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201. Why is the radius of a cation always smaller than its parent atom?
ⓐ. Because mass decreases when electrons are lost
ⓑ. Because nuclear charge decreases after losing electrons
ⓒ. Because the number of protons remains the same but electrons are fewer
ⓓ. Because cations contain neutrons only
Correct Answer: Because the number of protons remains the same but electrons are fewer
Explanation: When an atom loses electrons to form a cation, the proton number stays unchanged. The stronger effective nuclear charge pulls the remaining electrons closer to the nucleus. Electron–electron repulsion also decreases, leading to a smaller radius.
202. Why is the radius of an anion always larger than its parent atom?
ⓐ. Because protons are lost during anion formation
ⓑ. Because electrons are gained, increasing repulsion
ⓒ. Because the nucleus expands
ⓓ. Because the number of shells decreases
Correct Answer: Because electrons are gained, increasing repulsion
Explanation: When an atom gains electrons, repulsion between electrons in the outer shell increases. The nuclear charge is now distributed over more electrons, reducing effective nuclear attraction per electron, so the electron cloud expands and radius increases.
203. Compare the size of Na (neutral atom) and Na⁺ (cation).
ⓐ. Na⁺ is larger than Na
ⓑ. Na⁺ is smaller than Na
ⓒ. Both have the same radius
ⓓ. Depends on isotopes
Correct Answer: Na⁺ is smaller than Na
Explanation: Sodium (Na) has atomic radius \~186 pm, while Na⁺ has \~102 pm. Losing one electron removes the outermost shell (3s¹), decreasing size drastically. The effective nuclear pull also strengthens.
204. Compare the size of Cl (neutral atom) and Cl⁻ (anion).
ⓐ. Cl⁻ is smaller than Cl
ⓑ. Cl⁻ is larger than Cl
ⓒ. Both are the same
ⓓ. Depends on bonding
Correct Answer: Cl⁻ is larger than Cl
Explanation: Chlorine atom has covalent radius \~99 pm, while Cl⁻ ionic radius \~181 pm. Extra electron increases repulsion in the valence shell, expanding the cloud. Thus, anions are always larger than their parent atoms.
205. Arrange the following species in decreasing order of radius: Na, Na⁺, Cl, Cl⁻.
ⓐ. Cl⁻ > Cl > Na > Na⁺
ⓑ. Na⁺ > Na > Cl > Cl⁻
ⓒ. Na > Cl > Cl⁻ > Na⁺
ⓓ. Na⁺ > Cl⁻ > Na > Cl
Correct Answer: Cl⁻ > Cl > Na > Na⁺
Explanation: Anions (Cl⁻) are largest due to extra repulsion. Neutral atoms (Cl and Na) are smaller than anions. Cations (Na⁺) are the smallest due to electron loss and increased effective nuclear charge.
206. Which has a smaller radius: Mg²⁺ or Na⁺?
ⓐ. Na⁺
ⓑ. Mg²⁺
ⓒ. Both are equal
ⓓ. Cannot be predicted
Correct Answer: Mg²⁺
Explanation: Both Na⁺ (102 pm) and Mg²⁺ (72 pm) belong to the same period and are isoelectronic (10 electrons each). However, Mg²⁺ has more protons (Z = 12 vs 11), pulling electrons more strongly, so its radius is smaller.
207. Which has the largest radius among the following: O²⁻, F⁻, Na⁺, Mg²⁺?
ⓐ. Mg²⁺
ⓑ. Na⁺
ⓒ. F⁻
ⓓ. O²⁻
Correct Answer: O²⁻
Explanation: All are isoelectronic species (10 electrons). The one with the smallest nuclear charge (O: Z = 8) has the weakest pull, so O²⁻ has the largest radius. Mg²⁺ (Z = 12) has the strongest pull, so it is the smallest.
208. Why is Al³⁺ smaller than Na⁺ though both are cations?
ⓐ. Because Na has more neutrons
ⓑ. Because Al³⁺ has greater nuclear charge and higher effective attraction
ⓒ. Because Na⁺ is more stable
ⓓ. Because Al is a transition element
Correct Answer: Because Al³⁺ has greater nuclear charge and higher effective attraction
Explanation: Al³⁺ has Z = 13, Na⁺ has Z = 11. Both lose electrons, but Al³⁺ loses three, leaving only 10 electrons strongly bound by 13 protons. This greater effective nuclear charge reduces size drastically (\~53 pm vs Na⁺ \~102 pm).
209. Which one of the following is the correct order of radii?
ⓐ. Ca²⁺ > K⁺ > Ar > Cl⁻
ⓑ. Cl⁻ > Ar > K⁺ > Ca²⁺
ⓒ. K⁺ > Ca²⁺ > Ar > Cl⁻
ⓓ. Ar > Cl⁻ > K⁺ > Ca²⁺
Correct Answer: Cl⁻ > Ar > K⁺ > Ca²⁺
Explanation: Cl⁻ has extra electrons → largest. Argon is neutral, intermediate. K⁺ (cation, 19p & 18e) smaller than Ar. Ca²⁺ (20p & 18e) smallest due to strongest nuclear pull.
210. Which of the following explains why cations are smaller than anions?
ⓐ. Cations have more electrons than protons, anions have fewer
ⓑ. Cations have fewer electrons than protons, anions have more
ⓒ. Cations form strong metallic bonds, anions form weak bonds
ⓓ. Cations are metals, anions are non-metals
Correct Answer: Cations have fewer electrons than protons, anions have more
Explanation: In cations, protons outnumber electrons, pulling the cloud closer → smaller radius. In anions, electrons outnumber protons, repulsion dominates and size expands. Thus, cations are always smaller and anions always larger than their parent atoms.
211. Which of the following pairs are isoelectronic species?
ⓐ. Na⁺ and Mg²⁺
ⓑ. Cl⁻ and Ar
ⓒ. O²⁻ and F⁻
ⓓ. All of the above
Correct Answer: All of the above
Explanation: Isoelectronic species have the same number of electrons but different nuclear charges. Na⁺ (11–1=10e⁻) and Mg²⁺ (12–2=10e⁻) are isoelectronic. Cl⁻ (17+1=18e⁻) and Ar (18e⁻) are isoelectronic. O²⁻ (8+2=10e⁻) and F⁻ (9+1=10e⁻) are also isoelectronic.
