Which Rule Is Violated In The Following Orbital Diagram: Complete Guide

Which Rule Is Violated in the Following Orbital Diagram?
The short version is: you’ve probably broken Hund’s rule, but let’s dig into why that matters.

Ever stared at an orbital diagram that just didn’t look right and thought, “Did I just break a rule I never even knew existed?” You’re not alone. Most of us have drawn a sketch of carbon’s 2p orbitals, only to later realize we’ve stacked two electrons into the same box and left another empty. The moment that happens, the whole picture feels off—like a puzzle piece that refuses to click.

In practice, those little mistakes can snowball. A single mis‑placed electron means an incorrect electron configuration, which then ripples through everything: predicted magnetic properties, bonding behavior, even the color of a compound. So, let’s unpack the rules that keep orbital diagrams honest, spot the common offender, and learn how to avoid the slip‑up that trips up even seasoned chemistry students.

What Is an Orbital Diagram?

Think of an orbital diagram as a quick‑draw cartoon of where electrons live inside an atom. Each box represents an orbital—s, p, d, or f—and each little arrow is an electron with a specific spin. The arrows point up or down to indicate “spin‑up” or “spin‑down,” a convention that helps us keep track of the Pauli exclusion principle.

When you’re sketching, you start with the lowest‑energy orbitals (1s, then 2s, then 2p, and so on) and fill them according to a handful of rules. So the diagram itself is nothing more than a visual shorthand for the atom’s electron configuration. It’s the kind of thing you see on a whiteboard during a freshman chemistry lecture, but it’s also the foundation for everything from molecular orbital theory to predicting reactivity.

The Core Rules That Govern Filling

1. Aufbau Principle – “Build up” from the lowest‑energy orbital to the highest. In shorthand, you follow the n + l rule (lower n + l gets filled first; if equal, lower n wins).
2. Pauli Exclusion Principle – No two electrons in the same atom can have identical sets of four quantum numbers. In the diagram, that means a single orbital can hold at most two electrons, and they must have opposite spins.
3. Hund’s Rule – For orbitals of equal energy (like the three 2p orbitals), electrons fill singly first, with parallel spins, before any pairing occurs.

If you break any of these, the diagram no longer reflects a physically possible ground‑state arrangement. The question “which rule is violated?” usually points to the one that’s easiest to overlook: Hund’s rule Small thing, real impact..

Why It Matters / Why People Care

You might wonder why a tiny arrow pointing the wrong way matters at all. The answer is simple: electron arrangement dictates an atom’s chemistry. Here’s a quick rundown of the real‑world stakes Still holds up..

• Magnetism – Unpaired electrons give rise to paramagnetism. If you mistakenly pair electrons that should be unpaired, you’ll predict a non‑magnetic substance when it’s actually magnetic. Think oxygen: O₂ is paramagnetic because of two unpaired electrons in its π* orbitals. Miss those, and you’d call air “non‑magnetic”—which is technically true, but you’d miss the nuance that makes O₂ interesting.

• Bonding Patterns – Hybridization models (sp³, sp², etc.) assume a certain number of unpaired electrons ready to form sigma bonds. A diagram that violates Hund’s rule can suggest the wrong hybridization, leading you to predict the wrong molecular geometry.

• Spectroscopy – Transition energies depend on the exact electron distribution. An incorrect diagram throws off your calculations for UV‑Vis or IR spectra, meaning you could misinterpret experimental data.

• Chemical Reactivity – Unpaired electrons are the “reactive sites.” If you think a carbon atom has a paired set in its 2p orbitals, you might underestimate its tendency to form radicals or engage in addition reactions Practical, not theoretical..

Bottom line: the rule you break determines whether you’ll end up with a plausible chemical story or a fictional one.

How It Works (or How to Do It)

Let’s walk through building a correct orbital diagram step by step, then see where the most common mistake slips in.

1. List the Electron Count

First, know how many electrons you’re dealing with. For a neutral atom, that’s simply the atomic number. Example: nitrogen (Z = 7) has seven electrons Less friction, more output..

2. Order the Orbitals by Energy

Use the n + l rule:

When n + l ties, the lower n fills first (so 2p before 3s).

3. Apply the Pauli Exclusion Principle

Each box (orbital) can hold up to two arrows, and they must be opposite (↑↓). No more than two arrows per box, period.

4. Follow Hund’s Rule for Degenerate Orbitals

Here’s where most people trip. For a set of orbitals with the same energy (the three 2p boxes, the five 3d boxes, etc.):

• Put one electron in each box first.
• All those single electrons must have the same spin (usually drawn as ↑).
• Only after every box has one electron do you start pairing (↑↓).

5. Fill Until You Run Out of Electrons

Continue the process until you’ve placed all electrons. Let’s do nitrogen as a concrete example Easy to understand, harder to ignore..

