Redox Behaviour, Periodic Trends and Extraction Reactions

Cu+ (a d10 configuration) is not stable in aqueous solution: it spontaneously disproportionates, 2Cu+ → Cu2+ + Cu0 (metallic copper), because the hydration energy released by forming the smaller, more highly charged Cu2+ ion more than compensates for the energy cost of the disproportionation, making the reaction thermodynamically favourable in water. This is why copper(I) salts are essentially never encountered as stable species in aqueous solution, only in the solid state or in non-aqueous/complexed environments that suppress this disproportionation.

The Fe3+/Fe2+ reduction potential is favourable enough to oxidize iodide ion to iodine: Fe3+ + I- → Fe2+ + 1/2 I2 (written per iodide;

This is a standard qualitative test/reaction pairing a transition-metal redox couple (Fe3+/Fe2+) against a halide redox couple (I2/I-), and it works specifically because iodide is among the more easily oxidized halides — the same logic would not necessarily hold for fluoride or chloride, which are much harder to oxidize.

Paramagnetism tracks the number of unpaired d-electrons, which does NOT increase monotonically across a simple d1-d2-d3... series once you reach the half-filled d5 configuration and beyond: V2+ (d3, 3 unpaired) < Cr2+ (d4, 4 unpaired) < Mn2+ (d5, 5 unpaired, the maximum possible for these five ions) — but Fe2+ (d6, only 4 unpaired again, since the 6th electron must pair up in an already-occupied orbital) actually drops back BELOW Mn2+, not above it. So V2+ < Cr2+ < Mn2+ < Fe2+ is NOT a correct order for paramagnetism — it breaks exactly at the Mn2+-to-Fe2+ step. Ionic size across a row of same-charge transition-metal ions generally DECREASES from left to right (increasing nuclear charge pulls the d-electrons in more tightly, only partially shielded by other d-electrons), so Ni2+ < Co2+ < Fe2+ < Mn2+ correctly describes ionic size increasing as you move left (Mn2+ is the largest of this set). The number of distinct oxidation states a first-row transition metal can access tends to increase from the early transition metals toward the middle of the row (peaking around manganese) before decreasing again, so Sc < Ti < Cr < Mn is a correct increasing order for the number of accessible oxidation states.

Metallic bonding strength across the first-row transition metals depends heavily on how many d-electrons are available to participate in delocalised metallic bonding. Zinc has a full, non-bonding d10 configuration: its d-electrons are all paired and effectively core-like, contributing essentially nothing extra to metallic bonding beyond its two 4s electrons. Metals in the middle of the row (like Cr, with a half-filled d5 configuration ideal for maximum d-orbital overlap and bonding) have far more d-electrons available for strong metallic bonding, and so have much higher melting and boiling points. This is why zinc has an unusually LOW boiling point compared to its first-row transition-metal neighbours, despite being a d-block element.

In copper metallurgy, heating copper(I) oxide (Cu2O, Cu at +1) together with copper(I) sulfide (Cu2S, Cu at +1) drives a comproportionation-type reaction where sulfur is oxidized (from Cu2S, S at -2, up to +4 in SO2) while copper is reduced from +1 down to 0 (metallic copper): Cu2S + 2Cu2O → 6Cu + SO2. This reaction, sometimes called self-reduction or auto-reduction, is a key step in the pyrometallurgical extraction of copper, where the sulfide ore is partially roasted to oxide first, and the leftover sulfide then reduces that oxide directly on further heating, with no separate external reducing agent needed.