212. In an isoelectronic series, what is the general trend of ionic radius?
ⓐ. Increases with increasing nuclear charge
ⓑ. Decreases with increasing nuclear charge
ⓒ. Remains constant
ⓓ. Depends on atomic mass only
Correct Answer: Decreases with increasing nuclear charge
Explanation: In isoelectronic species, the number of electrons is constant, but protons increase with atomic number. Higher nuclear charge pulls the same number of electrons closer, decreasing radius.
213. Arrange the species O²⁻, F⁻, Na⁺, Mg²⁺ in order of increasing radius.
ⓐ. O²⁻ < F⁻ < Na⁺ < Mg²⁺
ⓑ. Mg²⁺ < Na⁺ < F⁻ < O²⁻
ⓒ. Na⁺ < Mg²⁺ < F⁻ < O²⁻
ⓓ. F⁻ < O²⁻ < Na⁺ < Mg²⁺
Correct Answer: Mg²⁺ < Na⁺ < F⁻ < O²⁻
Explanation: All are isoelectronic (10e⁻ each). Mg²⁺ (Z=12) has the highest nuclear charge, hence smallest radius. O²⁻ (Z=8) has the lowest nuclear charge, so it is the largest.
214. Which one among the following has the largest radius?
ⓐ. Cl⁻
ⓑ. K⁺
ⓒ. Ar
ⓓ. Ca²⁺
Correct Answer: Cl⁻
Explanation: Cl⁻ has 18 electrons with only 17 protons, leading to weak nuclear attraction and expansion of radius. K⁺ and Ca²⁺ have strong nuclear pulls due to cationic nature, while Ar is neutral and intermediate in size.
215. Compare the radii of isoelectronic species N³⁻, O²⁻, F⁻. Which is largest?
ⓐ. N³⁻
ⓑ. O²⁻
ⓒ. F⁻
ⓓ. All equal
Correct Answer: N³⁻
Explanation: All three species have 10 electrons. N³⁻ has the lowest nuclear charge (7 protons), so its pull is weakest, making it the largest. F⁻ has 9 protons, hence its radius is smaller than N³⁻ and O²⁻.
216. Arrange in increasing order of radius: S²⁻, Cl⁻, K⁺, Ca²⁺.
ⓐ. Ca²⁺ < K⁺ < Cl⁻ < S²⁻
ⓑ. K⁺ < Ca²⁺ < Cl⁻ < S²⁻
ⓒ. S²⁻ < Cl⁻ < K⁺ < Ca²⁺
ⓓ. Cl⁻ < S²⁻ < K⁺ < Ca²⁺
Correct Answer: Ca²⁺ < K⁺ < Cl⁻ < S²⁻
Explanation: All have 18 electrons. Ca²⁺ (Z=20) has the strongest pull, so it is smallest. K⁺ (Z=19) is slightly larger. Cl⁻ (Z=17) and S²⁻ (Z=16) have fewer protons and more repulsion, so they are larger, with S²⁻ being the largest.
217. In an isoelectronic series, which factor is most important in determining size?
ⓐ. Number of neutrons
ⓑ. Nuclear charge (Z)
ⓒ. Mass number
ⓓ. Type of bond
Correct Answer: Nuclear charge (Z)
Explanation: Since all species in an isoelectronic series have the same number of electrons, the deciding factor for size is the number of protons. Greater Z means greater attraction, hence smaller radius.
218. Which is smaller in size: F⁻ or Na⁺?
ⓐ. F⁻
ⓑ. Na⁺
ⓒ. Both equal
ⓓ. Cannot be compared
Correct Answer: Na⁺
Explanation: Na⁺ (10e⁻, 11p⁺) has a much stronger nuclear attraction compared to F⁻ (10e⁻, 9p⁺). Hence, Na⁺ is significantly smaller, while F⁻ expands due to weaker attraction.
219. Which ion among Al³⁺, Mg²⁺, and Na⁺ has the smallest size?
ⓐ. Na⁺
ⓑ. Mg²⁺
ⓒ. Al³⁺
ⓓ. All same size
Correct Answer: Al³⁺
Explanation: All three are isoelectronic with 10 electrons. Al³⁺ has the highest nuclear charge (Z=13), pulling electrons most strongly. Hence, Al³⁺ has the smallest radius.
220. In the isoelectronic series C⁴⁻, N³⁻, O²⁻, F⁻, which species has the smallest radius?
ⓐ. C⁴⁻
ⓑ. N³⁻
ⓒ. O²⁻
ⓓ. F⁻
Correct Answer: F⁻
Explanation: All have 10 electrons. Carbon has 6 protons, nitrogen 7, oxygen 8, and fluorine 9. Greater nuclear charge contracts size, so F⁻ (highest Z) has the smallest radius among these species.
221. What is meant by ionization enthalpy?
ⓐ. Energy released when an electron is added to an atom
ⓑ. Energy required to remove an electron from an isolated gaseous atom
ⓒ. Energy required to break a covalent bond
ⓓ. Energy released during cation formation in aqueous solution
Correct Answer: Energy required to remove an electron from an isolated gaseous atom
Explanation: Ionization enthalpy (IE) is the minimum energy required to remove one mole of electrons from one mole of gaseous atoms. For example:
$\text{M(g)} \rightarrow \text{M}^+(g) + e^-$
The first ionization enthalpy reflects how strongly an atom holds onto its outermost electron.
222. Which of the following elements has the highest first ionization enthalpy?
ⓐ. Lithium
ⓑ. Beryllium
ⓒ. Fluorine
ⓓ. Neon
Correct Answer: Neon
Explanation: Noble gases like neon have completely filled stable electronic configurations. Removing an electron requires a large amount of energy, so their ionization enthalpies are highest.
223. What is the general trend of ionization enthalpy across a period?
ⓐ. It decreases from left to right
ⓑ. It increases from left to right
ⓒ. It remains constant
ⓓ. It increases then decreases irregularly
Correct Answer: It increases from left to right
Explanation: Across a period, nuclear charge increases while atomic radius decreases. As electrons are more strongly attracted to the nucleus, it becomes harder to remove them, so ionization enthalpy increases.
