The thing nobody tells you about Lewis structures
You draw dots around element symbols to show valence electrons. That's the whole premise. It works fine for carbon, nitrogen, oxygen, fluorine. You get through a dozen problems and think you've got it. Then you hit a transition metal or an ion and the whole system starts looking like a guessing game. I spent three years in introductory chemistry tutoring before I stopped treating electron dot diagrams as sacred and started understanding where they actually break down. Here's the method. Find the element's group number. For main group elements, that number equals the valence electrons. Group 1 gets one dot. Group 14 gets four. Group 17 gets seven. Group 18 gets eight. Place dots around the symbol, one per side before pairing, starting at the top and going clockwise. Single dot, then another on each side until all four sides have one, then start pairing them up. I know that sounds straightforward but most people skip the pairing order and just draw random dots. They draw two pairs and two singles for nitrogen instead of three singles and one pair. The difference matters when you're predicting bonding geometry later. Chemists don't care about your diagram's aesthetic. They care whether you represented the unpaired electrons correctly because those unpaired electrons are what form bonds.
Here's a practical example. Sulfur is in group 16. Six valence electrons. You put one dot on each of the four sides, then pair two of them. You end up with two lone pairs and two single electrons ready to bond. That's why sulfur forms two bonds in hydrogen sulfide but can expand to six in sulfur hexafluoride. The dots tell you the story before you even write a formula.
Where it gets complicated
Transition metals are the first place this system quietly fails. Iron sits in group 8 but its common oxidation states are plus 2 and plus 3. Does iron have eight valence electrons or six or five? The periodic table group number doesn't answer that cleanly because d-orbital electrons complicate everything. I once had a student draw an iron electron dot diagram with eight dots and then watch him get confused when iron chloride only uses two of those electrons in bonding. The truth is most introductory courses don't even ask you to draw transition metal dot diagrams for a reason. The model loses predictive power past the p-block. Another edge case that trips people up is polyatomic ions. Take the ammonium ion. Nitrogen has five valence electrons, each hydrogen contributes one, and you subtract one for the positive charge. That gives you eight electrons total, four N-H bonds, zero lone pairs on the nitrogen. Drawing that as a dot diagram requires you to understand coordinate covalent bonding, which most textbooks introduce three chapters after they start teaching Lewis structures. You'll draw it wrong if you just do the arithmetic without understanding where the electrons actually came from. Resonance structures are another place where the single-diagram approach falls apart. Ozone cannot be accurately represented by one Periodic Table Electron Dot Diagram. You need two resonance forms because the actual molecule has delocalized electrons. Drawing one static diagram and calling it done is technically incorrect, though most high school classes accept it. If you're in college chemistry, your professor will mark it wrong.
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The actual limitation you should know about
Electron dot diagrams are a visual shortcut, not a rigorous model. They work for predicting simple molecular geometries with VSEPR theory. They fail when you need to explain magnetic properties, spectral data, or reactivity patterns in organometallic chemistry. I've seen students rely on dot diagrams so heavily that they couldn't explain why oxygen is paramagnetic when the diagram shows all paired electrons. Molecular orbital theory is what actually answers that question, and it has nothing to do with dot placement. For routine homework and general chemistry exams, the method is adequate. You should know it cold. But don't confuse it with an accurate picture of how electrons actually behave. It's a bookkeeping system, not a quantum mechanical description. The best use of this skill is as a stepping stone to understanding electronegativity differences, bond polarity, and formal charge calculations. Skip that foundation and you'll struggle when organic chemistry asks you to track electron movement in reaction mechanisms.