Step‑by‑step nitrogen (7e⁻):

1. 1s: ↑↓ (2 electrons)
2. 2s: ↑↓ (2 electrons) – total 4
3. 2p: three boxes → place one ↑ in each (3 electrons) – total 7

Result:

1s  ↑↓
2s  ↑↓
2p  ↑   ↑   ↑

All three 2p electrons are unpaired and have parallel spins—perfect Hund compliance.

6. Spotting a Violation

Now imagine you drew nitrogen’s 2p as:

2p  ↑↓   ↑   (empty)

You paired two electrons in the first 2p box and left another box empty. That’s a classic Hund’s rule violation. The Pauli principle is still satisfied (no more than two per box), but you ignored the “maximum multiplicity” part of Hund’s rule.

Common Mistakes / What Most People Get Wrong

Mistake #1 – Pairing Too Early

As shown, the urge to “fill the first box completely” is strong, especially when you’re used to the 1s and 2s orbitals, which are single boxes. Degenerate sets are different; they want you to spread the electrons out first Worth knowing..

Mistake #2 – Forgetting the “Same Spin” Part

Even if you do put one electron in each box, you might accidentally draw opposite spins (↑, ↓, ↑). That still breaks Hund’s rule because the rule specifically calls for parallel spins to maximize total spin multiplicity.

Mistake #3 – Mis‑ordering Energy Levels

Sometimes students think 4s fills before 3d (which is true) but then reverse it later, putting a 3d electron before the 4s is full. That’s an Aufbau violation, not a Hund one, but it often shows up in the same diagrams and confuses the analysis.

Not obvious, but once you see it — you'll see it everywhere Worth keeping that in mind..

Mistake #4 – Ignoring Subshell Capacity

A common slip is treating a p subshell as if it only has one orbital. Remember: p = 3 orbitals, d = 5, f = 7. Over‑crowding a single box is a Pauli violation, while under‑using the available boxes is a Hund violation That alone is useful..

Not obvious, but once you see it — you'll see it everywhere.

Mistake #5 – Applying the Rule to Excited States

Hund’s rule strictly applies to ground‑state configurations. On top of that, if you’re deliberately drawing an excited state (say, promoting an electron from 2p to 3s), you can break Hund’s rule—but you must label the diagram as excited. Forgetting to do that leads reviewers to think you made a mistake.

Practical Tips / What Actually Works

1. Draw a quick “box map” first. Sketch three empty boxes for any p subshell before you start placing arrows. It forces you to think about distribution.

2. Use the same arrow direction for all unpaired electrons. Pick ↑ as your default spin for the first pass; only switch to ↓ when you’re pairing.

3. Count as you go. After each step, tally the electrons you’ve placed. If the total doesn’t match the atomic number, you’ve missed something.

4. Double‑check with electron configuration notation. Write the shorthand (e.g., 1s² 2s² 2p³) and see if the diagram matches. If the superscripts don’t line up with the number of arrows, you’ve slipped.

5. Practice with odd‑electron atoms first. Elements like nitrogen, phosphorus, and chlorine are perfect for spotting Hund violations because they have partially filled p subshells.

6. Use a colored pencil or marker. One color for ↑, another for ↓. Visual contrast makes accidental spin mismatches obvious.

7. When in doubt, ask “Are any orbitals of the same energy partially filled?” If yes, make sure each has at most one electron before any pairing Worth keeping that in mind..

FAQ

Q1: Can a diagram violate more than one rule at once?
A: Absolutely. A common combo is pairing electrons early and putting a 3d electron before the 4s is full. That breaks both Hund’s and Aufbau principles. Spotting the first violation usually leads you to the second The details matter here..

Q2: Does Hund’s rule apply to d and f subshells?
A: Yes. Any set of degenerate orbitals—p (3), d (5), f (7)—follows the same rule: fill each singly with parallel spins before pairing.

Q3: What if I’m drawing an ion?
A: The same rules apply; just adjust the electron count. Here's one way to look at it: O²⁻ has ten electrons, so its 2p subshell ends up fully paired (↑↓ ↑↓ ↑↓), which does violate Hund’s rule—but that’s okay because the ion’s ground state has a filled subshell, not a partially filled one.

Q4: How do I know when a diagram is for an excited state?
A: Excited‑state diagrams are usually labeled (e.g., “*” or “excited”). If you’re just sketching a ground‑state configuration, any Hund violation signals an error That's the part that actually makes a difference..

Q5: Why do textbooks sometimes show Hund‑violating diagrams?
A: Rarely, but it can happen in examples meant to illustrate what not to do or to discuss excited configurations. The caption should make that clear—if it doesn’t, it’s probably a typo.

So, you’ve seen the pattern: most “which rule is violated?” questions point straight at Hund’s rule. The mistake is easy to make, hard to spot, and surprisingly impactful. Now, next time you pull out a pen and start filling boxes, remember the three‑step checklist: spread, align spins, then pair. Your orbital diagrams will stay honest, and the chemistry that follows will make a lot more sense. Happy sketching!