Acetanilide melts at approximately 114 °C (237 °F), with most reference sources listing the range as 113–115 °C. Precise calorimetric measurements put the onset of melting right around 387–388 K, which converts to roughly 114–115 °C. That narrow, well-defined melting range is the reason acetanilide shows up so often in chemistry laboratories, where it serves as both a teaching tool and a calibration standard for melting-point apparatus.
Why a Range Instead of a Single Number
You will sometimes see acetanilide’s melting point reported as a single value and other times as a span of a couple of degrees. Both are correct in their own way. A perfectly pure crystal has a single thermodynamic melting point, and careful thermal analysis of acetanilide places that point near 387.5 K, or about 114.4 °C.1Thermochimica Acta. Thermochemistry of fusion, solution and hydrogen bonding in benzamide, N-methylbenzamide, and acetanilide In practice, though, a small sample observed in a capillary tube does not all transition from solid to liquid at exactly the same instant. You first notice the crystals starting to soften or become translucent, then the last speck of solid disappears a degree or two later. That is why lab reports typically quote a melting range rather than a single temperature. For high-purity acetanilide, that range is tight: 113–115 °C.2International Journal of Scientific Research in Chemical Sciences. Identification of an Unknown Organic Compound by Determination of its Melting Point
If your sample is less pure, the range widens and the onset drops. This is the classic “melting-point depression” effect. Even a small amount of a foreign substance disrupts the orderly crystal packing, making the solid easier to break apart. A chemistry student who measures acetanilide melting at, say, 105–110 °C can be fairly sure the sample has a significant impurity. The sharper and closer to 114 °C the range, the cleaner the compound.
Acetanilide as a Calibration and Teaching Standard
Walk into almost any undergraduate organic chemistry lab and you will find acetanilide in the stockroom. There are a few reasons it has earned that role. Its melting point sits comfortably in the middle of the range that benchtop melting-point devices are designed to measure, typically around 50–300 °C. It is inexpensive, stable at room temperature, and not particularly hazardous compared with many other organic solids. And because the literature value is so well established and reproducible, it makes an ideal reference for checking whether a melting-point apparatus is reading accurately. If your instrument says pure acetanilide melts at 118 °C, you know the thermometer or heating rate needs adjusting before you trust it with unknowns.
Acetanilide also has a convenient property for recrystallization exercises. It dissolves readily in hot water and much less so in cold water, which makes it a textbook candidate for the classic “dissolve in hot solvent, filter, cool, and collect crystals” procedure. Students can then confirm that their recrystallized product is pure by checking whether the melting range is sharp and close to the literature value. The whole sequence, from impure starting material to pure crystals verified by melting point, teaches fundamental lab technique in a single afternoon.
The Mixed Melting-Point Test
One of the more elegant uses of melting-point data is the mixed melting-point method for identifying unknown compounds. The idea is simple: if you have a mystery white powder that melts around 113–115 °C, you look up which known compounds share that range. You then physically grind your unknown together with one of those candidates and measure the melting point of the mixture. If the two substances are the same compound, the mixture melts at the same temperature the pure compound would. If they are different, the mixture’s melting point drops noticeably, because each acts as an impurity in the other.
A published lab exercise illustrates this neatly. Researchers had an unknown substance melting at 112–116 °C and narrowed it down to either acetanilide (113–115 °C) or 4-nitrophenol (113–114 °C). When they mixed the unknown with acetanilide and measured the result, the melting range held steady at 113–116 °C. No depression meant the unknown was acetanilide.2International Journal of Scientific Research in Chemical Sciences. Identification of an Unknown Organic Compound by Determination of its Melting Point Had the unknown been 4-nitrophenol, the mixed sample would have melted several degrees lower. This technique predates modern spectroscopy and is still taught because it reinforces how melting-point data encode identity and purity in a single measurement.