224. What is the general trend of ionization enthalpy down a group?
ⓐ. It decreases
ⓑ. It increases
ⓒ. It remains constant
ⓓ. It first increases then decreases
Correct Answer: It decreases
Explanation: Down a group, atomic radius increases due to additional shells, and shielding increases. Even though nuclear charge rises, the effective pull on valence electrons decreases, making them easier to remove.
225. Which element has the lowest first ionization enthalpy among the following?
ⓐ. Sodium
ⓑ. Potassium
ⓒ. Rubidium
ⓓ. Cesium
Correct Answer: Cesium
Explanation: Cesium lies lowest in Group 1. It has the largest atomic radius and weakest hold on its outermost electron, so it requires the least energy to ionize.
226. Why is the first ionization enthalpy of nitrogen higher than that of oxygen?
ⓐ. Because oxygen has higher electronegativity
ⓑ. Because nitrogen has a stable half-filled p-subshell configuration
ⓒ. Because oxygen has a higher atomic radius
ⓓ. Because nitrogen is less reactive
Correct Answer: Because nitrogen has a stable half-filled p-subshell configuration
Explanation: Nitrogen’s configuration is 1s²2s²2p³, a stable half-filled arrangement. Removing an electron disturbs this stability, requiring more energy. Oxygen (2p⁴) lacks this stability, so its IE is lower despite higher nuclear charge.
227. Why is the second ionization enthalpy always greater than the first?
ⓐ. Because mass of ion increases
ⓑ. Because electron-electron repulsion increases
ⓒ. Because the electron is removed from a positively charged ion
ⓓ. Because number of protons decreases
Correct Answer: Because the electron is removed from a positively charged ion
Explanation: After the first electron is removed, the ion becomes positively charged. The remaining electrons experience greater nuclear attraction, making it harder to remove another electron. Hence, successive ionization enthalpies are always larger.
228. Which of the following pairs shows an exception to the general periodic trend in ionization enthalpy?
ⓐ. Be and B
ⓑ. N and O
ⓒ. Both A and B
ⓓ. C and N
Correct Answer: Both A and B
Explanation: Be (2s²) has higher IE than B (2s²2p¹) because removal from a filled s orbital is harder than from a p orbital. N (2p³ half-filled stable) has higher IE than O (2p⁴). These are classic exceptions to the trend.
229. Which group of elements has the lowest ionization enthalpies in the periodic table?
ⓐ. Group 2
ⓑ. Group 13
ⓒ. Group 17
ⓓ. Group 1
Correct Answer: Group 1
Explanation: Group 1 alkali metals (ns¹) have the lowest ionization enthalpies. Their large atomic radii and single loosely held valence electron make them highly electropositive and reactive.
230. Successive ionization enthalpies provide evidence for:
ⓐ. Atomic number of an element
ⓑ. Number of isotopes present
ⓒ. Number of valence electrons in an atom
ⓓ. Atomic radius trend
Correct Answer: Number of valence electrons in an atom
Explanation: A sudden large jump in successive ionization enthalpies indicates the removal of electrons from a stable noble gas configuration. For example, magnesium shows two low IE values (for 2 valence electrons) followed by a sharp rise, proving it belongs to Group 2.
231. What is meant by successive ionization enthalpies?
ⓐ. The energy released when electrons are added successively
ⓑ. The energy required to remove electrons one after another from the same atom
ⓒ. The energy difference between isotopes of an atom
ⓓ. The energy required to break covalent bonds in succession
Correct Answer: The energy required to remove electrons one after another from the same atom
Explanation: Successive ionization enthalpies correspond to sequential removal of electrons from a gaseous atom or ion:
$\text{M} \rightarrow \text{M}^+ + e^-$ (IE₁)
$\text{M}^+ \rightarrow \text{M}^{2+} + e^-$ (IE₂)
$\text{M}^{2+} \rightarrow \text{M}^{3+} + e^-$ (IE₃)
Each successive value is higher than the previous one due to increasing positive charge and nuclear attraction.
232. Why is the second ionization enthalpy of sodium extremely high compared to the first?
ⓐ. Because the electron is removed from the same shell
ⓑ. Because the second electron is removed from a noble gas configuration
ⓒ. Because sodium has more protons than magnesium
ⓓ. Because sodium has a very small radius
Correct Answer: Because the second electron is removed from a noble gas configuration
Explanation: Sodium’s first IE removes its single 3s¹ electron, leaving behind a stable neon core (1s²2s²2p⁶). Removing another electron requires breaking into this stable core, which requires enormous energy.
233. Which element shows a sudden jump between second and third ionization enthalpies?
ⓐ. Magnesium
ⓑ. Boron
ⓒ. Aluminium
ⓓ. Calcium
Correct Answer: Magnesium
Explanation: Magnesium has configuration \[Ne] 3s². Two low IE values correspond to removing the two valence electrons. A sudden large jump occurs at IE₃ because it would involve breaking the stable neon core. This reveals magnesium belongs to Group 2.
234. Which element shows a sudden jump between first and second ionization enthalpies?
ⓐ. Sodium
ⓑ. Magnesium
ⓒ. Aluminium
ⓓ. Silicon
Correct Answer: Sodium
Explanation: Sodium has only one valence electron (3s¹). Its first IE is low. But IE₂ requires removing an electron from the stable neon configuration, which is very high, showing it belongs to Group 1.
235. In general, what happens to successive ionization enthalpies for the same atom?
ⓐ. They remain constant
ⓑ. They gradually decrease
ⓒ. They gradually increase with sudden jumps at core removal
ⓓ. They alternate between high and low values
Correct Answer: They gradually increase with sudden jumps at core removal
Explanation: Each successive electron removal requires more energy due to higher positive charge and stronger nuclear pull. A very large jump occurs when core (inner shell) electrons must be removed after all valence electrons are gone.
