Getting the Electron Dots Right
Nitrogen sits in group 15 of the periodic table, which means it has five valence electrons. The Lewis structure is built from that single fact. You place the symbol N in the center and distribute those five dots around it, one per side before any pairing happens. Left, right, top, bottom, then back to the first side. That gives you three single unpaired dots and one pair. This arrangement matters because those three unpaired electrons are what nitrogen uses when it forms bonds. The diagram itself is straightforward. Nitrogen's symbol with five dots: a lone pair on one side and a single dot on each of the other three sides. In practice I draw it clockwise starting from the top, but the order doesn't affect the chemistry. What matters is that you can clearly see the three available bonding sites. I ran into a real issue once when I was grading student work and noticed nearly everyone was pairing up the dots too early. They'd put two on the top, two on the right, and one on the left, completely skipping the bottom position. That gave the wrong visual impression of the bonding geometry. The workaround is simple but worth enforcing: force yourself to place exactly one dot on each of the four sides before you go back and add the second dot to any side. It takes three seconds longer and eliminates the error entirely.
One thing beginners consistently miss is that the Lewis dot diagram for nitrogen alone is not the same thing as the diagram for a nitrogen compound. When nitrogen forms ammonia, for instance, it shares those three unpaired electrons with three hydrogen atoms, and the lone pair stays on the nitrogen. Students sometimes draw all five dots still floating around after bonding, which makes no sense because the shared pairs become bonding pairs, not nonbonding ones. The lone pair is the only thing that stays strictly as dots on the central atom after bond formation. Another nuance that doesn't get enough attention is formal charge. If you draw nitrogen with just the five valence dots and no bonds, the formal charge is zero. That's correct for the isolated atom. But when nitrogen forms three bonds and retains one lone pair, like in NH3, the formal charge calculation is still zero: five valence electrons minus two nonbonding electrons minus three bonding electrons (the six shared electrons divided by two). It comes out clean. However, if a student draws nitrogen with four bonds and no lone pair, like in the ammonium ion NH4+, the formal charge becomes positive one. That structure is valid but only when there's an accompanying negative charge somewhere else to balance it. Drawing that configuration without the charge present is a common mistake that signals a misunderstanding of the whole point of the exercise. The biggest limitation of Lewis dot diagrams is that they don't actually show geometry. They suggest it by implying where bonds can form, but the three-dimensional arrangement of NH3 is trigonal pyramidal, not flat, and the diagram gives no hint of that. VSEPR theory handles the geometry piece, and you should learn both together rather than expecting the dot diagram to do all the work.
A download link isn't really necessary for something this small, but if you need a reference image, most general chemistry textbooks have a clear version in their appendix, and LibreTexts hosts a public domain set you can print out. The Chemistry LibreTexts nitrogen page has it alongside related elements for comparison, which is useful if you're trying to see the pattern across the whole group. When you move past the isolated atom to actual molecules, the same five-electron rule governs everything nitrogen does. Triple bonds like in N2 use all three unpaired electrons plus the lone pair participates in resonance descriptions. The diagram stays the same starting point regardless of what molecule you're dealing with. Understanding that starting point is what separates people who can draw these structures from people who can look at them and actually predict reactivity.
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