Kaplan MCAT General Chemistry Review

Our discussion in this chapter will focus on the three phases of matter, with particular emphasis on liquids and solids. (The gaseous phase has been discussed extensively in the previous chapter.) When the attractive forces between molecules (i.e., van der Waals forces, etc.) overcome the kinetic energy that keeps them apart in the gas phase, the molecules move closer together, entering the liquid or solid phase. The liquid and solid phases are often referred to as the condensed phases because of their higher densities compared to that of the gaseous phase. Molecules in the liquid and solid phases have lower degrees of freedom of movement than those in the gaseous phase as a result of the stronger intermolecular forces that dominate in the liquid and solid phases. After characterizing fluids and solids, we will review phase equilibria and phase diagrams. We will conclude our consideration of the phases of matter by reviewing the colligative properties of solutions.

Solids

Key Concept

Because the molecules in liquids and solids are much closer together than those in a gas, intermolecular forces must be considered as though there is no such thing as “ideal” behavior. These forces do allow us to predict behavior, although differently so.

We recognize solids by their rigidity and resistance to flow. The intermolecular attractive forces among the atoms, ions, or molecules of the solid matter hold them in a rigid arrangement. Although particles in the solid phase do not demonstrate linear motion, this does not mean that they do not possess any kinetic energy. The motion of particles in the solid phase, however, is mostly limited to vibration. As a result, solids have definite shapes (usually independent of the shape of a container) and volumes. Note that on the MCAT, solids (and liquids, for that matter) are considered to be incompressible; that is, a given mass of any solid or liquid will have a constant volume regardless of changes in pressure.

For most substances, the solid phase is the densest phase. A notable, and tested, exception to this generalization is water. Water in its solid phase (ice) is less dense than it is in its liquid phase due to the grater spacing between the molecules in the crystalline structure of ice. The spacious lattice of ice crystals is stabilized by the hydrogen bonds between water molecules. Water molecules in the liquid phase also interact through hydrogen bonds, but because the water molecules are moving around, the lattice arrangement is absent, and the molecules are able to move closer to each other. In fact, water’s density reaches a maximum around 4°C. The density decreases at temperatures above 4°C because the increasing kinetic energy of the water molecules causes the molecules to move further apart. Between 4°C and 0°C, the density decreases because the lattice organization of hydrogen bonds is beginning to form.

Key Concept

Crystalline structures allow for a balance of both attractive and repulsive forces to minimize energy.

The molecular arrangement of particles in the solid phase can be either crystalline or amorphous. Crystalline solids, such as the ionic compounds (e.g., NaCl), possess an ordered structure; their atoms exist in a specific three-dimensional geometric arrangement or lattice with repeating patterns of atoms, ions, or molecules. Amorphous solids, such as glass, plastic, and candle wax, lack an ordered three-dimensional arrangement. The particles of amorphous solids are fixed in place but not in the lattice arrangement that characterizes crystalline solids. Most solids are crystalline in structure. The two most common forms of crystals are metallic and ionic crystals.

Figure 8.1

Ionic solids are aggregates of positively and negatively charged ions that repeat according to defined patterns of alternating cations and anions. As a result, the solid mass of a compound, such as NaCl, does not contain discrete molecules (see Chapter 4, Compounds and Stoichiometry). The physical properties of ionic solids include high melting points, high boiling points, and poor electrical conductivity in the solid state but high conductivity in the molten state or in aqueous solution. These properties are due to the compounds’ strong electrostatic interactions, which also cause the ions to be relatively immobile in the solid phase. Ionic structures are given by empirical formulas that describe the ratio of atoms in the lowest possible whole numbers. For example, the empirical formula BaCl₂ gives the ratio of barium to chloride atoms within the crystal.

Bridge

Ionic solids often have extremely strong attractive forces, thereby causing extremely high melting points.