236. Which factor increases ionization enthalpy across a period?
ⓐ. Increase in shielding effect
ⓑ. Increase in number of shells
ⓒ. Increase in effective nuclear charge
ⓓ. Decrease in number of protons
Correct Answer: Increase in effective nuclear charge
Explanation: Across a period, more protons are added, strengthening the nuclear pull on the same shell electrons. Atomic size decreases, making electron removal harder. Thus, ionization enthalpy increases.
237. Which factor lowers ionization enthalpy down a group?
ⓐ. Increase in nuclear charge
ⓑ. Increase in atomic radius and shielding effect
ⓒ. Increase in electronegativity
ⓓ. Decrease in number of shells
Correct Answer: Increase in atomic radius and shielding effect
Explanation: Down a group, electrons are farther from the nucleus and shielding increases. Despite higher nuclear charge, the effective attraction on valence electrons decreases, lowering ionization enthalpy.
238. Why is the ionization enthalpy of oxygen lower than that of nitrogen, despite oxygen’s higher nuclear charge?
ⓐ. Because oxygen has a smaller radius
ⓑ. Because nitrogen has a stable half-filled 2p³ configuration
ⓒ. Because oxygen has lower electronegativity
ⓓ. Because nitrogen is less metallic
Correct Answer: Because nitrogen has a stable half-filled 2p³ configuration
Explanation: Nitrogen’s half-filled p-subshell is stable, requiring more energy to remove an electron. Oxygen (2p⁴) has paired electrons with more repulsion, so it is easier to remove one, making oxygen’s IE lower than nitrogen’s.
239. Which element has a higher ionization enthalpy: Be or B?
ⓐ. Beryllium
ⓑ. Boron
ⓒ. Both equal
ⓓ. Cannot be predicted
Correct Answer: Beryllium
Explanation: Be (1s²2s²) has a filled s-subshell, which is stable. Boron (1s²2s²2p¹) has an electron in a higher-energy 2p orbital, which is easier to remove. Hence, Be has higher IE than B despite having lower nuclear charge.
240. Which of the following factors does NOT affect ionization enthalpy?
ⓐ. Atomic radius
ⓑ. Nuclear charge
ⓒ. Shielding effect
ⓓ. Number of neutrons
Correct Answer: Number of neutrons
Explanation: Ionization enthalpy depends on how strongly the nucleus attracts outer electrons. This depends on effective nuclear charge, shielding, and atomic size. Neutrons do not influence electron attraction, so they have no direct effect on ionization enthalpy.
241. What is meant by electron gain enthalpy?
ⓐ. Energy required to remove an electron from a gaseous atom
ⓑ. Energy released when a gaseous atom gains an electron
ⓒ. Energy required to break a covalent bond
ⓓ. Energy released when a cation is formed in aqueous solution
Correct Answer: Energy released when a gaseous atom gains an electron
Explanation: Electron gain enthalpy (ΔegH) is the energy change when one mole of gaseous atoms gains one mole of electrons to form anions:
$\text{X(g) + e⁻ → X⁻(g)}$
A negative value indicates energy release. The more negative ΔegH, the greater the tendency to accept electrons.
242. Which element has the most negative electron gain enthalpy?
ⓐ. Oxygen
ⓑ. Fluorine
ⓒ. Chlorine
ⓓ. Neon
Correct Answer: Chlorine
Explanation: Although fluorine is more electronegative, chlorine has a more negative electron gain enthalpy (–349 kJ/mol). This is because in fluorine, the small 2p orbital has strong electron–electron repulsion, while chlorine’s larger 3p orbital better accommodates the extra electron.
243. What is the general trend of electron gain enthalpy across a period?
ⓐ. Becomes less negative
ⓑ. Becomes more negative
ⓒ. Remains constant
ⓓ. Changes irregularly without pattern
Correct Answer: Becomes more negative
Explanation: Across a period, effective nuclear charge increases, atomic size decreases, and atoms have a greater tendency to accept electrons. Therefore, electron gain enthalpy becomes more negative from left to right (except noble gases with positive values).
244. What is the general trend of electron gain enthalpy down a group?
ⓐ. Becomes more negative
ⓑ. Becomes less negative
ⓒ. Remains constant
ⓓ. Changes alternately
Correct Answer: Becomes less negative
Explanation: Down a group, atomic size increases and the added electron enters farther shells. The attraction from the nucleus decreases, making electron addition less favorable, so ΔegH becomes less negative.
245. Why do noble gases have positive electron gain enthalpies?
ⓐ. Because they are metals
ⓑ. Because they have incomplete shells
ⓒ. Because they have stable ns²np⁶ configurations
ⓓ. Because they have high nuclear charge
Correct Answer: Because they have stable ns²np⁶ configurations
Explanation: Noble gases already have a complete octet. Adding an electron would require entering a higher energy orbital, which is energetically unfavorable. Hence, their ΔegH values are positive.
246. Why does oxygen have less negative electron gain enthalpy than sulfur?
ⓐ. Because oxygen is more electronegative
ⓑ. Because oxygen is smaller and has greater electron–electron repulsion
ⓒ. Because sulfur has fewer protons
ⓓ. Because sulfur is a metal
Correct Answer: Because oxygen is smaller and has greater electron–electron repulsion
Explanation: Oxygen’s small size causes high repulsion when an extra electron enters the compact 2p orbital. In sulfur, the 3p orbital is larger, so repulsion is less and electron addition is easier, making sulfur’s ΔegH more negative.
247. Which group of elements has the least negative electron gain enthalpies?
ⓐ. Group 1 (alkali metals)
ⓑ. Group 2 (alkaline earth metals)
ⓒ. Group 17 (halogens)
ⓓ. Group 18 (noble gases)
Correct Answer: Group 18 (noble gases)
Explanation: Noble gases have stable octets and resist electron addition. Their electron gain enthalpies are positive (unfavorable), unlike halogens which have the most negative values due to their strong tendency to complete the octet.
248. Arrange the following in decreasing order of electron gain enthalpy: Cl, F, Br, I.
ⓐ. Cl > F > Br > I
ⓑ. F > Cl > Br > I
ⓒ. I > Br > Cl > F
ⓓ. Cl > Br > I > F
Correct Answer: Cl > F > Br > I
Explanation: Chlorine has the most negative ΔegH. Fluorine is less negative than chlorine due to high repulsion in small 2p orbitals. As size increases down the group (Br, I), attraction decreases, making ΔegH values less negative.
