arranged elements by increasing atomic mass, grouped elements with similar properties into vertical columns, left gaps for undiscovered elements, and predicted properties of some unknown elements.
Q5. How did Mendeleev explain anomalies where order by atomic mass did not match properties?
Mendeleev sometimes reversed the order of elements based on properties rather than strict atomic mass to keep similar elements in the same group (e.g., Te and I).
Q6. State one advantage of Mendeleev’s periodic table.
It predicted the existence and properties of elements not yet discovered (e.g., gallium, germanium), which was later confirmed by experiments.
Q7. Mention one limitation of Mendeleev's periodic table.
It could not explain the position of isotopes and relied on atomic mass, which sometimes led to incorrect placements.
Q8. Which discovery led to the revision of Mendeleev's table into the modern periodic table?
The discovery of the proton and the concept of atomic number (number of protons) led to arranging elements by atomic number instead of atomic mass.
Q9. What is meant by 'period' in the periodic table?
A period is a horizontal row of the periodic table in which properties change gradually from left to right as atomic number increases.
Q10. What is meant by a 'group' in the periodic table?
A group is a vertical column of elements that have similar chemical properties and the same number of valence electrons in their outer shell.
Modern Periodic Law & Long Form (Q11–Q20)
Q11. State the modern periodic law.
The modern periodic law states: "Properties of elements are periodic functions of their atomic numbers." This means elements with similar properties recur at regular intervals when arranged by atomic number.
Q12. What is the long form of the periodic table?
The long form (modern periodic table) is the standard layout with 18 vertical groups and 7 horizontal periods arranged by increasing atomic number, including separate rows for lanthanoids and actinoids.
Q13. How many groups and periods are there in the modern periodic table?
There are 18 groups and 7 periods in the modern periodic table.
Q14. Why are noble gases placed separately in Group 18?
Noble gases have complete valence shells (stable electronic configuration) and are chemically inert; hence they occupy Group 18 as a distinct family of unreactive gases.
Q15. Where are metals, non-metals and metalloids located in the periodic table?
Metals are generally on the left and in the centre, non-metals on the top-right, and metalloids (elements with intermediate properties) lie along a zig-zag diagonal between metals and non-metals.
Q16. What are transition elements and where are they found?
Transition elements are d-block elements found in groups 3–12; they typically have partially filled d-orbitals and show variable oxidation states and coloured compounds.
Q17. Why are lanthanoids and actinoids placed separately?
They are f-block elements with electrons filling the 4f and 5f orbitals; placing them separately keeps the table compact and highlights their unique properties.
Q18. Define atomic number.
Atomic number (Z) is the number of protons in an atom's nucleus and determines the identity and position of an element in the periodic table.
Q19. What is isotopes' effect on the periodic table?
Isotopes are atoms of the same element with different mass numbers. They have the same atomic number so they occupy the same position in the periodic table.
Q20. Explain why elements in the same group show similar properties.
Elements in the same group have the same number of valence electrons, causing them to exhibit similar chemical reactivity and bonding patterns.
Electronic Configuration & Valency (Q21–Q28)
Q21. How does electronic configuration determine the period of an element?
The period number equals the highest principal quantum number (n) of the element's occupied electron shell. For example, elements with outer electrons in n=3 belong to period 3.
Q22. Write the electronic configuration of sodium (Na, Z=11).
Na: 1s² 2s² 2p⁶ 3s¹.
Q23. Write the electronic configuration of chlorine (Cl, Z=17).
Cl: 1s² 2s² 2p⁶ 3s² 3p⁵.
Q24. Define valency with an example.
Valency is the combining capacity of an element, equal to the number of electrons gained, lost or shared to attain noble gas configuration. Example: Na has valency 1 (loses one electron), O has valency 2 (gains two electrons).
Q25. How is valency related to group number for main-group elements?
For p-block and s-block main-group elements: valency often equals the group number for metals or (8 - group number) for non-metals to attain octet. Example: Group 1 elements have valency 1; Group 16 elements typically have valency 2.
Q26. What is the valency of carbon and why?
Carbon has valency 4 because it has four valence electrons (2s²2p²) and needs four more electrons to complete its octet, allowing it to form four covalent bonds.
Q27. How do you determine valency from electronic configuration for an element with atomic number 12 (Mg)?
Mg (Z=12) configuration: 1s² 2s² 2p⁶ 3s². It has two valence electrons in 3s, so valency = 2 (it tends to lose two electrons).
