Q.Which of the following species is not expected to be a ligand?

The key idea is that a ligand must have at least one lone pair of electrons available for donation to a metal centre.

1. Check each species for a lone pair:
   • $NO$: Has an unpaired electron and a lone pair on nitrogen — can act as a ligand.
   • $NH_4^{+}$: All four hydrogens are bonded; nitrogen has no lone pair — cannot donate electrons.
   • $NH_2CH_2CH_2NH_2$ (ethylenediamine): Each nitrogen has a lone pair — a classic bidentate ligand. …

To decide whether a species can be a ligand, you need to recall the fundamental definition: a ligand is an ion or molecule that donates a pair of electrons to a central metal atom to form a coordinate bond. That means the species must possess at least one lone pair of electrons that is available for donation.

Let’s examine each option with this idea in mind.

1. Option (i): $NO$ (Nitric oxide)

   Nitric oxide has an unpaired electron and a lone pair on nitrogen. It is a well-known ligand (often called nitrosyl) that can donate its lone pair to a metal. In fact, it forms stable complexes like $[Fe(NO)(H_2O)_5]^{2+}$. So it is expected to be a ligand.

2. Option (ii): $NH_4^{+}$ (Ammonium ion)

   In $NH_4^{+}$, nitrogen is bonded to four hydrogen atoms via four sigma bonds. All four valence electrons of nitrogen are used in these bonds — there is no lone pair left. Without a lone pair, it cannot donate electrons to a metal.

   Watch out

   A common mistake is to confuse $NH_4^{+}$ with $NH_3$. Ammonia ($NH_3$) has a lone pair and is a classic ligand, but the ammonium ion has lost that lone pair by forming a fourth bond to $H^{+}$. They are chemically different species.

   Therefore, $NH_4^{+}$ cannot act as a ligand.

3. Option (iii): $NH_2CH_2CH_2NH_2$ (Ethylenediamine) …

Concept: Ligands and Donor Atoms

A ligand is a molecule or ion that donates a lone pair of electrons to a central metal atom/ion to form a coordinate bond. For a species to act as a ligand, it must have at least one lone pair of electrons available for donation.

Method: Lone Pair Availability Check

Step 1: Identify the donor atom(s) in each species.

Step 2: Check if the donor atom has a lone pair that is free to donate.

Step 3: If no lone pair exists, the species cannot act as a ligand.

Applying the method:

Common Mistakes: Identifying Ligands in Coordination Chemistry

The Concept in Brief

A ligand is a molecule or ion that donates a lone pair of electrons to a central metal atom/ion to form a coordinate bond. For a species to act as a ligand, it must have at least one lone pair available for donation.

Mistake #1: Assuming All Neutral Molecules Are Ligands

The error: Students think that because $NO$ and $CO$ are neutral, they automatically qualify as ligands. They forget to check if the lone pair is actually available for donation.

Why it’s wrong:

• $CO$ has a lone pair on carbon and is a classic ligand (carbonyl).
• $NO$ also has a lone pair on nitrogen and acts as a ligand (nitrosyl). Both are valid ligands — the mistake is ruling them out without checking.

How to avoid:

• Always check for lone pairs on the atom that would donate.
• Neutral molecules can be ligands, but only if they have a free lone pair.

Mistake #2: Confusing $NH_4^+$ with $NH_3$

The error: Students see nitrogen and hydrogen and assume $NH_4^+$ can donate like $NH_3$.

Why it’s wrong:

• $NH_3$ has a lone pair on nitrogen → acts as a ligand.
• $NH_4^+$ has no lone pair — all four bonds are used, and the positive charge means no extra electrons to donate.

How to avoid:

• Count total valence electrons and bond pairs.
• If every valence electron is used in bonding (and there’s a positive charge), there is no lone pair left.
• Key check: $NH_4^+$ is not a ligand; $NH_3$ is.

Mistake #3: Overlooking the Role of Charge

The error: Students ignore the charge on a species when deciding if it can donate electrons.

Why it’s wrong:

• A positive charge often means the species has lost electrons — making lone pairs less likely.
• $NH_4^+$ is a classic example: it’s electron-deficient, not electron-rich.

How to avoid:

• For cations, ask: Does this ion still have a lone pair?
• For anions (like $Cl^-$, $CN^-$), they usually have extra lone pairs → good ligands.
• For neutral species, check the Lewis structure carefully.

Mistake #4: Thinking “Bigger Molecule = Always a Ligand” …