249. What happens to electron gain enthalpy across a period from sodium to chlorine?
ⓐ. It becomes less negative
ⓑ. It becomes more negative
ⓒ. It first decreases then increases
ⓓ. It stays constant
Correct Answer: It becomes more negative
Explanation: From Na (Group 1) to Cl (Group 17), effective nuclear charge increases and atomic radius decreases, so the tendency to accept electrons strengthens. Hence ΔegH becomes more negative, with Cl having the maximum value in that period.
250. Which factor does NOT directly affect electron gain enthalpy?
ⓐ. Atomic size
ⓑ. Nuclear charge
ⓒ. Electron configuration stability
ⓓ. Number of neutrons
Correct Answer: Number of neutrons
Explanation: Electron gain enthalpy depends on how strongly the nucleus attracts the incoming electron. Smaller atomic size, higher nuclear charge, and stable/unstable configurations matter. Neutrons do not affect electron–electron interactions directly, so they do not influence ΔegH.
251. What is meant by the first electron gain enthalpy of an element?
ⓐ. Energy released when one mole of electrons is removed from one mole of gaseous atoms
ⓑ. Energy required to remove an electron from a gaseous ion
ⓒ. Energy change when one mole of gaseous atoms gains one mole of electrons
ⓓ. Energy change when one mole of ions is dissolved in water
Correct Answer: Energy change when one mole of gaseous atoms gains one mole of electrons
Explanation: The first electron gain enthalpy (ΔegH₁) refers to the process:
$\text{X(g)} + e⁻ \rightarrow \text{X⁻(g)}$
This is usually exothermic (negative) because an electron is attracted to the positively charged nucleus, releasing energy.
252. What is meant by the second electron gain enthalpy?
ⓐ. Energy released when two electrons are gained simultaneously
ⓑ. Energy required to remove a second electron from an ion
ⓒ. Energy change when an electron is added to a negatively charged ion
ⓓ. Energy change when a gaseous atom gains two electrons at once
Correct Answer: Energy change when an electron is added to a negatively charged ion
Explanation: The second electron gain enthalpy (ΔegH₂) involves:
$\text{X⁻(g)} + e⁻ \rightarrow \text{X²⁻(g)}$
Since the electron is added to an already negatively charged ion, repulsion makes the process endothermic (positive).
253. Why is the first electron gain enthalpy of most non-metals negative?
ⓐ. Because they have small atomic size and high nuclear charge
ⓑ. Because they have full octets already
ⓒ. Because they have low electronegativity
ⓓ. Because they have many inner shells
Correct Answer: Because they have small atomic size and high nuclear charge
Explanation: Non-metals like halogens have strong nuclear attraction for electrons. Adding an electron stabilizes them by completing the octet, releasing energy, hence ΔegH₁ is negative (exothermic).
254. Why is the second electron gain enthalpy always positive?
ⓐ. Because nuclear charge decreases after first electron addition
ⓑ. Because an electron is being added to a negative ion, causing repulsion
ⓒ. Because protons repel electrons strongly
ⓓ. Because the atom becomes unstable after first electron gain
Correct Answer: Because an electron is being added to a negative ion, causing repulsion
Explanation: The incoming electron experiences repulsion from the already negative ion. Overcoming this repulsion requires energy input, so ΔegH₂ is always positive (endothermic), even for halogens.
255. Which of the following has the most negative first electron gain enthalpy?
ⓐ. Oxygen
ⓑ. Fluorine
ⓒ. Chlorine
ⓓ. Neon
Correct Answer: Chlorine
Explanation: Chlorine has the most negative ΔegH₁ (–349 kJ/mol). Fluorine is slightly less negative due to strong electron–electron repulsion in its small 2p orbital. Neon has positive ΔegH as it has a stable octet.
256. What is the first electron gain enthalpy of oxygen?
ⓐ. Positive
ⓑ. Negative
ⓒ. Zero
ⓓ. Very high positive
Correct Answer: Negative
Explanation: Oxygen has configuration 1s²2s²2p⁴. Addition of one electron gives O⁻ with a more stable 2p⁵ configuration, so energy is released. Thus ΔegH₁ is negative.
257. What is the second electron gain enthalpy of oxygen?
ⓐ. Positive
ⓑ. Negative
ⓒ. Same as the first
ⓓ. Zero
Correct Answer: Positive
Explanation: Adding a second electron to O⁻ forms O²⁻. Repulsion between the incoming electron and the already negative ion makes the process endothermic. Hence, ΔegH₂ is positive.
258. Which of the following best explains why halogens have highly negative first electron gain enthalpies?
ⓐ. They are metals
ⓑ. They have ns²np⁵ configuration, needing just one electron for octet
ⓒ. They have filled octets already
ⓓ. They have the largest atomic radii
Correct Answer: They have ns²np⁵ configuration, needing just one electron for octet
Explanation: Halogens require only one electron to complete the stable noble gas configuration. This strong drive makes ΔegH₁ highly negative, indicating high energy release when they accept an electron.
259. Which of the following has a positive first electron gain enthalpy?
ⓐ. Fluorine
ⓑ. Chlorine
ⓒ. Neon
ⓓ. Sulfur
Correct Answer: Neon
Explanation: Neon (1s²2s²2p⁶) already has a stable octet. Adding an electron requires entering a new orbital (3s), which is energetically unfavorable. Hence, ΔegH₁ is positive.
260. Which statement is correct about electron gain enthalpies?
ⓐ. ΔegH₁ is usually negative, ΔegH₂ is always positive
ⓑ. Both ΔegH₁ and ΔegH₂ are always negative
ⓒ. Both ΔegH₁ and ΔegH₂ are always positive
ⓓ. ΔegH₁ is positive only for halogens
Correct Answer: ΔegH₁ is usually negative, ΔegH₂ is always positive
Explanation: The first electron gain generally stabilizes an atom (energy release, negative ΔegH₁). The second electron gain involves adding to a negative ion, requiring energy input (positive ΔegH₂). This distinction is universal across elements.
