The P-Block Elements NDA Questions

NDA Chemistry

The P-Block Elements MCQ Questions


13.
PCl₃ and PCl₅ are two important halides. Why does PCl₅ act ONLY as an oxidising agent (not as reducing agent)?
A.
In PCl₅, phosphorus is in its HIGHEST oxidation state (+5) — it can only DECREASE its oxidation state (get reduced) by acting as an oxidising agent; it cannot increase further (no higher state available)
B.
PCl₅ contains Cl in −1 state which prevents reduction
C.
PCl₅ has weaker P-Cl bonds than PCl₃ making reduction impossible
D.
PCl₅ is a gas and gases cannot be reducing agents
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ANSWER :
A. In PCl₅, phosphorus is in its HIGHEST oxidation state (+5) — it can only DECREASE its oxidation state (get reduced) by acting as an oxidising agent; it cannot increase further (no higher state available)
14.
The structure of PCl₅ in solid state is:
A.
Ionic [PCl₄]⁺[PCl₆]⁻ — in solid state PCl₅ is not molecular but exists as tetrachlorophosphonium hexachlorophosphate
B.
Octahedral structure with lone pair
C.
Trigonal bipyramidal molecular structure (same as gas phase)
D.
Square pyramidal molecular structure
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ANSWER :
A. Ionic [PCl₄]⁺[PCl₆]⁻ — in solid state PCl₅ is not molecular but exists as tetrachlorophosphonium hexachlorophosphate
15.
The OXOACIDS OF PHOSPHORUS include H₃PO₂, H₃PO₃ and H₃PO₄. Their basicity is 1, 2 and 3 respectively because:
A.
Basicity is determined by the number of P=O double bonds
B.
Basicity equals the total number of oxygen atoms in the molecule
C.
Basicity = number of ionisable O-H (P-OH) groups, NOT the total number of OH groups; P-H bonds are not ionisable
D.
Basicity equals the number of P-H bonds in the molecule
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ANSWER :
C. Basicity = number of ionisable O-H (P-OH) groups, NOT the total number of OH groups; P-H bonds are not ionisable
16.
The general valence shell electronic configuration of Group 16 elements is:
A.
ns²np⁵
B.
ns²np³
C.
ns²np²
D.
ns²np⁴ — two electrons short of the noble gas configuration
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ANSWER :
D. ns²np⁴ — two electrons short of the noble gas configuration
17.
Oxygen shows ONLY −2 oxidation state (except in OF₂). Why can oxygen NOT show positive oxidation states in most compounds?
A.
Oxygen has no d-orbitals so it cannot expand its valence shell to positive states
B.
Oxygen is the MOST ELECTRONEGATIVE element after fluorine (O = 3.5, F = 4.0) — it always attracts electrons toward itself in any bond, giving it negative character; only F (more electronegative) can force O to be positive
C.
Oxygen always exists as O₂ and cannot form other bonding types
D.
Oxygen's ionisation energy is too low for it to lose electrons
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ANSWER :
B. Oxygen is the MOST ELECTRONEGATIVE element after fluorine (O = 3.5, F = 4.0) — it always attracts electrons toward itself in any bond, giving it negative character; only F (more electronegative) can force O to be positive
18.
Molecular OXYGEN (O₂) is PARAMAGNETIC despite having an even number of electrons. This is explained by:
A.
The two O atoms exert different electronegativity making one electron spin differently
B.
Valence bond theory explains O₂ paramagnetism by a double bond
C.
O₂ paramagnetism is due to the presence of ionic O²⁻ species
D.
Molecular orbital theory: O₂ has two unpaired electrons in the degenerate π*2p antibonding orbitals — each π*2p orbital gets one electron (Hund's rule in MO theory)
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ANSWER :
D. Molecular orbital theory: O₂ has two unpaired electrons in the degenerate π*2p antibonding orbitals — each π*2p orbital gets one electron (Hund's rule in MO theory)