The Energy It Takes to Melt Acetanilide
Melting point is a temperature, but it only tells half the story. The other half is how much energy the solid absorbs at that temperature before it becomes liquid. For acetanilide, this heat of fusion has been measured many times and consistently comes in around 21–22 kilojoules per mole. A 2023 calorimetric study found a value of about 21.3 kJ/mol at the melting temperature, in good agreement with a weighted average from earlier literature of roughly 21.9 kJ/mol.1Thermochimica Acta. Thermochemistry of fusion, solution and hydrogen bonding in benzamide, N-methylbenzamide, and acetanilide
Why does this matter outside a thermodynamics classroom? The heat of fusion determines how useful a substance is for storing thermal energy. Materials that absorb a lot of heat while melting, called phase-change materials, can act like thermal batteries. You pump heat in, the solid melts and soaks up energy without getting hotter, and later, when it solidifies again, it releases that stored energy. The melting point dictates the temperature at which this energy exchange happens, and the heat of fusion dictates how much energy you get per cycle.
Acetanilide in Thermal Energy Storage
Acetanilide’s melting point near 114 °C puts it in a temperature range useful for certain solar cooking and low-grade heat-storage applications. Researchers have tested it through hundreds of melt-freeze cycles to see how well it holds up. A study on indoor solar cooking put acetanilide through 500 thermal cycles and tracked whether its melting temperature, melting range, and latent heat changed over that span.3Energy Conversion and Management. Fast thermal cycling of acetanilide and magnesium chloride hexahydrate for indoor solar cooking The fact that the compound was tested at all for this purpose tells you something about its appeal: it is cheap, nontoxic at the quantities involved, and melts at a useful cooking-adjacent temperature.
Compared with common alternatives like paraffin waxes (which tend to melt at lower temperatures, often 40–70 °C) or salt hydrates (which can be corrosive and sometimes degrade unpredictably), acetanilide occupies a niche for medium-temperature heat storage. Its main limitation is that 114 °C is still well below the temperatures needed for many industrial processes, and its heat of fusion per gram is moderate rather than exceptional. But for applications where you want to capture and slowly release heat in the 100–120 °C window, it is a candidate worth considering.
How Acetanilide Became a Chemical Workhorse
Acetanilide’s place in chemistry goes back to the late nineteenth century, when it was one of the first synthetic drugs to reach widespread use. It is a simple chemical derivative of aniline, which itself was a byproduct of coal-tar processing. The early pharmaceutical industry seized on aniline derivatives because they were abundant, easy to modify, and showed interesting biological activity. Acetanilide was marketed as a fever reducer and painkiller under the trade name Antifebrin starting in 1886.4PubMed. Early drug discovery and the rise of pharmaceutical chemistry
It worked, but it also caused a dangerous side effect: it could convert hemoglobin in the blood to methemoglobin, which cannot carry oxygen effectively. Patients taking acetanilide sometimes developed a bluish discoloration of the skin, and in serious cases the oxygen-carrying capacity of their blood dropped to harmful levels. Chemists eventually traced the painkilling activity to a metabolite, acetaminophen (paracetamol), which the body produces when it processes acetanilide. Acetaminophen turned out to give the therapeutic benefit without the same degree of blood toxicity, and it gradually replaced acetanilide in medicine over the first half of the twentieth century.
So while acetanilide itself is no longer used as a drug, it is, in a sense, the grandparent of one of the most widely consumed medications on the planet. Its molecular structure is closely related to acetaminophen’s: the difference is a single hydroxyl group on the benzene ring that acetaminophen has and acetanilide lacks. That small structural change improves safety enormously, a fact that helped establish the principle that tiny molecular modifications can have outsized effects on how a drug behaves in the body.
Why Impurities Depress the Melting Point
The melting-point depression effect mentioned earlier deserves a bit more explanation because it comes up constantly in practical lab work. When a pure substance crystallizes, its molecules pack into a regular repeating lattice. Every molecule is surrounded by identical neighbors, and the attractive forces between them are uniform throughout. To melt the crystal, you have to supply enough thermal energy to overcome those attractions all at once, and that happens at a well-defined temperature.
Now imagine a few molecules of something else scattered through that lattice. They do not fit perfectly, and the forces between them and their neighbors differ from the pure-crystal interactions. The lattice is slightly destabilized at those spots, so it takes less energy to start breaking apart. The result is that melting begins at a lower temperature. And because the disruption is uneven across the crystal, different regions start melting at slightly different temperatures, widening the observed melting range. The more impurity, the more the temperature drops and the broader the range becomes.
