The MOT Diagram of N2: Why Most Students Get the Energy Levels Wrong

Ever stared at a chemistry textbook and wondered why the mot diagram of n2 looks just a little "off" compared to oxygen? You're not alone. Honestly, it’s one of those things that trips up even the brightest students because it breaks the pattern we’re taught for the rest of the periodic table.

Nitrogen is a weird one.

While O2 and F2 follow a very straightforward energy progression, nitrogen decides to throw a wrench in the gears with something called s-p mixing. If you've ever felt like the energy levels in dinitrogen were a bit of a chaotic mess, there’s actually a very logical (and cool) reason for it. Let's get into the weeds of how these orbitals actually stack up.

The Weird Reality of S-P Mixing

Most people think the $\sigma_{2p}$ orbital should always be lower in energy than the $\pi_{2p}$ orbitals. Why? Because sigma bonds involve head-on overlap, which is typically stronger and more stable than the side-on overlap of pi bonds.

But nitrogen doesn't play by those rules.

In the mot diagram of n2, the $\pi_{2p}$ orbitals actually sit below the $\sigma_{2p}$ orbital. This happens because the 2s and 2p atomic orbitals in nitrogen are surprisingly close in energy. Since they’re so close, they "talk" to each other. This interaction—the s-p mixing—pushes the $\sigma_{2s}$ and $\sigma^*_{2s}$ levels down while shoving the $\sigma_{2p}$ level way up.

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Basically, the $\sigma_{2p}$ gets a massive "energy penalty," leaving the $\pi$ orbitals as the more stable option for those electrons.

It’s a bit like a traffic detour. Because the 2s and 2p are basically neighbors, they interfere with each other's construction. By the time you get to oxygen, the gap between 2s and 2p is so huge that they don't interact anymore, and the "normal" order returns. But for N2? The detour is the only way to go.

Counting the 14 Electrons

To build the diagram, you've gotta track 14 total electrons.

• 7 from the first Nitrogen.
• 7 from the second Nitrogen.

If we're just looking at the valence shell (which is where the magic happens), we're dealing with 10 electrons.

1. First, the 1s orbitals fill up—but we usually ignore those since they’re core electrons.
2. Then the $\sigma_{2s}$ takes two.
3. The $\sigma^*_{2s}$ (antibonding) takes another two.
4. Next come the $\pi_{2p}$ orbitals. There are two of them, so they hold four electrons total.
5. Finally, the $\sigma_{2p}$ takes the last two.

Wait. Notice something?

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All 10 valence electrons are tucked away in bonding or low-energy antibonding orbitals, and every single one of them is paired up. This is exactly why nitrogen is diamagnetic. If you put a tube of liquid nitrogen between two magnets, it won't stick. Compare that to oxygen, which has unpaired electrons and is paramagnetic (it'll actually hang out between magnet poles like a tiny liquid bridge).

Calculating the Triple Bond

Chemistry teachers love "Bond Order." It's just a fancy way of asking, "How many bonds are holding these atoms together?"

The math is simple:
$Bond Order = \frac{(Bonding Electrons - Antibonding Electrons)}{2}$

In the mot diagram of n2, we have:

• 2 electrons in $\sigma_{2s}$ (bonding)
• 4 electrons in $\pi_{2p}$ (bonding)
• 2 electrons in $\sigma_{2p}$ (bonding)
• 2 electrons in $\sigma^*_{2s}$ (antibonding)

That gives us $8 - 2 = 6$ bonding-advantaged electrons. Divide by two, and you get 3.

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Three bonds. A triple bond. That's why nitrogen is so incredibly hard to break apart and why it makes up 78% of our atmosphere without reacting with every single thing it touches. It’s basically the "Fort Knox" of molecules.

The Difference Between Theory and Reality

While we draw these neat little lines in the mot diagram of n2, the reality is a bit more fluid. Research in recent years, including work cited in Dipak Kumar Mandal’s Molecular Orbital Theory and Frontier Orbitals (2025/2026 editions), suggests that the degree of s-p mixing can vary slightly depending on the environment and the ionic state of the molecule.

For example, if you look at $N_2^+$, you’ve removed an electron from the $\sigma_{2p}$ orbital (the HOMO, or Highest Occupied Molecular Orbital). Suddenly, the bond order drops to 2.5, and the molecule becomes paramagnetic. The bond gets longer and weaker.

It’s weird to think that removing just one tiny electron can turn one of the strongest bonds in nature into something significantly more fragile.

Key Takeaways for Your Next Exam

If you're trying to memorize this for a test or a project, don't just memorize the "shape." Understand the "why."

• The Swap: Always remember that $\pi$ comes before $\sigma$ in N2. This is the biggest mistake students make.
• Diamagnetism: No lonely electrons. Everyone has a partner.
• Bond Order: It's 3. Always 3 for neutral N2.
• S-P Mixing: This is the "secret sauce" that explains the energy shift. It only happens significantly for elements before Oxygen in the second row.

To master this, try drawing the diagram from scratch without looking at a reference. Start with the atomic orbitals (2s and 2p) on the sides, then build the molecular orbitals in the middle. Remember the "2-1-2-1" pattern for the p-block: two $\pi$, one $\sigma$, then the antibonding versions ($\pi^$ then $\sigma^$). If you can explain the s-p mixing to a friend, you've actually understood it, not just memorized it.

Next Steps for Mastery
Check your understanding by comparing the N2 diagram directly against the O2 diagram. Specifically, look at the placement of the $\sigma_{2p}$ level. Once you see the "crossover" where the $\sigma$ level drops below the $\pi$ level as you move from Nitrogen to Oxygen, the concept of effective nuclear charge and orbital energy will finally click. For more complex systems, look into how this theory applies to heteronuclear molecules like CO (Carbon Monoxide), which is isoelectronic with N2 but has a skewed diagram because Oxygen is more electronegative than Carbon.