To determine if a substance acts as an oxidising or reducing agent, we need to examine the oxidation state of its central atom. An oxidising agent itself gets reduced (its oxidation state decreases), while a reducing agent itself gets oxidised (its oxidation state increases). The maximum possible oxidation state for an element limits its ability to be further oxidised.
Correct Option: 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).
The question asks about the structure of PCl₅ in the solid state. It's important to differentiate between the gas phase and solid phase structures of compounds, especially those that can undergo disproportionation or ionization.
A) Ionic [PCl₄]⁺[PCl₆]⁻ — in solid state PCl₅ is not molecular but exists as tetrachlorophosphonium hexachlorophosphate
This option correctly describes the ionic nature and the specific ions formed by PCl₅ in the solid state. The tetrachlorophosphonium cation \([\text{PCl}_4]^+\) is tetrahedral, and the hexachlorophosphate anion \([\text{PCl}_6]^-\) is octahedral.
The basicity of an oxoacid is determined by the number of ionizable hydrogen atoms it possesses. For phosphorus oxoacids, only the hydrogen atoms attached to oxygen (P-OH groups) are acidic and can be ionized. Hydrogen atoms directly bonded to phosphorus (P-H bonds) are generally not acidic.
Correct Option: C) Basicity = number of ionisable O-H (P-OH) groups, NOT the total number of OH groups; P-H bonds are not ionisable
To determine the general valence shell electronic configuration of Group 16 elements, we need to understand the periodic table trends and how electron configurations are built. Group 16 elements are also known as Chalcogens.
Correct Option: D) ns²np⁴ — two electrons short of the noble gas configuration
This option correctly states the general valence shell electronic configuration for Group 16 elements and accurately describes that they are two electrons short of achieving a stable noble gas configuration (\(ns^2 np^6\)).
The question asks why oxygen typically does not exhibit positive oxidation states, except in compounds like OF₂. This relates to oxygen's fundamental chemical properties, particularly its electronegativity.
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. This statement accurately explains why oxygen typically has negative oxidation states and why OF₂ is an exception. Its high electronegativity means it pulls electrons towards itself in most bonds, leading to negative oxidation states. Only fluorine, being more electronegative, can cause oxygen to have a positive oxidation state.
The paramagnetism of a molecule, especially one with an even number of electrons like O₂, is a classic example where Valence Bond Theory (VBT) fails and Molecular Orbital Theory (MOT) provides the correct explanation. Paramagnetism arises from the presence of unpaired electrons.
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). This statement accurately describes why O₂ is paramagnetic. Molecular Orbital Theory correctly predicts the presence of two unpaired electrons in the antibonding \(\pi^{*}_{2p}\) orbitals, which are responsible for its paramagnetic behavior.