The s & P Block Elements -

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

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Correct Option: C
Option A Explained: Incorrect. Electron affinity is the energy released when an electron is added to a neutral atom, not removed. It is the opposite process of what the question asks.
Option B Explained: Incorrect. Lattice energy is the energy needed to break an ionic solid into its separate gaseous ions. It has nothing to do with removing an electron from a single atom.
Option C Explained: Correct. Ionization energy is defined exactly as the energy needed to remove an electron from an isolated gaseous atom in its neutral state.
Option D Explained: Incorrect. Crystal energy is not the standard term used for electron removal. This term is sometimes loosely used for lattice energy, which is unrelated here.

2. The elements for which the value of ionization energy is low can:
A. Gain electrons readily
B. Gain electrons with difficulty
C. Lose electron less readily
D. Lose electron readily

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Correct Option: D
Option A Explained: Incorrect. Low ionization energy means the atom holds its electrons loosely, so it prefers losing electrons, not gaining them.
Option B Explained: Incorrect. Gaining electrons with difficulty is linked to metals in general, but this is not the direct meaning of low ionization energy.
Option C Explained: Incorrect. This is the opposite of the truth. Low ionization energy means electrons leave the atom easily, not with difficulty.
Option D Explained: Correct. Low ionization energy means very little energy is needed to remove an electron, so such elements lose electrons readily. This is a typical property of metals.

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

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3. The diagram below is a plot of melting points of elements of second period against their atomic numbers. Lithium and fluorine are placed at the extreme ends of the plot. On the basis of melting points where would you place carbon among the empty slots on the plot?
A. 1
B. 4
C. 2
D. 3

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Correct Option: A
Option A Explained: Correct. Carbon has a giant covalent (network) structure with very strong covalent bonds throughout, giving it the highest melting point in period 2. So it is placed at the topmost position on the plot.
Option B Explained: Incorrect. Position 4 would represent a much lower melting point, which does not match carbon’s extremely high melting point.
Option C Explained: Incorrect. Position 2 is too low for carbon since carbon’s melting point is far higher than most other period 2 elements.
Option D Explained: Incorrect. Position 3 also underestimates carbon’s melting point, which is the highest among all period 2 elements.
Note: This answer depends on the original diagram in the source image. Please confirm the position numbers match your image before printing.

4.Which one remains same along a period?
A. Atomic radius
B. Number of shells (orbits)
C. Melting point
D. Electrical conductivity

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Correct Option: B
Option A Explained: Incorrect. Atomic radius decreases across a period because nuclear charge increases while the shell number stays the same, pulling electrons closer.
Option B Explained: Correct. All elements in the same period have the same number of electron shells since a new period begins only when a new shell starts filling.
Option C Explained: Incorrect. Melting point changes significantly across a period, first increasing then decreasing, so it does not remain constant.
Option D Explained: Incorrect. Electrical conductivity changes a lot across a period as elements go from metals to metalloids to non-metals.

5. More the ionization energy of an element:
A. More the electro positivity
B. Less the metallic character
C. More the reducing power
D. Bigger the atomic radius

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Correct Option: B
Option A Explained: Incorrect. High ionization energy means the atom holds electrons tightly, which is the opposite of electropositive (metallic) behavior.
Option B Explained: Correct. High ionization energy means electrons are hard to remove, which is a key feature of non-metals, so metallic character decreases.
Option C Explained: Incorrect. Reducing power comes from easily losing electrons, so high ionization energy actually lowers reducing power.
Option D Explained: Incorrect. High ionization energy is generally linked with smaller atomic radius, not bigger, since electrons are held closer to the nucleus.

6. Along a period, atomic radius decreases. This gradual decrease in radius is due to:
A. Increase in number of shells
B. Increase in number of protons in the nucleus
C. Melting and boiling points first decrease then increase
D. Melting and boiling points first increase then decrease

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Correct Option: B
Option A Explained: Incorrect. The number of shells stays the same across a period; a new shell only appears when a new period begins.
Option B Explained: Correct. As protons increase across a period, nuclear charge increases, pulling the same shell of electrons closer, which shrinks atomic radius.
Option C Explained: Incorrect. This describes melting/boiling point trend, not the actual reason behind radius decrease.
Option D Explained: Incorrect. This is also about melting/boiling points, which is unrelated to why radius decreases across a period.

7. What is the trend of melting and boiling points of the elements of short periods as we move from left to right in a periodic table?
A. Melting and boiling points decrease gradually
B. Melting and boiling points first decrease then increase
C. Melting and boiling points increase gradually
D. Melting and boiling points first increase then decrease

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Correct Option: D
Option A Explained: Incorrect. Melting and boiling points do not fall gradually throughout the period; they rise first due to stronger metallic and covalent bonding.
Option B Explained: Incorrect. This is the reverse of the actual trend seen in short periods.
Option C Explained: Incorrect. Points do not increase throughout; they drop sharply after group 14 due to weaker molecular structures.
Option D Explained: Correct. Melting and boiling points rise up to group 14 (strong giant structures like carbon/silicon) then fall sharply as we reach simple molecular non-metals and noble gases.

