The change from solid to liquid is a physical rather than chemical change because no chemical bonds have been broken. The individual particles—atoms, ions, or molecules—that made up the solid are the same individual particles that make up the liquid. What does change is the arrangement of the particles. In the liquid, the particles are at a higher temperature, having more energy than in the solid, and this allows them to move away from their nearest neighbors. The attractions between liquid particles, though less than those of solids, is still fairly strong. This keeps the particles close to each other and touching, even though they can around past one another. They cannot be pushed closer together, and so, like solids, liquids maintain their volume and cannot be compressed. Because their particles move freely around, liquids can flow, and they will assume the shape of any container. Like solids, the particles of liquids are close to each other; therefore, the amount of space occupied by liquids is quite close to that of their corresponding solids. However, because of the disorderly arrangement, the empty space between the liquid particles is usually slightly greater than that between the particles of the solid. Therefore, liquids usually have a slightly larger volume-that is, they are less dense-than solids. A very unusual exception to this is the case of ice melting to form water, when the volume actually decreases. The crystalline lattice of ice has a cage-like structure of H2O molecules with big, open spaces in the middle of the cages. When the ice melts and the crystal breaks down, the cages collapse and the molecules move closer together, taking up less space. Consequently, a given weight of water occupies more volume as ice than as liquid. In other words, ice is less dense than water. Therefore, ice floats on liquid water. Also, a full, closed container of water will break as it freezes because the ice must expand. A water pipe may break if it freezes in winter because of this unusual property of water.
Arrangement of particles in matter depends on its phase. They're close and organized in solids, close but irregular in liquids, and far apart and irregular in gases. The state of a substance depends on the balance between the kinetic energy of the individual particles (molecules or atoms) and the intermolecular forces. The kinetic energy keeps the molecules apart and moving around, and is a function of the temperature of the substance and the intermolecular forces try to draw the particles together.
Molecules in liquids are held to other molecules by intermolecular interactions, which are weaker than the intramolecular interactions that hold molecules and polyatomic ions together. The three major types of intermolecular interactions are dipole–dipole interactions, London dispersion forces (these two are often referred to collectively as van der Waals forces), and hydrogen bonds.
11.2.1 Ion-Dipole Forces
11.2.2 Dipole-Dipole Forces
11.2.3 London Dispersion Forces
11.2.4 Hydrogen Bonding
11.2.5 Comparing Intermolecular Forces
Surface tension, capillary action, and viscosity are unique properties of liquids that depend on the nature of intermolecular interactions. Surface tension is the energy required to increase the surface area of a liquid. Surfactants are molecules that reduce the surface tension of polar liquids like water. Capillary action is the phenomenon in which liquids rise up into a narrow tube called a capillary. The viscosity of a liquid is its resistance to flow.
11.3.1 Viscosity
11.3.2 Surface Tension
Fusion, vaporization, and sublimation are endothermic processes, whereas freezing, condensation, and deposition are exothermic processes. Changes of state are examples of phase changes, or phase transitions. All phase changes are accompanied by changes in the energy of a system. Changes from a more-ordered state to a less-ordered state (such as a liquid to a gas) are endothermic. Changes from a less-ordered state to a more-ordered state (such as a liquid to a solid) are always exothermic. 11.4.1 Energy Changes Accompanying Phase Changes
11.4.2 Heating Curves
11.4.3 Critical Temperature and Pressure
Because the molecules of a liquid are in constant motion and possess a wide range of kinetic energies, at any moment some fraction of them has enough energy to escape from the surface of the liquid to enter the gas or vapor phase. This process, called vaporization or evaporation, generates a vapor pressure above the liquid. Molecules in the gas phase can collide with the liquid surface and reenter the liquid via condensation. Eventually, a steady state or dynamic equilibrium is reached. vapor pressure – measures tendency of a liquid to evaporate 11.5.1 Explaining Vapor Pressure on the Molecular Level
11.5.2 Volatility, Vapor Pressure, and Temperature
11.5.3 Vapor Pressure and Boiling Point
The states of matter exhibited by a substance under different temperatures and pressures can be summarized graphically in a phase diagram, which is a plot of pressure versus temperature. Phase diagrams contain discrete regions corresponding to the solid, liquid, and gas phases. The solid and liquid regions are separated by the melting curve of the substance, and the liquid and gas regions are separated by its vapor pressure curve, which ends at the critical point.
11.6.1 the Phase diagrams of H2O and CO2
A crystalline solid can be represented by its unit cell, which is the smallest identical unit that when stacked together produces the characteristic three-dimensional structure. Solids are characterized by an extended three-dimensional arrangement of atoms, ions, or molecules in which the components are generally locked into their positions. The components can be arranged in a regular repeating three-dimensional array. The smallest repeating unit of a crystal lattice is the unit cell.
11.7.1 Unit Cell
11.7.2 The Crystal structure of Sodium Chloride
11.7.3 Close Packing of Spheres
The major types of solids are ionic, molecular, covalent, and metallic. Ionic solids consist of positively and negatively charged ions held together by electrostatic forces; the strength of the bonding is reflected in the lattice energy. Ionic solids tend to have high melting points and are rather hard. Molecular solids are held together by relatively weak forces, such as dipole–dipole interactions, hydrogen bonds, and London dispersion forces. Metallic solids have unusual properties.
11.8.1 Molecular Solids
11.8.2 Covalent-Network Solids
11.8.3 Ionic Solids
11.8.4 Metallic Solids
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