Working with organic and inorganic compounds in the lab
Separating Organic And Inorganic Compounds by Solubility
I started learning this stuff because my university lab assigned us a mixed-sample practical exam. We had to identify and separate compounds without being handed a step-by-step manual. That was my introduction to the fact that textbooks describe ideal conditions, and real samples are never ideal. Organic compounds are built around carbon frameworks. They include things like hydrocarbons, alcohols, carboxylic acids, esters, and amides. Inorganic compounds cover everything else. Salts, metal oxides, mineral acids, cyanides, carbonates. The carbon boundary is the dividing line, though even carbon has inorganic exceptions like CO2 and the carbonate family. The key thing nobody emphasizes enough is that solubility isn't just about polarity. It's about lattice energy versus hydration energy for inorganics, and intermolecular forces for organics. If you just remember "like dissolves like," you'll run into trouble fast. Sodium chloride dissolves in water because the hydration energy of Na+ and Cl- overcomes its lattice energy. Acetone dissolves in water because it can accept and donate hydrogen bonds. But sodium acetate, which is an organic salt, has different solubility behavior than acetic acid, even though they share the same carbon backbone. The ionized form interacts completely differently.
Here's a practical separation workflow I use. Take your mixed sample. Add distilled water. Stir and let it sit. What dissolved went into the aqueous layer. What didn't is either an insoluble organic or an insoluble inorganic. Filter through qualitative filter paper. The solid on the paper goes one way. The filtrate goes the other. Then process each stream separately. For the aqueous filtrate, evaporate carefully. What crystallizes out could be an inorganic salt like NaCl, CuSO4, or CaCl2. It could also be something like sodium acetate, which is technically organic but behaves like an inorganic salt in water. You test by igniting a small amount. Inorganics don't char. Most organics char or burn. That's your quick split. For the solid residue, you need an organic solvent. Diethyl ether is standard but dangerous if you're working near open flames. Ethyl acetate is safer and works well for most neutral organics. Wash the residue with the solvent. The organics dissolve. Evaporate the ether under reduced pressure. What's left is your organic fraction. Recrystallize from ethanol if you need purity.
I ran into a problem last year that cost me three hours. A student sample claimed to contain both calcium carbonate and naphthalene. Classic separation exercise. I added water, filtered, evaporated the filtrate, and got nothing crystalline. I rechecked the original mixture. The calcium carbonate wasn't dissolving. It was finely powdered and had caked into a filter cake that water couldn't penetrate. The carbonate stayed trapped in the solid phase along with the naphthalene. When I tested the "aqueous" fraction later with silver nitrate, I got zero precipitation, confirming no ions had leached out. The workaround was simple but not obvious if you haven't seen it. I took the wet filter cake and added dilute HCl directly to it before filtering again. The acid dissolved the carbonate, releasing CO2 bubbles so I could see the reaction happening. Then I filtered. The filtrate now contained Ca2+ and Cl- ions. Confirming with ammonium oxalate gave me the white calcium oxalate precipitate. Meanwhile, the naphthalene stayed behind in the solid residue and came off cleanly with ether wash. Here's a counter-intuitive point that trips people up. Silver nitrate is commonly used to test for chloride ions in inorganic chemistry. But if your sample contains any organochlorine compound, like chloroform or a chlorinated solvent residue, adding silver nitrate won't give you a precipitate immediately. The chlorine is covalently bound, not ionic. You'd need to do a sodium fusion test first to break those C-Cl bonds and convert them to chloride ions. I've seen students misidentify a chlorobenzene contamination as pure water because they skipped that step.
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Another common error involves distinguishing between organic and inorganic carbonates. Both release CO2 with acid. The difference is structural. Inorganic carbonates like CaCO3 and Na2CO3 are ionic lattices. Organic carbonates like ethylene carbonate are covalent molecular structures. They behave completely differently in solvents and under heat. Heating Na2CO3 doesn't decompose it at normal lab temperatures. Heating an organic carbonate might decompose it at 150C. Thermal behavior is a better differentiator than acid reactivity here. For identification beyond simple solubility tests, you'll want infrared spectroscopy. O-H stretches around 3300 per cm indicate alcohols or carboxylic acids. C=O stretches around 1700 per cm point to ketones, aldehydes, or acids. Ionic compounds don't show these peaks the same way. Their spectra are dominated by lattice vibrations in the far infrared region, below 400 per cm, which standard IR instruments don't always capture well. Melting point determination works for pure organic compounds. Impure samples melt lower and over a broader range. Inorganic salts typically have much higher melting points. NaCl melts at 801C. Sugar decomposes before melting. If you're heating a unknown solid and it turns black before reaching 200C, it's almost certainly organic. Most inorganics survive well beyond that.
The limitation of this whole approach is that some compounds exist in gray areas. Metal organic frameworks blur the line completely. Urea contains carbon but behaves like an inorganic salt in many reactions. Ammonium cyanate was the compound Wöhler accidentally synthesized from inorganic precursors in 1828, and that single experiment collapsed the old vitalism theory. You can't always neatly classify something before you've tested it. If you need a reference for standard solubility tables, the CRC Handbook of Chemistry and Physics has the most reliable data. The Merck Index is better for organic compounds specifically. Neither is free, but university libraries carry both. Online alternatives like PubChem are adequate for quick checks but sometimes list incomplete data for obscure derivatives. For a hands-on workflow, I keep a simple spreadsheet tracking sample origin, appearance, solubility results in water and ether, pH response, and ignition test outcome. It takes about ten minutes to set up per sample and saves you from repeating steps when you get ambiguous results. The data becomes useful later when you're trying to figure out why two samples that looked identical behaved differently.