For acetanilide, this effect is easy to observe. A student who accidentally contaminates their recrystallized product with even a few percent of another substance will see the melting range drop below 113 °C and stretch over five or more degrees. Experienced chemists use this as a quick quality check: if the melting range is sharp and hits the expected value, the sample is clean enough for the next step. If not, another round of purification is needed.
Factors That Can Shift Your Reading in the Lab
Even with a pure sample, a few practical factors can make your measured melting point drift from the textbook value. The most common culprit is the heating rate. If you ramp the temperature too fast, the thermometer reading lags behind the actual sample temperature, and you overshoot. Most lab protocols recommend a slow heating rate of about one degree per minute near the expected melting point to avoid this. Some modern melting-point instruments automate this by ramping quickly through the initial range and then slowing down as the target zone approaches.
Sample preparation also matters. If the crystals are packed too loosely in the capillary tube, the heat transfer through the sample is uneven, and some parts melt before others in a way that has nothing to do with purity. Conversely, grinding crystals into a fine powder and tapping the tube gently to pack them gives more consistent results. How much sample you use can also affect things: too much, and the center stays cooler than the edge; too little, and the onset of melting is hard to see.
Atmospheric pressure has a tiny theoretical effect on melting point, but for most organic solids, including acetanilide, the shift is negligible at the elevations where labs normally operate. You would need to be conducting the experiment at a wildly different pressure, like deep underwater or in a near-vacuum, to notice a meaningful change.
Acetanilide Compared With Related Compounds
Acetanilide belongs to a family of aniline-derived amides and related compounds, and comparing melting points within this family reveals how molecular structure influences crystal packing. Acetaminophen, with that extra hydroxyl group on the ring, melts at about 169 °C, substantially higher than acetanilide’s 114 °C. The hydroxyl group allows for additional hydrogen bonding in the crystal, holding the lattice together more tightly and demanding more thermal energy to break apart.
Phenacetin, another historical painkiller closely related to both, melts at around 134 °C. It has an ethoxy group on the ring instead of a hydroxyl, and that gives intermediate hydrogen-bonding ability. All three compounds were studied together in thermodynamic work comparing their sublimation properties, and the pattern is consistent: more capacity for hydrogen bonding in the solid state generally means a higher melting point.1Thermochimica Acta. Thermochemistry of fusion, solution and hydrogen bonding in benzamide, N-methylbenzamide, and acetanilide
Benzamide, another simple amide with a similar molecular weight to acetanilide, melts at about 130 °C. Its amide group can form stronger intermolecular hydrogen bonds than acetanilide’s because the nitrogen in benzamide carries two hydrogens rather than one. These comparisons underscore a useful rule of thumb: when you see two structurally similar organic molecules with very different melting points, the one with the higher melting point almost always has stronger or more numerous intermolecular attractions holding its crystal together.
Polymorphism and Whether It Affects Acetanilide
Some organic compounds can crystallize in more than one crystal structure, a phenomenon called polymorphism. Different polymorphs of the same compound can have noticeably different melting points because the molecules pack differently and the lattice energies differ. Chocolate, for instance, has multiple polymorphic forms of cocoa butter, and getting the wrong one is what causes that white bloom on poorly tempered bars.
Acetanilide, as it turns out, is not particularly prone to polymorphism under normal conditions. Its most common crystal form is orthorhombic, and this is the form you will almost always encounter in a laboratory setting. There has been some research into high-pressure crystal behavior of acetanilide, but for standard benchtop work at atmospheric pressure, you can treat its melting point as belonging to a single, stable crystal form. This is part of what makes it so reliable as a calibration standard: you do not have to worry about accidentally growing a different polymorph that melts at a different temperature, a headache that plagues calibration with some other compounds.
If you ever do suspect you have an unusual crystal form of acetanilide, the mixed melting-point test would still flag it. Mixing a standard sample with the suspect sample would show depression if the crystal structures differed, just as it would for two different chemical compounds. In practice, though, this situation is rare enough that most chemists never encounter it with acetanilide.