Chapter 6 Gases, Liquids, and Solids States of matter: the physical state of matter depends on a balance between the kinetic energy of particles, which tends to keep them apart, and the attractive forces between them, which tend to bring them together. he attractive forces between molecules are the same in all three states (gas, liquid, and solid). However, in the gaseous state, the kinetic energy of the molecules is great enough to overcome the attractive forces between them. Kinetic energy increases with increasing temperature. Kinetic molecular theory (properties of an ideal gas):. Gases consist of particles, either atoms or molecules, constantly moving through space in straight lines, in random directions.. he average kinetic energy of gas particles is proportional to the temperature in kelvins. he higher temperature, the faster they move, and the greater their kinetic energy. 3. Molecules collide with each other, much as billiard balls do, bouncing off each other and changing directions (they may exchange kinetic energies). 4. Gas particles have no volume. 5. here are no attractive forces between gas particles (they do not stick together after a collision occurs). 6. Molecules collide with the walls of the container, and these collisions constitute the pressure of the gas. Pressure: the force per unit area. Force (F) P Area (A) Pressure has the different units: atmosphere, torr, mm Hg, in. Hg, Pascal, bar. atm 760 torr 760 mm Hg 0,35 Pascals 9.9 in. Hg bar Note: we use a barometer to measure atmospheric pressure (see Figure 5. on page 43). Note: we use a manometer to measure the pressure of a gas in a container (Figure 5.3 on page 44). Gas laws: Boyle s law: for a fixed mass of an ideal gas at a constant temperature, the volume of the gas is inversely proportional to the applied pressure. P P P P Charles s law: for a fixed mass of an ideal gas at a constant pressure, the volume of gas is directly proportional to the temperature in Kelvin (K).
Gay-Lussac s law: for a fixed mass of an ideal gas at constant volume, the pressure is directly proportional to the temperature in Kelvin (K). P P P P Combined gas law: the three gas laws can be combined: P Avogadro s law: equal volumes of gases at the same temperature and pressure contain equal numbers of molecules (regardless of their identity). Ideal gas law: under most experimental conditions, real gases behave sufficiently like ideal gases. n: amount of the gas in moles (mol) R: a constant for all gases (ideal gas constant) : volume of the gas in liters (L) : temperature of the gas in Kelvins (K) P: pressure of the gas in atmospheres (atm) P nr Note: at the Standard emperature and Pressure (SP) ( 0 C (73K) and P atm), one mole of any gas occupies a volume of.4 L. P (.00 atm) (.4 L) R n (.00 mol) (73 K) L.atm 0.08 mol.k Dalton s law: in a mixture of gases, each molecule acts independently of all the others. he total pressure, P, of a mixture of gases is the sum of the partial pressures of each individual gas. P P + P + P 3 + Partial pressure: the pressure that a gas in a mixture of gases would exert if it were alone in the container.
Attractive forces: London dispersion forces: extremely weak attractive forces between atoms or molecules caused by the electrostatic attraction between temporary induced dipoles. London dispersion forces exist between all molecules (polar or nonpolar). However, these forces are the only forces of attraction between nonpolar molecules. Ne Ne Dipole-Dipole interactions: the attraction between the positive end of a dipole of one molecule and the negative end of another dipole in the same or different molecule. Dipole- Dipole interactions exist between polar molecules. Hydrogen bonds: a noncovalent force of attraction between the partial positive charge on a hydrogen atom bonded to an atom of high electronegativity (most commonly oxygen or nitrogen) and the partial negative charge on a nearby oxygen or nitrogen. H H O H O O Note: hydrogen bonds increase the surface tension and the boiling points of the liquids (higher boiling point than dipole-dipole interactions and London dispersion forces). apor pressure: the partial pressure of a gas in equilibrium with its liquid form in a closed container. Equilibrium: a condition in which two opposing physical forces are equal. In a closed container with an equilibrium condition, the number of vapor molecules reentering the liquid equals the number escaping from it (as long as the temperature does not change).
Boiling point: the temperature at which the vapour pressure of a liquid is equal to the atmospheric pressure. At this temperature, a liquid starts boiling. Normal boiling point: the temperature at which a liquid boils under a pressure of atm. For example, 00 C is the normal boiling point of water because that is the temperature at which water boils at atm pressure. he boiling point of covalent compounds depends on three factors:. Intermolecular forces: higher attractive forces higher boiling point London dispersion forces < Dipole-Dipole interactions < Hydrogen bonds. Number of sites for intermolecular interaction (surface area): Larger surface area of the molecule (more electrons) more sites for London dispersion forces higher boiling point. For example: C 6 H 4 (hexane) has a larger surface area and it has more electrons than CH 4 (methane). Because of its larger surface area, there are more sites for London dispersion forces. herefore, hexane has the higher boiling point than methane. 3. Molecule shape: shape of the molecules can affect the boiling point. A linear molecule has a higher boiling point than a spherical molecule (if both compounds have the same intermolecular forces and the same molecular weight). Because, a linear molecule has a larger surface area than a spherical molecule. As surface area increases, contact between adjacent molecules, the strength of London dispersion forces, and boiling points all increase. Fusion (melting): changing phase from the solid state into the liquid state. Sublimation: a transition from the solid state directly into the vapour state without going through the liquid state (example: sublimation of dry ice (CO )) Amorphous solid: its atoms are arranged randomly (examples: wax and glass). hey have much lower melting points than network solids. Network solid or network crystal: in these solids, a very large number of atoms are connected by covalent bonds. hese atoms are arranged in a symmetric form (crystalline form). hey are hard and they have very high melting points (examples: diamond and quartz). Heating curve: it shows us the phase changes (a change from one physical state (gas, liquid, or solid) to another). Note: during the phase changing (evaporation, melting, sublimation, condensation, freezing), the temperature of the sample stays constant.
Heat added (cal) Heating curve of ice Phase diagram: by using this diagram, we can show all phase changes for any substance. emperature is plotted on the x-axis and pressure on the y-axis. he lines separating the different states and contain all the boiling points (A-C), all the melting points (A-B) and all the sublimation points (A-D). A phase diagram illustrates how one may go from one phase to another by changing the temperature and the pressure. riple Points: At this unique point on the phase diagram, all three phases coexist.