Melting point does not follow a single direction down every group in the periodic table. In some groups it increases, in others it decreases, and in a few it does both. The trend depends entirely on the type of bonding holding the atoms together: metallic bonds get weaker going down a group, while the weak attractions between individual molecules get stronger.
Why Metals Melt at Lower Temperatures Down a Group
In metallic elements, atoms are held together by a “sea” of shared electrons that acts as a kind of glue between positively charged metal cores. As you move down a group, each atom has more electron shells, making the atom physically larger. That extra size pushes the outer electrons further from the nucleus, weakening the pull between the positive cores and the shared electrons. The result: less energy is needed to break the structure apart, so the melting point drops.
The alkali metals (Group 1) are the textbook example. Lithium, at the top, melts at about 180 °C. By the time you reach cesium at the bottom, the melting point has fallen to just 28 °C, barely above room temperature. Each atom still contributes only one electron to the metallic bond, but the growing distance between the nucleus and those bonding electrons makes each successive bond weaker. The increasing nuclear charge you might expect to help is completely offset by additional layers of inner electrons that shield the outer ones from the nucleus.
Why Nonmetals Melt at Higher Temperatures Down a Group
Nonmetals that exist as small molecules rely on a completely different force to hold them in a solid. These van der Waals forces (also called London dispersion forces) arise from the constant, random motion of electrons around each atom. At any given instant, electrons can cluster slightly to one side of an atom, creating a tiny temporary charge imbalance. That fleeting imbalance induces a matching one in a neighboring atom, and the two briefly attract each other.
Larger atoms have more electrons and more surface area for these temporary charge shifts to occur. The bigger the atom, the stronger and more frequent these fleeting attractions become. Research using atomic force microscopy has confirmed this directly: the measured attractive force between noble gas atoms scales with atomic size, with xenon interactions being significantly stronger than those of krypton or argon.
This is why melting points climb steadily down groups of nonmetals. Among the halogens (Group 17), fluorine melts at a frigid -220 °C while iodine melts at 114 °C, a solid you can hold in your hand at room temperature. Noble gases (Group 18) follow the same pattern. Helium solidifies only near absolute zero (about 1 K), while radon’s melting point is 202 K, more than 200 times higher.
Noble Gas Melting Points
- Helium: 0.95 K
- Neon: 24.7 K
- Argon: 83.6 K
- Krypton: 115.8 K
- Xenon: 161.7 K
- Radon: 202.2 K
Groups Where the Trend Reverses Partway Down
Some groups contain elements that change their bonding character from top to bottom, creating a melting point trend that rises, then falls, or the reverse. Group 14 (the carbon family) is the clearest case. Carbon, at the top, forms a giant covalent network where every atom is locked to four neighbors by strong bonds. This gives diamond one of the highest melting points of any substance, above 3,500 °C. Silicon and germanium, the next two elements, are metalloids with similar (though slightly weaker) network structures and still-high melting points.
But tin and lead, at the bottom of Group 14, are metals. They conduct electricity, bend easily, and melt at much lower temperatures because their atoms are held together by metallic bonds rather than a rigid covalent framework. The shift from giant covalent structure to metallic bonding means the melting point drops dramatically as you go from carbon to lead. This isn’t a simple “increases” or “decreases” trend; it’s a change in the fundamental way atoms stick together.
Group 2: When the Trend Gets Messy
Not every group follows a clean line even when all its members share the same bond type. The alkaline earth metals (Group 2) are all held together by metallic bonding, so you might expect a smooth decrease like the alkali metals. In practice, the trend is irregular. Beryllium melts at about 1,287 °C, magnesium at 650 °C, calcium at 842 °C (higher than magnesium despite being lower in the group), strontium at 777 °C, and barium at 727 °C.
These irregularities come from differences in crystal structure. The way atoms pack together in the solid state varies between Group 2 elements, and some packing arrangements produce stronger metallic bonding than others. The general direction is still downward from beryllium to barium, but individual elements can buck the trend because of these structural differences.
How to Predict the Trend for Any Group
The key question is always: what type of bonding holds the solid together?
- Metallic bonding (Groups 1, 2, and transition metals): Melting points generally decrease down the group. Larger atoms mean weaker attraction between metal cores and shared electrons.
- Molecular substances held by van der Waals forces (Groups 17, 18): Melting points increase down the group. Larger atoms with more electrons create stronger temporary attractions between molecules.
- Mixed bonding character (Groups 14, 15): The trend depends on where in the group you are. Elements near the top may form giant covalent networks with very high melting points, while metallic elements near the bottom melt at much lower temperatures.
If you’re answering an exam question, state the bonding type first, then explain the direction. A blanket statement that melting points “increase” or “decrease” down a group will only be correct for specific groups. The periodic table’s melting point trends are one of the clearest illustrations that the same underlying principle, the growing size of atoms, can push a property in opposite directions depending on context.