How do you prepare a solid sample for melting point determination?

Weigh approximately 1 mg of the solid substance and transfer it to a watch glass. Crush the solid using a metal spatula until it becomes a fine powder. Load the powder into a capillary tube by tapping the open end on the powder and dropping it through a glass tube repeatedly until 1–2 mm of packed solid fills the bottom of the capillary.

Why is the heating rate important when measuring melting points?

Heating rate control prevents experimental errors that can distort results. Heat at 5°C/min initially, then reduce to 1–2°C/min as you approach the expected melting point. This slower rate near the transition allows accurate observation of when the first liquid drop appears and when the last solid melts, ensuring precise melting point range determination.

What intermolecular forces explain the difference between naphthalene and urea melting points?

Naphthalene exhibits primarily London dispersion forces, which are weak intermolecular interactions. Urea, by contrast, is capable of hydrogen bonding, creating much stronger intermolecular forces in the solid state. These stronger forces require higher temperatures to overcome, resulting in urea's significantly higher melting point compared to naphthalene.

How does an impurity affect the melting point of a pure substance?

Impurities cause freezing point depression, lowering and broadening the melting point range. A pure substance melts over a narrow range at a specific temperature, but when an unknown impurity is added to urea, the melting point range becomes broader and shifts to a lower temperature, reflecting the disrupted crystal structure.

What is the purpose of using a mineral oil bath in melting point determination?

A mineral oil bath provides uniform, controlled heating around the capillary sample. The oil distributes heat evenly and allows precise temperature control via a hotplate and magnetic stir bar. This setup ensures the thermometer and sample reach thermal equilibrium, enabling accurate temperature readings at the onset and completion of melting.

What safety precautions should you take during a melting point experiment?

Wear a lab coat, splash-proof safety glasses, and nitrile gloves throughout the experiment. Conduct the heating setup in a fume hood to contain any vapors. Ensure the rubber band securing the capillary to the thermometer is positioned at least 8–9 cm above the solid and above the mineral oil level to prevent contact with heat sources.

How do you record accurate melting point data during heating?

Observe the capillary tube continuously as temperature increases. Record the oil bath temperature when the first drop of liquid appears in the capillary, marking the melting point start. Continue heating slowly and record the temperature when the last piece of solid melts, marking the melting point end. This range defines the substance's melting point.

Melting Points

Concepts Instructor Prep Student Protocol

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The melting point of a substance is the temperature at which that substance starts changing from the solid phase to the liquid phase. At this temperature, the liquid and solid phases are in equilibrium. With additional heat, the substance will melt completely. But what determines a substance’s melting point? Let’s think about solids and liquids. A solid’s molecules hold each other in a rigid, ordered structure called a lattice, while a liquid’s molecules have weaker interactions and move around.

Heating a solid transfers energy to the molecules. With enough energy, the molecules overcome the forces keeping them in the lattice and start moving around. In other words, if we heat a solid enough, it melts into a liquid. So, the melting point depends on the energy it takes to overcome the forces between the molecules, or the intermolecular forces, holding them in the lattice. The stronger the intermolecular forces are, the more energy is required, so the higher the melting point is.

Many intermolecular forces depend on how strongly atoms in the molecule attract electrons — or their electronegativity. Nitrogen, oxygen, fluorine, and chlorine are highly electronegative, while carbon, hydrogen, and sulfur are only moderately electronegative. Bonds between atoms with significantly different electronegativities are polar. For instance, a typical carbon-oxygen bond is polar, but a typical carbon-hydrogen bond is not.

A molecule’s electrons spend more time around its most electronegative atoms, giving it a slight negative charge on that side and a slight positive charge on the other side. This is called a dipole. If the dipole isn’t canceled out by an equal and opposite dipole in the same molecule, the molecule has a permanent dipole and is polar.

Now, let’s discuss three important intermolecular forces: hydrogen bonding, dipole-dipole interactions, and London dispersion forces. Hydrogen bonding occurs between an electron-withdrawing atom with a lone pair of electrons and a hydrogen bound to a more electronegative atom. Hydrogen bonds are among the strongest intermolecular forces.

Dipole-dipole interactions occur between polar molecules. In an attractive dipole-dipole interaction, the negative side of one dipole aligns with the positive side of another dipole. Dipole-dipole interactions are generally weaker than hydrogen bonds.

London dispersion forces come from brief, random shifts in a molecule’s electron distribution, which cause corresponding shifts in nearby molecules. These random shifts happen in every molecule, so this is one of the few interactions available to nonpolar molecules. London dispersion forces are among the weakest intermolecular forces.

Earlier, we predicted that stronger intermolecular forces corresponded to higher melting points. We can see this in action with hexadecane, 2-hexadecanone, and hexadecanoic acid. As the strength of the intermolecular interactions available to each molecule increases, so does the melting point.

Intermolecular forces aren’t the only factor that determines the melting point of a substance. Its purity significantly affects its melting and freezing points in an effect called ‘freezing-point depression’. This effect means that a solution has a lower freezing point than the pure solvent does. That’s why streets are sprinkled with salt when it gets very cold. If any water collects on the street, the salt quickly dissolves to make a solution with a much lower freezing point than pure water.

In a solid, impurities are incorporated into the lattice structure. These areas often have weaker intermolecular interactions, making parts of the structure easier to disrupt. So, compared to a pure solid, melting starts at a lower temperature and occurs over a wider temperature range.