261. Which of the following factors makes electron gain enthalpy more negative?
ⓐ. Large atomic radius
ⓑ. High effective nuclear charge
ⓒ. Strong electron–electron repulsion
ⓓ. Stable noble gas configuration
Correct Answer: High effective nuclear charge
Explanation: When effective nuclear charge is high, the nucleus attracts incoming electrons strongly, making the process energetically favorable. This leads to more negative ΔegH. Small atomic radius also enhances this effect.
262. Why is the electron gain enthalpy of fluorine less negative than chlorine?
ⓐ. Because chlorine is less electronegative
ⓑ. Because fluorine has strong inter-electron repulsion in its small 2p orbital
ⓒ. Because fluorine has fewer protons than chlorine
ⓓ. Because chlorine has a smaller radius than fluorine
Correct Answer: Because fluorine has strong inter-electron repulsion in its small 2p orbital
Explanation: Fluorine is very small, so incoming electrons experience strong repulsion in its compact 2p orbital. In chlorine, the 3p orbital is larger and accommodates the extra electron more easily, making ΔegH more negative.
263. Which group has the least negative electron gain enthalpy values?
ⓐ. Group 1 (alkali metals)
ⓑ. Group 2 (alkaline earth metals)
ⓒ. Group 17 (halogens)
ⓓ. Group 18 (noble gases)
Correct Answer: Group 18 (noble gases)
Explanation: Noble gases have stable ns²np⁶ configurations. Adding another electron requires entering a higher shell, which is energetically unfavorable. Thus, their electron gain enthalpies are positive or least negative.
264. Which of the following factors does NOT affect electron gain enthalpy?
ⓐ. Atomic size
ⓑ. Nuclear charge
ⓒ. Stable electronic configuration
ⓓ. Number of neutrons in the nucleus
Correct Answer: Number of neutrons in the nucleus
Explanation: Neutrons affect mass but not nuclear attraction for electrons. ΔegH mainly depends on atomic size, effective nuclear charge, and electron configuration stability.
265. Which element has the most negative electron gain enthalpy among the following?
ⓐ. Oxygen
ⓑ. Sulfur
ⓒ. Chlorine
ⓓ. Fluorine
Correct Answer: Chlorine
Explanation: Chlorine’s ΔegH is –349 kJ/mol, the most negative of all elements. Fluorine is slightly less negative due to electron repulsion, while oxygen and sulfur are less negative than halogens in general.
266. What is electronegativity?
ⓐ. The energy released when an atom gains an electron
ⓑ. The ability of an atom to attract a shared electron pair in a covalent bond
ⓒ. The energy required to remove an electron from an atom
ⓓ. The distance between nucleus and valence shell
Correct Answer: The ability of an atom to attract a shared electron pair in a covalent bond
Explanation: Electronegativity is not an energy but a relative tendency of atoms to attract bonding electrons. It depends on atomic size, nuclear charge, and bonding environment.
267. Which element has the highest electronegativity on the Pauling scale?
ⓐ. Oxygen
ⓑ. Nitrogen
ⓒ. Fluorine
ⓓ. Chlorine
Correct Answer: Fluorine
Explanation: Fluorine has the highest electronegativity (3.98 on the Pauling scale) due to its small size and high nuclear charge. This makes it the strongest attractor of bonding electrons.
268. What is the general trend of electronegativity across a period?
ⓐ. It decreases from left to right
ⓑ. It increases from left to right
ⓒ. It remains constant
ⓓ. It first decreases then increases
Correct Answer: It increases from left to right
Explanation: Across a period, atomic radius decreases and effective nuclear charge increases, so atoms attract bonding electrons more strongly. Thus, electronegativity rises left to right.
269. What is the general trend of electronegativity down a group?
ⓐ. It decreases
ⓑ. It increases
ⓒ. It remains constant
ⓓ. It changes irregularly
Correct Answer: It decreases
Explanation: Down a group, atomic size increases and valence electrons are farther from the nucleus. This weakens nuclear attraction on bonding electrons, so electronegativity decreases.
270. Which element has the lowest electronegativity?
ⓐ. Hydrogen
ⓑ. Fluorine
ⓒ. Cesium
ⓓ. Oxygen
Correct Answer: Cesium
Explanation: Cesium (Group 1, Period 6) has the lowest electronegativity (\~0.7). Its very large atomic radius and weak hold on valence electrons make it the least effective in attracting bonding electrons.
271. Who introduced the most widely used scale of electronegativity?
ⓐ. J.J. Thomson
ⓑ. Robert Mulliken
ⓒ. Linus Pauling
ⓓ. Henry Moseley
Correct Answer: Linus Pauling
Explanation: Linus Pauling introduced the Pauling scale of electronegativity in 1932. It is based on bond energy differences between covalent and heteronuclear bonds. Pauling’s scale remains the most commonly referenced electronegativity scale in chemistry.
272. What is the electronegativity value of fluorine on the Pauling scale?
ⓐ. 3.44
ⓑ. 3.98
ⓒ. 4.0
ⓓ. 2.55
Correct Answer: 3.98
Explanation: Fluorine has the highest electronegativity (≈3.98 on Pauling scale). This value reflects its very small atomic radius and strong nuclear attraction, making it the most powerful attractor of bonding electrons.
273. According to Pauling, electronegativity difference ($\Delta \chi$) between two atoms is related to which measurable property?
ⓐ. Ionization enthalpy
ⓑ. Bond length
ⓒ. Bond dissociation energy
ⓓ. Atomic radius
Correct Answer: Bond dissociation energy
Explanation: Pauling’s method is based on the observation that the bond energy of a heteronuclear bond (A–B) is higher than the average of homonuclear bonds (A–A, B–B). The extra stability is attributed to electronegativity difference, which he quantified mathematically.
274. According to Mulliken’s definition, electronegativity is:
ⓐ. The geometric mean of ionization enthalpy and electron gain enthalpy
ⓑ. The arithmetic mean of ionization enthalpy and electron gain enthalpy
ⓒ. The difference between ionization enthalpy and electron gain enthalpy
ⓓ. Twice the ionization enthalpy
Correct Answer: The arithmetic mean of ionization enthalpy and electron gain enthalpy
Explanation: Mulliken proposed that electronegativity is directly related to the tendency of an atom to gain or lose electrons. He defined it as:
$\chi = \frac{IE + EA}{2}$
where IE = ionization enthalpy, EA = electron affinity.
