Virtually everyone has seen the same substance in solid, liquid and gaseous states by the age of perhaps five at the latest: That substance is water. Below a certain temperature (0 °C or 32 °F), water exists in a "frozen" state as a solid. Between 0 °C and 100 °C (32 °F to 212 °F), water exists as a liquid, and past its boiling point of 100 °C/212 °F, water exists as water vapor, a gas.

Other substances that you may think of existing only in one physical state or another, such as a chunk of metal, also have characteristic melting and boiling points, which can be quite extreme in relation to everyday temperatures on Earth.

The melting and boiling points of elements, like many of their physical characteristics, depend largely on their position in the periodic table of elements and therefore on their atomic number. But this is a loose relationship, and other information you can gather from the periodic table of elements helps to determine a given element's melting point.

When a solid moves from a very cold temperature to a warmer one, its molecules gradually assume more kinetic energy. When molecules in the solid achieve a sufficient average kinetic energy, the substance becomes a liquid, wherein the substance is free to change shape in accordance with its container as well as gravity. The liquid has melted. (Going the other way, from liquid to solid, is called freezing.)

In the liquid state, molecules can "slide" past each other, and are not fixed in place, but lack the kinetic energy to escape into the environment. However, once the temperature becomes sufficiently high, the molecules can escape and move far apart, and the substance is now a gas. Only collisions with the walls of the container, if any, and with each other limit the gas molecules' movement.

Most solids assume a form at the molecular level called a crystalline solid, made from a repeated arrangement of molecules fixed in place to create a crystal lattice. The central nuclei of the involved atoms remain spaced a fixed distance apart in a geometric pattern, such as a cube. When sufficient energy is added to a uniform solid, this overcomes the energy "locking" the atoms in place, and they are free to jostle about.

A variety of factors contribute to the melting points of individual elements, such that their position on the periodic table is only a rough guide, and other issues must be considered as well. Ultimately, you should consult a table like the one in the Resources.

You might ask if larger atoms have inherently higher melting points, being perhaps harder to break apart because of more matter in them. In fact, this trend is not observed, as other aspects of individual elements prevail.

The atomic radii of atoms tend to increase from one row to the next but decrease across the length of the row. Melting points, meanwhile, increase across rows to a point, then sharply drop off at certain points. Carbon (atomic number 6) and silicon (14) can form four bonds with relative ease, but the atoms a step up on the table cannot, and they have far lower melting points as a result.

There is a rough relationship between atomic number and boiling point of elements as well, with the "jumps" to lower boiling points within rows followed by an increase happening in about the same places. Notably, however, the boiling points of the noble gases in the rightmost column (period 18) are barely higher than their melting points. Neon, for example, exists as a liquid only between 25 °C and 27 °C!