Matter exists in one of three physical forms: a solid, a liquid, or a gas. These states are determined by how constituent particles are arranged and move relative to one another. Changing a substance from one state to another, such as turning ice into water, is known as a phase transition. These transitions are physical alterations, meaning the chemical identity of the substance remains unchanged.
The Essential Role of Energy and Particle Movement
The ability to change a substance’s state depends on managing the energy possessed by its particles. All particles are constantly in motion, described as kinetic energy. When energy is added, particles move faster and vibrate more vigorously, increasing their kinetic energy. The state of a substance balances this particle movement against the attractive forces that hold molecules close together.
To cause a phase change, the energy input must be sufficient to overcome these attractive forces. In a solid, strong forces lock particles into fixed positions where they only vibrate slightly. In a liquid, forces are weaker, allowing particles to slide past one another. The shift to gas occurs when energy nullifies the attractive forces, resulting in highly energetic and widely spaced particles.
The amount of energy required for a phase transition is unique to every substance, depending on the strength of molecular attractions. For example, iron requires more energy to overcome its strong metallic bonds than water needs for its weaker molecular attractions. Controlling the energy provided tips the balance between attraction and motion, forcing a transition.
Causing the Solid-to-Liquid Transition
The transformation from a solid to a liquid, known as fusion, begins when energy is applied to the solid structure. This energy initially increases the vibration of the particles, causing the temperature to rise. Particles eventually reach a specific point where their increased kinetic energy destabilizes the fixed positions of the crystal lattice. This temperature is the melting point, a characteristic physical property of a pure substance.
At the melting point, the substance absorbs additional energy without any further temperature rise. This absorbed energy is the latent heat of fusion, dedicated entirely to breaking the attractive bonds holding the solid structure. This energy increases the potential energy of the molecules, allowing them to overcome their fixed constraints. For water, 334 kilojoules are required to melt one kilogram of ice at its melting point.
Melting continues until every particle has absorbed enough latent energy to break free. Once converted to a liquid, particles flow and slide past one another, remaining relatively close. Further energy added will increase the liquid’s temperature, as the energy is no longer needed to break intermolecular attractions. Industrial processes like smelting rely on controlling this energy input to reach the high melting points of substances like iron, which melts at approximately 1,538 degrees Celsius.
Causing the Liquid-to-Gas Transition
To transition from a liquid to a gas, the remaining attractive forces between molecules must be overcome entirely, a process called vaporization. This phase change is accomplished in two ways: through boiling (a bulk process) or through evaporation (a surface phenomenon). Both require a significant energy input, often greater than that needed for the solid-to-liquid transition.
Boiling occurs when the liquid reaches its specific boiling point, where the liquid’s vapor pressure equals the surrounding atmospheric pressure. As energy is added, particles throughout the liquid gain enough kinetic energy to escape intermolecular attractions. This leads to the rapid formation of vapor bubbles that rise and escape the surface. Similar to melting, the liquid’s temperature remains constant while the latent heat of vaporization is absorbed to convert the remaining liquid into gas.
Evaporation is a slower process that can occur at any temperature below the boiling point. Within a liquid, particles have a range of kinetic energies. A small fraction of molecules near the surface randomly possess enough energy to break free from the liquid’s surface tension. These high-energy molecules escape as gas, leaving behind the lower-energy molecules. Because the highest-energy particles are constantly leaving, evaporation causes a cooling effect on the remaining liquid.