ICSE – Grade 10 – Chemistry – Ch 01 – QA

1. Explain the Modern Periodic Law. How does it differ from Mendeleev’s Periodic Law?
    Modern Periodic Law states that the physical and chemical properties of elements are periodic functions of their atomic numbers. Mendeleev’s law was based on atomic masses. Modern law resolved anomalies like the position of isotopes and certain element pairs by using atomic number as the basis.

    

2. What are periodic properties? List and define any four with trends in periods and groups.
    Periodic properties are properties that repeat at regular intervals in the periodic table. Four such properties are:

    

• Atomic size: decreases across a period, increases down a group.
• Ionization energy: increases across a period, decreases down a group.
• Electronegativity: increases across a period, decreases down a group.
• Metallic character: decreases across a period, increases down a group.

3. Describe how atomic size varies across a period and down a group. Give reasons.
   Across a period, atomic size decreases due to increased nuclear charge pulling electrons closer. Down a group, atomic size increases due to the addition of electron shells.

    

4. Define ionisation potential. How does it vary in a period and in a group? Why?
   Ionisation potential is the energy required to remove an electron from an isolated gaseous atom. It increases across a period due to stronger nuclear attraction, and decreases down a group due to larger atomic size and shielding effect.

    

5. What is electronegativity? State its trends in the periodic table and explain.
   Electronegativity is the tendency of an atom to attract electrons in a chemical bond. It increases across a period due to increased nuclear charge, and decreases down a group due to increased atomic size and shielding.

    

6. Differentiate between Mendeleev’s Periodic Table and Modern Periodic Table.

    

• Mendeleev’s table is based on atomic mass; Modern is based on atomic number.
• Mendeleev left gaps for undiscovered elements; Modern table is complete.
• In Mendeleev’s table, isotopes posed a problem; in the modern table, they don’t.
• Modern table explains periodicity better in terms of electronic configuration.

7. What are alkali metals? Discuss their position, electronic configuration, and properties.
   Alkali metals are Group 1 elements like Li, Na, K. They have 1 valence electron (ns1 configuration), are highly reactive, soft, have low ionization energies, and form strong bases with water.

    

8. Why are noble gases placed in Group 18? Describe their properties.
   Noble gases have full valence shells, making them stable and chemically inert. They are colorless, monoatomic gases with very low reactivity.

    

9. How do metallic and non-metallic characters vary in the periodic table?
   Metallic character decreases across a period and increases down a group. Non-metallic character increases across a period and decreases down a group due to variations in ionization energy and electronegativity.

    

10. Why are halogens highly reactive? Mention any three properties of halogens.
    Halogens have 7 valence electrons and need one more to complete their octet. They are highly electronegative, form diatomic molecules, and have high reactivity and electron affinity.

     

11. Discuss the importance of the periodic table in modern chemistry.
    It helps in classification of elements, prediction of properties, understanding trends, and organizing chemical behavior systematically for ease of study and research.

     

12. What is electron affinity? State and explain its periodic trends.
    Electron affinity is the energy released when an atom gains an electron. It becomes more negative across a period and less negative down a group.

     

13. Explain the concept of periodicity with reference to the modern periodic table.
    Periodicity refers to the repetition of properties at regular intervals due to similar valence shell configurations when elements are arranged by atomic number.

     

14. What is the basis of classification in the modern periodic table?
    Elements are arranged in increasing order of atomic number, which reflects their electronic configuration and periodic recurrence of chemical properties.

     

15. Describe the trends in valency across periods and down groups.
    Across a period, valency first increases then decreases. Down a group, elements have the same valency as they have the same number of valence electrons.

     

16. Why are elements in the same group similar in chemical properties?
    Because they have the same number of valence electrons, leading to similar bonding and chemical behavior.

     

17. State two advantages of the Modern Periodic Table over Mendeleev’s.
    It is based on atomic number, resolving isotopic and placement issues, and shows periodicity more consistently due to electronic configuration.

     

18. Why does atomic size decrease from left to right across a period?
    Due to increasing nuclear charge which pulls the electron cloud closer to the nucleus, reducing atomic radius.

     

19. Give reasons for the high reactivity of alkali metals.
    They have a single valence electron that is easily lost, low ionization energy, and large atomic size.

     

20. Give reasons for the low reactivity of noble gases.
    They have complete outer shells, are energetically stable, and do not tend to gain or lose electrons.

     

21. Define and compare the electron affinities of Group 17 and Group 1 elements.
    Group 17 has high (negative) electron affinity due to need to complete octet; Group 1 has low (less negative) affinity as adding electrons increases instability.

