Learning Objectives Review the states of matter and their properties Describe how change in temperature will affect the state of matter. Previously, you were introduced to the three states, also called phases, of matter; solid, liquid, and gas. A phase is a certain form of matter that includes a specific set of physical properties. That is, the atoms, the molecules, or the ions that make up the phase do so in a consistent manner throughout the phase. Science recognizes three stable phases: the solid phase , in which individual particles can be thought of as in contact and held in place; the liquid phase , in which individual particles are in contact but moving with respect to each other; and the gas phase , in which individual particles are separated from each other by relatively large distances (see Figure \(\PageIndex{1}\)). The state of a substance depends on the balance between the kinetic energy of the individual particles (molecules or atoms) and the attractive forces between molecules, called intermolecular forces . The kinetic energy keeps the molecules apart and moving around, and is a function of the temperature of the substance. The intermolecular forces draw the particles together. A discussed previously, gasses are very sensitive to temperatures and pressure. However, these also affect liquids and solids too. Heating and cooling can change the kinetic energy of the particles in a substance, and so, we can change the physical state of a substance by heating or cooling it. Increasing the pressure on a substance forces the molecules closer together, which increases the strength of intermolecular forces. We take advantage of changes between the gas, liquid, and solid states to cool a drink with ice cubes (solid to liquid), cool our bodies by perspiration (liquid to gas), and cool food inside a refrigerator (gas to liquid and vice versa). We use dry ice, which is solid CO 2 , as a refrigerant (solid to gas), and we make artificial snow for skiing and snowboarding by transforming a liquid to a solid. In this section, we examine what happens when any of the three forms of matter is converted to either of the other two. These changes of state are often called phase changes. The six most common phase changes are shown in Figure \(\PageIndex{2}\). Energy Changes That Accompany Phase Changes Phase changes are always accompanied by a change in the enthalpy, \(\Delta H\), of a system. For example, converting a liquid, in which the molecules are close together, to a gas, in which the molecules are, on average, far apart, requires an input of energy (heat) to give the molecules enough kinetic energy to allow them to overcome the intermolecular attractive forces. The stronger the attractive forces, the more energy is needed to overcome them. Solids, which are highly ordered, have the strongest intermolecular interactions, whereas gases, which are very disordered, have the weakest. Thus any transition from a more ordered to a less ordered state (solid to liquid, liquid to gas, or solid to gas) requires an input of energy; the \(\Delta H\) is positive (endothermic). Conversely, any transition from a less ordered to a more ordered state (liquid to solid, gas to liquid, or gas to solid) releases energy; the \(\Delta H\) is negative (exothermic). The energy change associated with each common phase change is shown in Figure \(\PageIndex{2}\). \(\Delta H\) is positive for any transition from a more ordered to a less ordered state and negative for a transition from a less ordered to a more ordered state. Previously, we defined the enthalpy changes associated with various chemical and physical processes. The molar enthalpy of fusion (\(ΔH_{fus}\)), is the energy required to convert a solid to a liquid, a process known as fusion (or melting). As noted above, the process of melting requires energy and therefore, the (\(ΔH_{fus}\)) is positive. The reverse process of freezing would release energy making the (\(ΔH_{fus}\)) negative. The molar enthalpy of vaporization (\(ΔH_{vap}\)), is the energy required to convert a liquid to a gas, known as vaporization. Melting points, enthalpies of fusion, boiling points, and enthalpies of vaporization for selected compounds are listed in Table \(\PageIndex{1}\). Substance Melting Point (°C) ΔHfus (kJ/mol) Boiling Point (°C) ΔHvap (kJ/mol) N2 −210.0 0.71 −195.8 5.6 HCl −114.2 2.00 −85.1 16.2 Br2 −7.2 10.60 58.8 30.0 CCl4 −22.6 2.56 76.8 29.8 CH3CH2OH (ethanol) −114.1 4.93 78.3 38.6 CH3(CH2)4CH3 (n-hexane) −95.4 13.10 68.7 28.9 H2O 0 6.01 100 40.7 Na 97.8 2.60 883 97.4 NaF 996 33.40 1704 176.1 The substances with the highest melting points usually have the highest enthalpies of fusion; they tend to be ionic compounds that are held together by very strong electrostatic interactions. Substances with high boiling points are those with strong intermolecular interactions that must be overcome to convert a liquid to a gas, resulting in high enthalpies of vaporization. The enthalpy of vaporization of a given substance is much greater than its enthalpy of fusion because it takes more energy to completely separate molecules (conversion from a liquid to a gas) than to enable them only to move past one another freely (conversion from a solid to a liquid).