Explain the trends of atomic size, ionization energy, and electron affinity in the periodic table with suitable examples.

In the periodic table, atomic size decreases as you move across a period from left to right. This happens because as the nuclear charge increases, electrons are pulled closer to the nucleus, reducing the size of the atom. For example, the atomic radius of sodium (Na) is larger than that of chlorine (Cl).

Ionization energy, the energy required to remove an electron from an atom, increases across a period due to the increased nuclear charge that holds the electrons more tightly. For example, the ionization energy of magnesium (Mg) is higher than that of sodium (Na).

Electron affinity, the energy change when an atom gains an electron, becomes more negative (indicating a greater affinity for electrons) across a period. Chlorine has a more negative electron affinity than sodium, meaning it more readily gains an electron to form a stable anion.

Related Questions:

  1. Arrange the elements in each of the following in order of decreasing shielding effect:
  2. Specify which of the following elements you would expect to have the greatest electron affinity: S, P, Cl
  3. Electronic configurations of some elements are given below. Group the elements in pairs that would represent similar chemical properties:
  4. Arrange the elements in groups and periods from Q. No. 6:
  5. For normal elements, the number of valence electrons of an element is equal to the group number. Find the group number of the following elements:
  6. Write the valence shell electronic configuration for the following groups:
  7. Write the valence shell electronic configuration of the atoms of the following elements:
  8. In which block, group, and period in the periodic table would you place each of the following elements with the given electronic configurations?
  9. Describe the significance of Mendeleev’s periodic law and how it helped in the development of the modern periodic table.
  10. Explain the position and properties of alkali metals and halogens in the periodic table. Why do they react so easily with each other?
  11. Write the valence shell electronic configuration of an element present in the 3rd period and Group IIA.
  12. Define halogens.
  13. Which atom has a higher shielding effect, Li or Na?
  14. Explain why Na has higher ionization energy than K.
  15. Alkali metals belong to the S-block in the periodic table. Why?
  16. Arrange the elements in each of the following groups in order of increasing ionization energy:
  17. Why does atomic radius decrease across a period from left to right?
  18. Why do noble gases have high ionization energy?
  19. Why are alkali metals highly reactive?
  20. What is periodicity in the periodic table?