Q28. What is the electronic configuration pattern for Group 18 elements?
Group 18 elements (noble gases) have complete outer shells: e.g., He: 1s², Ne: 1s²2s²2p⁶, Ar: 1s²2s²2p⁶3s²3p⁶ — full valence shells make them stable and unreactive.
Periodic Trends (Q29–Q36)
Q29. How does atomic radius change across a period?
Across a period (left → right), atomic radius decreases because increasing nuclear charge pulls electrons closer to the nucleus, reducing atomic size.
Q30. How does atomic radius change down a group?
Down a group, atomic radius increases because new electron shells are added, increasing the distance between the nucleus and outer electrons.
Q31. Define ionization energy.
Ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms (first ionization energy for removing the first electron).
Q32. Trend of ionization energy in the periodic table?
Ionization energy generally increases across a period (more nuclear charge, harder to remove electrons) and decreases down a group (outer electrons are farther and more shielded).
Q33. What is electron affinity in simple terms?
Electron affinity is the energy change when an atom in the gas phase gains an electron to form an anion; atoms that gain electrons easily have higher (more negative) electron affinity.
Q34. How does electronegativity vary in the periodic table?
Electronegativity increases across a period (towards fluorine) and decreases down a group. Fluorine is the most electronegative element.
Q35. Why are alkali metals highly reactive?
Alkali metals (Group 1) have one valence electron that is easily lost to achieve a stable noble gas configuration, making them highly reactive, especially with water and halogens.
Q36. Why does the reactivity of halogens decrease down the group?
Reactivity of halogens decreases down the group because the ability to gain an electron decreases as atomic size increases and nuclear attraction on incoming electrons reduces.
Metals, Non-metals & Group Properties (Q37–Q43)
Q37. List three general physical properties of metals.
Metals are generally lustrous (shiny), good conductors of heat and electricity, malleable and ductile, and usually solid at room temperature (except mercury).
Q38. List three chemical properties of non-metals.
Non-metals tend to gain electrons to form anions, form acidic oxides, and often are poor conductors of heat and electricity.
Q39. How does reactivity change for alkali metals down the group?
Reactivity of alkali metals increases down the group because the valence electron is farther from the nucleus and can be lost more easily.
Q40. Describe the general trend of metallic character in the periodic table.
Metallic character increases down a group and decreases across a period from left to right (elements on the left and bottom are most metallic).
Q41. Give one chemical property of noble gases.
Noble gases are chemically inert due to complete valence electron shells; they rarely form compounds (few exceptions under extreme conditions).
Q42. What is the typical oxidation state of Group 2 elements?
Group 2 elements (alkaline earth metals) typically have an oxidation state of +2 because they lose two valence electrons to achieve a stable configuration.
Q43. Which group contains elements that form coloured compounds and show variable oxidation states?
The transition elements (d-block, groups 3–12) often form coloured compounds and exhibit multiple oxidation states.
Miscellaneous Questions & Applications (Q44–Q50)
Q44. Why is hydrogen placed separately in the periodic table?
Hydrogen has properties that resemble both alkali metals (one valence electron) and halogens (needs one electron to complete its shell), so it is often placed separately at the top of Group 1 or alone.
Q45. How did Mendeleev’s predictions help in validating his periodic table?
Mendeleev predicted properties and masses of undiscovered elements (like eka-aluminium → gallium); their later discovery with similar properties validated his table.
Q46. Give an example where modern periodic law corrects a limitation of Mendeleev’s table.
Modern periodic law arranges elements by atomic number, which resolves the placement issue of isotopes and anomalies like Te and I that were misplaced when using atomic mass ordering.
Q47. What general rule helps to locate metals and non-metals in a period?
In a period, elements on the left are generally metals and on the right are non-metals, with metalloids near the diagonal dividing line.
Q48. How is the periodic table useful in predicting chemical behaviour?
The periodic table shows periodic trends (valency, reactivity, atomic size) so one can predict how an element will react, its likely ions, and bonding behaviour without detailed experiments.
Q49. Name an element with filled d-orbitals and state one typical property.
Copper (Cu) has filled 3d¹⁰ in its common state and typically forms coloured compounds and conducts electricity well; transition metals commonly show variable oxidation states.
Q50. How would you classify an element that forms ionic compounds and has high melting point?
Such an element is typically metallic (often an alkali/alkaline earth or transition metal) that forms ionic lattices — characteristic of metals with strong electrostatic forces and high melting points.