275. Which of the following is an advantage of the Mulliken scale over Pauling’s scale?
ⓐ. It is dimensionless
ⓑ. It is based on directly measurable atomic properties
ⓒ. It works only for metals
ⓓ. It uses bond energies of molecules
Correct Answer: It is based on directly measurable atomic properties
Explanation: Mulliken’s definition uses ionization enthalpy and electron gain enthalpy, which are directly measurable atomic properties. Pauling’s approach relies on bond energies, which may vary depending on molecular environment.
276. If the ionization enthalpy of chlorine is 1251 kJ/mol and electron gain enthalpy is –349 kJ/mol, its Mulliken electronegativity is approximately:
ⓐ. 901 kJ/mol
ⓑ. 1500 kJ/mol
ⓒ. 451 kJ/mol
ⓓ. 800 kJ/mol
Correct Answer: 901 kJ/mol
Explanation: Using Mulliken’s formula:
$\chi = \frac{IE + EA}{2} = \frac{1251 + 551}{2}$
(Note: electron gain enthalpy is taken as magnitude = 349 → 1251 + 349 = 1600/2 = 800 approx. Some references adjust sign conventions.) Depending on conventions, the Mulliken value lies between 800–900 kJ/mol.
277. Which statement correctly compares Pauling and Mulliken scales?
ⓐ. Pauling scale is based on bond energies; Mulliken on atomic properties
ⓑ. Mulliken scale is dimensionless; Pauling is not
ⓒ. Pauling scale is more accurate than Mulliken’s
ⓓ. Both give identical numerical values
Correct Answer: Pauling scale is based on bond energies; Mulliken on atomic properties
Explanation: Pauling’s scale uses bond dissociation energies as a measure of electronegativity difference, while Mulliken’s is derived from ionization enthalpy and electron gain enthalpy of atoms.
278. Which is a limitation of the Pauling scale?
ⓐ. It cannot compare metals and non-metals
ⓑ. It is not directly measurable for single atoms
ⓒ. It does not work for ionic compounds
ⓓ. It underestimates fluorine’s electronegativity
Correct Answer: It is not directly measurable for single atoms
Explanation: Pauling’s scale is indirect because it is based on bond energy data, which depends on the molecule considered. Hence, electronegativity values are relative, not absolute.
279. Mulliken’s electronegativity values are usually converted into Pauling’s scale by:
ⓐ. Subtracting electron gain enthalpy from ionization enthalpy
ⓑ. Taking geometric mean instead of arithmetic mean
ⓒ. Using an empirical proportionality constant
ⓓ. Dividing by bond length
Correct Answer: Using an empirical proportionality constant
Explanation: Mulliken values are dimensional (energy units), so to compare with Pauling’s dimensionless scale, an empirical constant is applied to normalize the values. This ensures consistency between scales.
280. Which of the following elements shows the highest Mulliken electronegativity?
ⓐ. Fluorine
ⓑ. Oxygen
ⓒ. Nitrogen
ⓓ. Neon
Correct Answer: Fluorine
Explanation: Fluorine, being the most electronegative element, has the highest Mulliken value as well. This is because it has a very high ionization enthalpy and strongly negative electron gain enthalpy, reflecting its great ability to attract electrons.
281. What is the general trend of electronegativity across a period?
ⓐ. Decreases from left to right
ⓑ. Increases from left to right
ⓒ. Remains constant
ⓓ. First decreases, then increases irregularly
Correct Answer: Increases from left to right
Explanation: Across a period, effective nuclear charge increases while atomic radius decreases. This strengthens the nucleus’ ability to attract shared electrons, making electronegativity rise from left to right, with halogens being highest in each period.
282. What is the general trend of electronegativity down a group?
ⓐ. It increases due to higher nuclear charge
ⓑ. It decreases due to increasing atomic radius and shielding
ⓒ. It remains constant
ⓓ. It fluctuates randomly
Correct Answer: It decreases due to increasing atomic radius and shielding
Explanation: Down a group, the number of shells increases, so valence electrons are farther from the nucleus. Shielding also increases, reducing effective nuclear attraction, hence electronegativity decreases.
283. Which element has the maximum electronegativity in the periodic table?
ⓐ. Oxygen
ⓑ. Chlorine
ⓒ. Fluorine
ⓓ. Neon
Correct Answer: Fluorine
Explanation: Fluorine (Pauling value ≈ 3.98) has the highest electronegativity due to its very small atomic radius and high effective nuclear charge, making it the strongest electron attractor.
284. Which of the following elements has the lowest electronegativity?
ⓐ. Sodium
ⓑ. Cesium
ⓒ. Barium
ⓓ. Magnesium
Correct Answer: Cesium
Explanation: Cesium (Group 1, Period 6) has a very large atomic radius and weak nuclear attraction on its outermost electron. Its electronegativity (\~0.7) is the lowest in the periodic table.
285. Why do noble gases have almost zero electronegativity values?
ⓐ. Because they have very small radii
ⓑ. Because they already have stable ns²np⁶ octets
ⓒ. Because their ionization enthalpies are low
ⓓ. Because they are metals
Correct Answer: Because they already have stable ns²np⁶ octets
Explanation: Noble gases have completely filled valence shells. They do not usually form bonds, hence their tendency to attract bonding electrons (electronegativity) is almost zero.
286. Which element has higher electronegativity, sulfur or chlorine?
ⓐ. Sulfur
ⓑ. Chlorine
ⓒ. Both equal
ⓓ. Depends on bond type
Correct Answer: Chlorine
Explanation: Both are in Period 3, but chlorine (Group 17) has higher nuclear charge and smaller atomic radius than sulfur (Group 16). Thus, chlorine’s electronegativity is greater (Cl = 3.16; S = 2.58 Pauling scale).
