If you've ever spent more than five minutes in an organic chemistry lab, you know that distinct, slightly sweet smell. It’s chloroform. Most people think of it as the stuff movie villains use to knock people out, but for scientists, it's a workhorse solvent. One of the most critical things to understand about it isn't its toxicity—it's the boiling point for chloroform.
The number is $61.15^\circ\text{C}$ (or about $142^\circ\text{F}$).
Why does that matter? Well, if you’re trying to recover a sample from a solution or run a Soxhlet extraction, knowing exactly when your solvent turns to vapor is the difference between a successful experiment and a ruined week. It’s a relatively low boiling point compared to water, which makes it incredibly easy to "strip" off using a rotary evaporator. But that ease comes with a hidden set of risks that most textbooks gloss over.
Understanding the Boiling Point for Chloroform under Pressure
Nature isn't a vacuum, except when it is. The $61.15^\circ\text{C}$ figure is what we call the "normal boiling point," measured at standard atmospheric pressure (1 atm or 760 mmHg). But move your lab to Denver, or put your flask under a vacuum, and that number evaporates. Literally.
When you decrease the pressure, the boiling point for chloroform drops significantly. This is the whole principle behind vacuum distillation. By pulling a vacuum of just 100 mmHg, you can get chloroform to boil at roughly $20^\circ\text{C}$. This is great for heat-sensitive compounds that might decompose if you cranked the heat up to sixty degrees. However, it also means that if your lab is too warm and your bottle isn't sealed tight, you’re losing expensive (and regulated) solvent to the atmosphere every second.
The Clausius-Clapeyron Relationship
We use the Clausius-Clapeyron equation to predict these changes. It looks like this:
$$\ln\left(\frac{P_2}{P_1}\right) = -\frac{\Delta H_{vap}}{R} \left(\frac{1}{T_2} - \frac{1}{T_1}\right)$$
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In this formula, $\Delta H_{vap}$ is the enthalpy of vaporization. For chloroform, that’s about $29.24 \text{ kJ/mol}$. If you’re a student, you've probably cursed this equation during an exam. If you’re a professional, you probably just use a nomograph or a digital pressure controller. But the math stays the same. Pressure and temperature are locked in a tight dance.
Why Chloroform Behaves Differently Than Water
Water is weird because of hydrogen bonding. Chloroform ($CHCl_3$) doesn't have that same internal "stickiness." While it is a polar molecule, its intermolecular forces are primarily dipole-dipole interactions and London dispersion forces.
Because these bonds are weaker than the hydrogen bonds in water, it takes much less energy—less heat—to kick a chloroform molecule out of the liquid phase and into the gas phase. That’s why the boiling point for chloroform is nearly $40^\circ\text{C}$ lower than water's.
It’s also heavy. Really heavy.
Chloroform has a density of about $1.49 \text{ g/cm}^3$. When you boil it, those heavy vapors don't just float away like steam; they tend to sink and "pool" on the floor or in the bottom of a fume hood. This creates a massive inhalation risk if your ventilation isn't perfect.
The Phosgene Problem: Heating Chloroform Safely
Here is the thing nobody mentions enough: chloroform is chemically unstable when exposed to light and oxygen. Over time, it degrades into phosgene gas ($COCl_2$). Phosgene was used as a chemical weapon in WWI. It’s nasty stuff.
When you reach the boiling point for chloroform, you are adding energy to the system. If your chloroform is old or hasn't been stabilized with ethanol or amylene, you could be accelerating the production of phosgene.
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- Always check for stabilizers. Most lab-grade chloroform contains $0.5%$ to $1%$ ethanol to neutralize any phosgene that forms.
- Never distill to dryness. If you boil off all the liquid, the residues left behind can be unstable or even explosive.
- Store it in amber glass. Light is the enemy. It provides the activation energy for degradation.
Azeotropes: When Chloroform Gets Stuck
Sometimes, the boiling point for chloroform isn't actually $61.15^\circ\text{C}$. If you mix it with other liquids, they can form an azeotrope. This is a mixture of liquids that maintains a constant boiling point and composition throughout distillation.