     

22. Why is fluorine more reactive than chlorine?
    Fluorine is smaller in size, has higher electronegativity and electron affinity, making it more reactive.

     

23. Explain the term ‘group’ in the periodic table with an example.
    Group is a vertical column with elements having same number of valence electrons; e.g., Group 1 – Li, Na, K.

     

24. Explain the term ‘period’ with reference to the periodic table.
    A period is a horizontal row in which elements have the same number of electron shells; e.g., Period 3 has Na to Ar.

     

25. How does the number of shells change as we go down a group?
    It increases by one for each element, resulting in larger atomic size and lower ionization energy.

     

26. Why do elements in the same group have different physical properties but similar chemical properties?
    Physical properties depend on atomic mass and size; chemical properties depend on valence electrons which are same in a group.

     

27. State the general electronic configuration of alkali metals and halogens.
    Alkali metals: ns1; Halogens: ns2 np5.

     

28. Why are halogens non-metals though they exist close to metals in the periodic table?
    They have high ionization energy and high electronegativity, tending to gain electrons.

     

29. Compare metallic and non-metallic properties of Period 3 elements.
    Na, Mg, Al are metallic; Si is metalloid; P, S, Cl are non-metallic. Metallic nature decreases across the period.

     

30. Why does ionization potential increase across a period?
    Due to increasing nuclear charge and decreasing atomic size, it becomes harder to remove an electron.

     

31. What is meant by shielding effect and how does it affect periodic properties?
    It is the repulsion of outer electrons by inner electrons, reducing nuclear pull. It increases atomic size and reduces ionization energy.

     

32. Compare reactivity of alkali metals and halogens.
    Alkali metals lose electrons easily; halogens gain electrons easily. Both are highly reactive but in opposite ways.

     

33. What are metalloids? Give two examples.
    Elements that show properties of both metals and non-metals. Examples: Silicon, Arsenic.

     

34. Why do elements in the same period have different properties?
    Because they have different numbers of valence electrons and hence form different types of bonds and compounds.

     

35. What is the significance of the position of an element in the periodic table?
    It indicates its electronic configuration, group and period, chemical reactivity, and bonding behavior.

     

36. Why are Group 1 elements stored under kerosene oil?
    To prevent reaction with air or moisture, as they are highly reactive metals.

     

37. What is a periodic trend? Give any three with their reasons.
    A recurring pattern in properties:

     

• Atomic size decreases across a period.
• Ionization energy increases across a period.
• Metallic character decreases across a period.

38. Why does the metallic character increase down a group?
    Because atoms become larger and lose electrons more easily due to weaker nuclear attraction.

     

39. Explain why the second period has only 8 elements.
    Because the second shell can hold a maximum of 8 electrons (2s and 2p orbitals).

     

40. Define valency. How does it change in a period and a group?
    Valency is the combining capacity of an atom. It increases from 1 to 4 then decreases to 0 in a period; remains same in a group.

     

41. Why do halogens form negative ions?
    Because they have 7 valence electrons and need one more to achieve noble gas configuration.

     

42. Give reasons for placing hydrogen separately in the periodic table.
    Hydrogen resembles both Group 1 and Group 17 elements in properties, but is unique in behavior and electronic structure.

     

43. Why is argon placed in Group 18 though its atomic mass is less than potassium?
    Because it has a complete octet and belongs to noble gases; placement is based on atomic number, not mass.

     

44. How is periodicity related to electron configuration?
    Elements with similar valence shell configuration show similar properties, leading to periodic trends.

     

45. Explain why neon and argon do not form compounds easily.
    They have complete outer shells, making them chemically inert and not tending to gain, lose or share electrons.

     

46. What is a shell? How does the number of shells change in a period and group?
    A shell is an energy level where electrons reside. Across a period, number of shells is constant; down a group, it increases.

     

47. Discuss the relationship between atomic number and chemical properties.
    Atomic number determines electronic configuration, which in turn defines chemical properties of an element.

     

48. How does electron affinity influence the chemical behavior of an element?
    Higher electron affinity means the atom readily gains electrons, often forming anions and reacting with electropositive elements.

     

49. Why does atomic radius increase down a group despite increased nuclear charge?
    Due to addition of electron shells and increased shielding, which outweigh the effect of nuclear charge.

     

50. How does the periodic table help in predicting the types of chemical bonds formed by elements?
 Based on electronegativity and valence electrons, one can predict whether elements will form ionic or covalent bonds.