Melting point

The melting point of a substance is the temperature at which it transitions from a solid to a liquid state. At this temperature, the solid and liquid phases coexist in thermodynamic equilibrium. The melting point depends on pressure and is usually measured under standard atmospheric conditions. When the same temperature is considered during the reverse transition—from liquid to solid—it is termed the freezing point or crystallisation point. In practice, melting points are often more precisely determined than freezing points because substances may supercool, forming liquids below their expected freezing temperatures until nucleation occurs.

Temperature, Pressure, and Phase Transition

Most substances have melting and freezing points that are approximately identical. For instance, mercury melts and freezes at the same temperature. However, some materials exhibit hysteresis, meaning their melting and solidification temperatures differ. Agar is a classic example: it melts at a higher temperature but solidifies at a significantly lower one.
Under standard pressure (1 atmosphere), the melting point of pure ice is extremely close to 0°C, measured with high precision at approximately 0.002519°C. Due to the possibility of supercooling, water without nucleation sites can remain liquid down to −40°C before freezing. At the opposite end of the phase spectrum, helium does not solidify at any temperature under atmospheric pressure, requiring substantially higher pressures for solidification.
The melting point varies with pressure. Substances that are denser in their solid state typically have melting points that increase with pressure. Water is a notable exception: ice is less dense than liquid water, so its melting point decreases as pressure rises. Silicon, germanium, gallium, and bismuth are among other substances exhibiting similar behaviour. Under extreme pressures, melting points may shift dramatically; for example, silicon melts at 1415°C at atmospheric pressure but at about 1000°C at pressures exceeding 10 GPa.

Examples of Melting Points

A wide range of melting points is observed in nature and materials science. Tungsten possesses the highest melting point among elemental metals at approximately 3422°C, making it ideal for high-temperature applications, such as in incandescent lamp filaments. Carbon, often cited as having an exceptionally high melting point, does not melt at atmospheric pressure. Instead, it sublimates at about 3900 K unless subjected to pressures above roughly 10 MPa, under which a liquid carbon phase can form.
Refractory compounds, particularly those combining transition metals and carbon or nitrogen, hold some of the highest known melting points. Hafnium carbonitride (Hf–C–N) possesses the highest experimentally confirmed melting point of any known compound, exceeding 4200 K. Computational predictions indicate that the alloy HfN₀.₃₈C₀.₅₁ may melt at around 4400 K, a value supported by later experimental confirmation.

Methods of Melting Point Determination

A variety of laboratory instruments and techniques are used to measure melting points, depending on the substance and its expected melting range.
General methods for ordinary solidsA simple melting point apparatus consists of an oil bath, magnifier, and a capillary tube containing grains of the test material. As the temperature rises, the observer records the point at which the crystals disappear. Modern instruments frequently feature automated optical detection. A Kofler bench, which contains a metal strip with a temperature gradient, allows multiple samples to be tested rapidly by placing them at different points along the strip. Differential scanning calorimetry (DSC) is a more advanced method providing accurate measurements of melting points along with enthalpy of fusion.
Industrial processes sometimes perform continuous online measurements, such as freeze-point testing of diesel fuels, enabling real-time quality control.
Techniques for refractory materialsExtremely high melting materials require specialised methods. Black-body furnaces heated electrically or by induction are paired with optical pyrometers, which determine temperature by matching observed radiance with known radiance-temperature relationships. Because direct calibration may be limited, temperatures above known reference points (such as gold’s melting point at 1063°C) are determined using extrapolation via Planck’s law. This introduces uncertainties but remains the standard approach.
To measure the melting point of refractory metals, black-body cavities may be engineered inside long rods of the test material. When heated by passing very large currents through the rod, the internal cavity radiates as a black body until melting disrupts the cavity’s behaviour. More recently, containerless laser heating techniques combined with spectro-pyrometry allow extremely rapid heating and precise high-temperature measurement, reducing complications like vaporisation or reaction with container materials.

Thermodynamic Considerations

Melting requires two distinct energy contributions: the heat required to raise the solid’s temperature to the melting point and the heat of fusion, a latent heat that enables the phase transformation. Thermodynamically, at the melting point the Gibbs free energy change (ΔG) equals zero, indicating equilibrium between solid and liquid. Both entropy (ΔS) and enthalpy (ΔH) increase during melting, and their relationship is given by:
ΔS = ΔH / T
where T is the absolute temperature at the melting point.
The sensitivity of melting point to pressure is defined by the Clausius–Clapeyron relation, which links the slope of the phase boundary to changes in enthalpy and volume. Because the solid–liquid transition typically involves small volume changes, melting points are less pressure-sensitive than boiling points. Exceptions, such as water, occur when the liquid state is denser than the solid.
Extremely high pressures can alter melting behaviour significantly. The melting point of some materials decreases with increasing pressure before rising again at higher pressures, reflecting complex structural transitions in dense phases.