8. The trends in melting points of the elements of 3rd period are depicted in figure below. The sharp decrease observed from ‘Si’ to ‘P’ is due to:
A. Decrease in atomic radius from ‘Si’ to ‘P’
B. Change in bonding and structure of two elements
C. Different densities of two elements
D. Increase in electron density from ‘Si’ to ‘P’

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Correct Option: B
Option A Explained: Incorrect. Atomic radius change alone is too small to cause such a huge drop in melting point.
Option B Explained: Correct. Silicon has a giant covalent network structure (strong bonds throughout), while phosphorus exists as small P4 molecules held by weak van der Waals forces, causing the sharp melting point drop.
Option C Explained: Incorrect. Density differences are a result of structure, not the direct cause of the melting point change.
Option D Explained: Incorrect. Electron density is not the reason; the real cause is the shift from a giant structure to a simple molecular structure.

9. Arrange the following elements according to the trends of ionization energies: C, N, Ne, B
A. Ne < N < C < B
B. B < C < N < Ne
C. B < N < C < Ne
D. Ne < B < C < N

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Correct Option: B
Option A Explained: Incorrect. This order is completely reversed compared to the actual increasing trend of ionization energy.
Option B Explained: Correct. Ionization energy generally increases across a period: B (801) < C (1086) < N (1402) < Ne (2081) kJ/mol, matching this order exactly.
Option C Explained: Incorrect. Nitrogen’s ionization energy is higher than carbon’s, so N cannot come before C in an increasing order.
Option D Explained: Incorrect. Neon has the highest ionization energy among these four, so it cannot be placed first.

10. Keeping in view the size of atoms, which order is correct?
A. N > C
B. Ar > Cl
C. P > Si
D. Li > Be

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Correct Option: D
Option A Explained: Incorrect. Carbon is actually larger than nitrogen since atomic radius decreases across a period, and carbon comes before nitrogen.
Option B Explained: Incorrect. Chlorine’s covalent radius is generally larger than argon’s, so this order is wrong.
Option C Explained: Incorrect. Silicon is bigger than phosphorus because atomic radius decreases from left to right across period 3.
Option D Explained: Correct. Lithium comes before beryllium in period 2, and since radius decreases across a period, Li is indeed larger than Be.

11. Which one of the following will have the smallest radius?
A. Al³⁺
B. Mg²⁺
C. Si⁴⁺
D. Na⁺

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Correct Option: C
Option A Explained: Incorrect. Al³⁺ has 10 electrons and 13 protons, giving a stronger pull than Na⁺ or Mg²⁺, but it is still bigger than Si⁴⁺.
Option B Explained: Incorrect. Mg²⁺ has fewer protons than Al³⁺ and Si⁴⁺, so the nuclear pull on its 10 electrons is weaker, making it larger than these two ions.
Option C Explained: Correct. All four ions have the same number of electrons (isoelectronic, 10 electrons), but Si⁴⁺ has the highest number of protons (14), pulling the electron cloud in the most, giving the smallest radius.
Option D Explained: Incorrect. Na⁺ has only 11 protons, which is the lowest among these four ions, so it has the weakest pull and the largest radius among them.

12. The ionic radius of fluoride ion is:
A. 72 pm
B. 136 pm
C. 95 pm
D. 157 pm

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Correct Option: B
Option A Explained: Incorrect. 72 pm is close to the atomic radius of fluorine, not the larger ionic radius of the fluoride ion after gaining an electron.
Option B Explained: Correct. When fluorine gains one electron to form F⁻, the extra electron increases electron-electron repulsion, expanding the radius to about 136 pm.
Option C Explained: Incorrect. 95 pm is a value closer to another halide or unrelated ion, not the standard fluoride ionic radius.
Option D Explained: Incorrect. 157 pm is too large a value and does not match the known ionic radius of fluoride.

13. Which one of the following physical properties decreases along a period from left to right?
A. Electron affinity
B. Non-metallic character
C. Atomic radius
D. Melting point up to group IVA

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Correct Option: C
Option A Explained: Incorrect. Electron affinity increases from left to right in a period in periodic table.
Option B Explained: Incorrect. Non-metallic property of elements increases along a period from left to right in periodic table.
Option C Explained: Correct. Atomic radius shows a clear, steady decrease across period 3 due to increasing nuclear charge.
Option D Explained: Incorrect. Melting point trend is not smooth; it rises then falls sharply, unlike a simple continuous decrease.

14. The following sketch shows the melting point of eight elements with consecutive atomic numbers. Which element is silicon?

A. A
B. C
C. B
D. D

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15. Which one of the following anions has the smallest ionic radius?
A. O^2-
B. Cl^–
C. F^–
D. N^3-

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16. How many elements are there in p-block of periodic table?
A. 14
B. 6
C. 42
D. 36

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1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

1. Energy required to remove an electron from gaseous neutral atom is:
A. Electron affinity
B. Lattice energy
C. Ionization energy
D. Crystal energy

Show Answer

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