Which one of the following constitutes a group of the isoelectronic species?
(a) C2(2-), O2(-), CO,

Q: Why do we subtract electrons when there is a positive charge?

Electrons carry a negative charge. If an atom or molecule loses a negatively charged electron, the overall balance shifts, leaving it with a net positive charge. Therefore, a $+1$ charge means one electron has been removed from the total count.

Q: Do we consider only valence electrons or total electrons for isoelectronic species?

You must count the **total** electrons. The term 'isoelectronic' strictly refers to the absolute total number of electrons present in the entire species, not just the outermost valence shell.

Q: Are isoelectronic species always isostructural (having the same shape)?

Not always, but very often! For diatomic molecules (like the ones in this question), being isoelectronic usually means they have similar molecular orbital diagrams and the same bond order. However, for larger polyatomic molecules, they might be isoelectronic but have different geometries depending on the central atom.

Q: What about option (a)? Why is it incorrect?

Let's check option (a): $\text{C}_2^{2-}$ has 14 electrons, $\text{CO}$ has 14, and $\text{NO}$ has 15 ($7+8$). Because $\text{NO}$ has 15 electrons, the group is not isoelectronic.

Q: Is there a quick way to count electrons during an exam to save time?

Yes! Memorize the atomic numbers of the first 10-15 elements perfectly. When you see common pairs like $\text{CO}$ ($6+8=14$) and $\text{N}_2$ ($7+7=14$), you will instantly recognize them as isoelectronic without having to write down the math.

Question

Which one of the following constitutes a group of the isoelectronic species?

Options:

(A) $	ext{C}_2^{2-}, 	ext{O}_2^-, 	ext{CO}, 	ext{NO}$

(B) $	ext{NO}^+, 	ext{C}_2^{2-}, 	ext{CN}^-, 	ext{N}_2$

(C) $	ext{CN}^-, 	ext{N}_2, 	ext{O}_2^{2-}, 	ext{C}_2^{2-}$

(D) $	ext{N}_2, 	ext{O}_2^-, 	ext{NO}^+, 	ext{CO}$

Sigma Percentile · Accepted Answer

### The Secret Society of 14 Electrons: Unveiling Isoelectronic Species

When you encounter the term **isoelectronic**, the secret to solving the problem is hidden right in the word itself. The prefix *iso-* comes from Greek, meaning "equal" or "same," and *electronic* refers to electrons. Therefore, isoelectronic species are simply atoms, molecules, or ions that possess the exact same total number of electrons.

In competitive exams like JEE and NEET, identifying isoelectronic groups is a classic, high-yield question. It tests your fundamental ability to recall atomic numbers and correctly account for ionic charges. Let's break down the math behind it.

### The Master Equation for Electron Counting

To find the total number of electrons in any given species, you don't need complex quantum mechanics. You just need basic arithmetic:

$	ext{Total } e^- = \sum Z \pm 	ext{charge}$

Here, $Z$ represents the atomic number (which equals the number of electrons in a neutral atom) for each element in the molecule.

**Crucial Trap:** Remember that electrons are negatively charged! If an ion has a positive charge (a cation), it means it has *lost* electrons, so you must subtract that number. Conversely, if an ion has a negative charge (an anion), it has *gained* electrons, so you must add that number.

### Analyzing the Winning Group

Let's put our formula to the test by evaluating the species given in Option (b):

**1. The Nitrosonium Ion ($	ext{NO}^+$)**
Nitrogen ($Z=7$) brings 7 electrons, and Oxygen ($Z=8$) brings 8 electrons. The $+1$ charge indicates the loss of one electron.
$	ext{Total } e^- = 7 + 8 - 1 = 14$

**2. The Acetylide Ion ($	ext{C}_2^{2-}$)**
Each Carbon ($Z=6$) brings 6 electrons. The $-2$ charge indicates a gain of two extra electrons.
$	ext{Total } e^- = 6 + 6 + 2 = 14$

**3. The Cyanide Ion ($	ext{CN}^-$)**
Carbon ($Z=6$) brings 6 electrons, and Nitrogen ($Z=7$) brings 7 electrons. The $-1$ charge adds one more electron.
$	ext{Total } e^- = 6 + 7 + 1 = 14$

**4. The Nitrogen Molecule ($	ext{N}_2$)**
Two Nitrogen atoms ($Z=7$) bring 7 electrons each. It is a neutral molecule, so there is no charge to adjust.
$	ext{Total } e^- = 7 + 7 = 14$

### The Verdict and The Way Forward

As we can clearly see, every single species in Option (b) contains exactly **14 electrons**. They form a perfect isoelectronic series!

Why does this matter beyond just answering the question? In the realm of **Molecular Orbital Theory**, isoelectronic diatomic species often exhibit strikingly similar chemical properties. Because they have the same number of electrons filling their molecular orbitals, they frequently share the same bond order and magnetic behavior. For instance, all the 14-electron species we just calculated have a robust **bond order of 3** (think of the triple bond in $	ext{N} \equiv 	ext{N}$ or $	ext{C} \equiv 	ext{O}$) and are diamagnetic. Recognizing these patterns will give you a massive speed advantage in your exams!

Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: Which one of the following constitutes a group of the isoelectronic species?

Select Answer:

Visualized Solution

$Concept of Isoelectronic Species$

$Electron Counting Formula$

$Electrons in NO^{+}$

$Electrons in C_{2}^{2 -}$

$Electrons in CN^{-}$

$Electrons in N_{2}$

$Conclusion$

$Pro Tip: Molecular Orbital Theory$

The Sigma Insight: Molecular Orbital Theory

The Secret Society of 14 Electrons

The Master Equation for Electron Counting

Analyzing the Winning Group

The Verdict and The Way Forward

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