For instance, chloroform and methanol form a positive azeotrope. The mixture boils at $53.5^\circ\text{C}$, which is lower than the boiling point of either pure chloroform or pure methanol ($64.7^\circ\text{C}$).
This is a nightmare for purification. If you have a mixture of these two, you can't just "boil off" the chloroform and leave the methanol behind. They are chemically "stuck" together at that specific temperature. You’d need to use a technique like extractive distillation or add a third component to break the azeotrope.
Common Chloroform Azeotropes
- Chloroform + Water: Boils at $56.1^\circ\text{C}$ (contains about $97%$ chloroform).
- Chloroform + Ethanol: Boils at $59.4^\circ\text{C}$ (contains about $93%$ chloroform).
- Chloroform + Acetone: This is a rare "negative azeotrope." The mixture boils at $64.7^\circ\text{C}$, which is higher than both components. This happens because of strong attractive forces between the hydrogen in chloroform and the carbonyl oxygen in acetone.
Practical Applications in Modern Industry
Despite the rise of greener solvents like dichloromethane (DCM), the specific boiling point for chloroform makes it irreplaceable in certain niches.
In the pharmaceutical industry, it’s used in the extraction of alkaloids from plant materials. Because it boils at $61^\circ\text{C}$, you can maintain a warm environment that increases solubility without destroying the delicate organic molecules you're trying to extract.
It's also a precursor to Teflon. The process involves reacting chloroform with hydrogen fluoride to create chlorodifluoromethane ($R-22$), which is then heated—way past its boiling point—to create tetrafluoroethylene.
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Real-World Safety: What to do at the Boiling Point
Honestly, the biggest mistake people make is underestimating the vapor pressure. Just because you aren't at the boiling point for chloroform doesn't mean it isn't dangerous. At room temperature ($20^\circ\text{C}$), chloroform has a vapor pressure of about 160 mmHg. That is extremely high.
If you are heating it:
- Use a heating mantle or a water bath. Never use a blowtorch or an open flame. Chloroform isn't flammable in the traditional sense, but high heat causes it to decompose into toxic fumes.
- Check your seals. Ground glass joints should be lightly greased or fitted with PTFE sleeves.
- Work in a certified fume hood. The face velocity should be at least 100 feet per minute.
Comparing Chloroform to Other Halogenated Solvents
| Solvent | Boiling Point ($^\circ\text{C}$) | Density ($\text{g/mL}$) | Vapor Pressure at $20^\circ\text{C}$ |
|---|---|---|---|
| Dichloromethane | 39.6 | 1.33 | 350 mmHg |
| Chloroform | 61.2 | 1.49 | 160 mmHg |
| Carbon Tetrachloride | 76.7 | 1.59 | 91 mmHg |
As you can see, chloroform sits in the "Goldilocks" zone. It's not as volatile (and hard to contain) as DCM, but it’s easier to remove than carbon tetrachloride. This middle-ground volatility is exactly why it remains a staple in analytical chemistry.
Actionable Steps for Lab Success
If you're planning on working with this solvent, don't just wing it. Follow these steps to ensure you're managing the boiling point for chloroform effectively:
- Calibrate your thermometer. In precise distillations, a $0.5$-degree error can lead to impurity in your distillate.
- Test for Peroxides and Phosgene. Use starch-iodide paper or specific phosgene test strips if the bottle has been sitting on the shelf for more than six months.
- Calculate your recovery. If you start with 100mL and recover only 70mL after boiling, your seals are leaking or your condenser isn't cold enough. Lower the temperature of your recirculating chiller to at least $5^\circ\text{C}$ when condensing chloroform.
- Dispose of it properly. Never pour chloroform down the sink. Halogenated waste must be separated from non-halogenated waste because the incineration process for chloroform requires higher temperatures to prevent the formation of dioxins.
Knowing the boiling point for chloroform is about more than just a number on a table. It's about understanding how energy, pressure, and chemistry collide in a flask. Whether you're a student trying to pass O-Chem or a veteran researcher, respecting that $61.15^\circ\text{C}$ threshold is the key to safe, efficient science.