287. What is meant by valency of an element?
ⓐ. The number of isotopes an element has
ⓑ. The number of protons in the nucleus
ⓒ. The combining capacity of an element, determined by valence electrons
ⓓ. The energy required to remove valence electrons
Correct Answer: The combining capacity of an element, determined by valence electrons
Explanation: Valency is the ability of an atom to combine with others, generally equal to the number of electrons lost, gained, or shared to achieve a stable octet/duplet.
288. What is the valency of Group 1 elements?
ⓐ. 0
ⓑ. 1
ⓒ. 2
ⓓ. 7
Correct Answer: 1
Explanation: Group 1 (alkali metals) have configuration ns¹. They lose one electron to form M⁺ ions, so their valency is always 1. Example: Na → Na⁺.
289. What is the usual valency of oxygen?
ⓐ. 1
ⓑ. 2
ⓒ. 3
ⓓ. 4
Correct Answer: 2
Explanation: Oxygen (ns²np⁴) requires 2 more electrons to complete its octet. Hence, it typically shows valency 2, forming compounds like H₂O and MgO.
290. Which group of elements has variable valency most commonly?
ⓐ. s-block elements
ⓑ. p-block elements
ⓒ. d-block (transition) elements
ⓓ. Noble gases
Correct Answer: d-block (transition) elements
Explanation: Transition metals often show variable valency due to the close energy of ns and (n–1)d orbitals. For example, Fe shows +2 and +3 states; Cu shows +1 and +2. This variability is one of their key characteristics.
291. How is the valency of Group 1 elements related to their group number?
ⓐ. Valency = Group number – 1
ⓑ. Valency = Group number
ⓒ. Valency = Group number ÷ 2
ⓓ. Valency = 8 – Group number
Correct Answer: Valency = Group number
Explanation: Group 1 elements (ns¹) lose one electron to form M⁺ ions. Their group number is 1, and valency is also 1. Example: Na (Z = 11) → \[Ne]3s¹ → Na⁺ (valency 1).
292. What is the general valency of Group 2 elements?
ⓐ. 1
ⓑ. 2
ⓒ. 3
ⓓ. 4
Correct Answer: 2
Explanation: Group 2 (alkaline earth metals) have ns² configuration. They lose two electrons to form M²⁺ ions, so their valency is 2. Example: Mg → Mg²⁺.
293. For p-block elements, how is group number related to valency?
ⓐ. Group number = valence electrons = valency
ⓑ. Group number = 10 + valence electrons; valency varies as (8 – valence electrons) or valence electrons
ⓒ. Group number = atomic number ÷ 2
ⓓ. Group number = valency × 2
Correct Answer: Group number = 10 + valence electrons; valency varies as (8 – valence electrons) or valence electrons
Explanation: For example, Group 15 (ns²np³) has 5 valence electrons. Valency may be 3 (by sharing/losing 5–8 = 3 electrons) or 5 (by using all valence electrons).
294. What is the valency of Group 17 elements?
ⓐ. 1
ⓑ. 7
ⓒ. Either 1 or 7 depending on bonding
ⓓ. 0
Correct Answer: 1
Explanation: Group 17 (halogens) have ns²np⁵ configuration. They need one electron to complete their octet, so their common valency is –1. Example: Cl → Cl⁻.
295. Which group elements generally exhibit valency 0?
ⓐ. Group 2
ⓑ. Group 14
ⓒ. Group 18
ⓓ. Group 1
Correct Answer: Group 18
Explanation: Group 18 (noble gases) have ns²np⁶ stable configurations. They usually do not combine under normal conditions, hence valency is considered 0.
296. What is the valency of carbon in methane (CH₄)?
ⓐ. 2
ⓑ. 3
ⓒ. 4
ⓓ. 8
Correct Answer: 4
Explanation: Carbon belongs to Group 14 (ns²np²). It shares 4 electrons with 4 hydrogen atoms to complete its octet, hence valency = 4.
297. How does the valency of elements change across a period in the periodic table?
ⓐ. It decreases continuously
ⓑ. It increases from 1 to 4, then decreases back to 0
ⓒ. It increases only
ⓓ. It decreases only
Correct Answer: It increases from 1 to 4, then decreases back to 0
Explanation: Across a period, valency starts at 1 (Group 1), increases to 4 (Group 14), then decreases (Group 15 = 3, Group 16 = 2, Group 17 = 1, Group 18 = 0). This reflects the octet rule.
298. Which of the following pairs shows the same valency as their group number?
ⓐ. Na (Group 1), Mg (Group 2)
ⓑ. Cl (Group 17), O (Group 16)
ⓒ. Ne (Group 18), Ar (Group 18)
ⓓ. Al (Group 13), Si (Group 14)
Correct Answer: Na (Group 1), Mg (Group 2)
Explanation: Na has valency 1 (Group 1), Mg has valency 2 (Group 2). For p-block elements, valency often differs from group number due to covalency and octet stability.
299. Which of the following elements has a valency equal to 3?
ⓐ. Boron
ⓑ. Beryllium
ⓒ. Oxygen
ⓓ. Fluorine
Correct Answer: Boron
Explanation: Boron (Group 13) has ns²np¹ configuration. It has 3 valence electrons, which it shares to form covalent compounds like BCl₃. Hence, its valency is 3.
300. What is the maximum valency shown by sulfur?
ⓐ. 2
ⓑ. 4
ⓒ. 6
ⓓ. 8
Correct Answer: 6
Explanation: Sulfur (Group 16) normally shows valency 2 (as in H₂S), but in compounds like SF₆, it uses its d-orbitals to expand the octet and shows a maximum valency of 6.
In Class 11 Chemistry (NCERT/CBSE syllabus), the chapter Classification of Elements and Periodicity in Properties plays a vital role in linking basic atomic structure to chemical reactivity. This chapter explains why elements of the same group show similar properties and how periodic variation helps in predicting the behavior of elements. Topics include trends in atomic radius, ionic radius, ionization energy, electron gain enthalpy, and electronegativity across periods and down groups. Special focus is given to valency, oxidation states, variable valency examples, diagonal relationship, anomalous behavior due to small size, high electronegativity, and absence of d-orbitals. The full set has 350 MCQs with correct answers, divided into 4 parts. Here in Part 3, you will get another 100 MCQs with solutions, specially useful for board exam preparation and competitive tests like JEE and NEET.
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