In this lab, you will measure the melting points of two known organic compounds and then analyze a mixture to explore how impurities affect the melting point range.

Print Procedure Steps

Source: Lara Al Hariri and Ahmed Basabrain at the University of Massachusetts Amherst, MA, USA

Determining the Melting Points of Naphthalene, Urea, and an Unknown Mixture
In this experiment, you will measure the melting point range of two known substances, naphthalene and urea, by observing the melting phenomenon during heating. You'll then analyze a mixture of urea and an unknown substance to observe how the unknown impurity affects the melting point range.
- Put on a lab coat, splash-proof safety glasses, and nitrile gloves.
- Weigh ~1 mg of naphthalene using the analytical balance, and carefully, bring it back to your bench.
- Transfer the naphthalene onto a watch glass, and then crush it using the metal spatula until it becomes a fine powder.
- Carefully pick up one capillary tube and tap the open end on the powder to force some of the solid into the tube.
- Obtain a long glass tube from the common bench and hold it upright over your bench. Pick up the capillary with the open end upward and then drop it through the glass tube. Allow the capillary to bounce off the benchtop to pack the solid at the bottom.
- Retrieve the capillary and tap its open end on the naphthalene powder again to force more of it into the capillary. Continue loading naphthalene into the capillary and dropping it through the glass tube until 1 – 2 mL of powder is packed at the bottom of the capillary.
- Use a rubber band to attach the capillary to a digital thermometer, with the bottom of the capillary just above the bottom of the thermometer. Note: Make sure that the rubber band is at least 8 – 9 cm above the solid in the capillary.
- Set up a stirring hotplate and a lab stand in a fume hood and clamp the thermometer above the hotplate.
- Place the provided mineral oil bath on the hotplate and add a magnetic stir bar. Start stirring the bath at a moderate speed over low heat.
- Lower the thermometer and capillary setup into the oil bath until the sample is about 5 cm deep in the bath. Note: Make sure that the rubber band is above the mineral oil.
- Increase the hotplate temperature to 60°C, which is about 20° below the melting point of naphthalene in the literature. Note: It will take about 10 min for the entire oil bath to reach the same temperature as the hotplate.
- Once the bath reaches 60°C, heat it to 75°C at a rate of 5°C/min. When the bath is close to the melting point of naphthalene (80°C), heat the bath by 1 – 2 °C/min while carefully observing the solid in the capillary.
- At the first drop of liquid in the capillary, record the temperature of the oil bath. Keep watching the capillary tube as you slowly increase the temperature. Record the temperature when the last piece of solid melts.
  Table 1: Melting Points of Naphthalene, Urea, and an Unknown Mixture
  
  Melting point temperature (°C)
  Substance Start End
  Napthalene
  Urea
  Urea + Unknown (trial 1)
  Urea + Unknown (trial 2)
  Click Here to download Table 1
- Carefully remove the capillary and thermometer from the oil bath and set them aside to cool.
- Weigh roughly 1 mg of urea and transfer it to a clean watch glass. Crush the urea to a powder. Load the powder into a capillary tube as before.
- Attach the capillary to the thermometer with the rubber band. Heat the oil bath to 110°C, which is about 20° below the reported melting point of 132°C for urea.
- Insert the thermometer and capillary tube into the oil bath. Heat the bath to about 125°C at 5°C/min while watching the capillary.
- When the bath reaches 125°C, heat the bath by 1 – 2°C /min. Record the temperature when the first drop of liquid appears and when the last bit of solid melts.
- Turn off the hotplate and remove the capillary and thermometer from the oil bath to allow them to cool.
- Test the sample of urea mixed with an unknown impurity, which was prepared by your instructor. Perform two runs of this sample, so weigh 2 mg of the sample on the balance and transfer it to a watch glass.
- Crush the sample to powder and load two capillaries by tapping them on the powder and dropping them through the glass tube like before.
- Once the oil bath has cooled to at least 80°C, turn on the stir motor and set the hotplate to 80°C. Attach one of the loaded capillaries to the thermometer and clamp them in the oil bath.
- In the first test run, increase the temperature by 10°C/min until the sample has melted to determine the rough melting point range. Record the temperature when you see the first drop of liquid form and the last drop of liquid form to determine the rough melting point range.
- Turn off the hotplate and allow it to cool. Once the oil bath has cooled to at least 80°C, attach the second sample to the thermometer and clamp them in the bath. Turn the stir motor back on and set the hotplate to 80°C.
- For the second test, heat the bath by 5°C/min until you are close to the rough melting point range, and then continue heating by 1 – 2°C/min to accurately capture the melting point range.
- When you have finished measuring the melting point of the urea mixed with the unknown substance, turn off the hotplate and detach the capillary tube from the thermometer. Dispose of your capillary tubes in the glass waste container.
- Return the mineral oil to your instructor, and clean all of your glassware using detergent and water.
Results
- First, look at the melting point range of naphthalene. These differences may be caused by experimental errors, such as heating the sample too quickly, or by impurities in the sample.
- Next, check the results for the pure urea sample. Why is urea's melting point so much higher than naphthalene's? Looking at the structure of naphthalene, the intermolecular forces must be primarily London dispersion forces. Urea is capable of hydrogen bonding, so the intermolecular forces in solid urea are much stronger than the forces in solid naphthalene.
- Finally, look at the results from the urea sample with the impurity. The melting point range should be broader and lower than the range for pure urea because of freezing point depression. As expected, the melting point range obtained was 127 – 130°C, which is lower than the range for pure urea.