Chapters 10, 11, 24

CHAPTER 10Section 10.1Substances that are gases at room temperature tend to be molecular substances with low molar masses. Air, a mixture composed mainly of N2 and O2, is the most common gas we encounter. Some liquids and solids can also exist in the gaseous state, where they are known as vapors. Gases are compressible; they mix in all proportions because their component molecules are far apart from each other.Section 10.2to describe the state or condition of a gas, we must specify four variables: pressure (P), volume (V), temperature (T), and quantity (n). Volume is usually measured in liters, temperature in kelvins, and quantity of gas in moles. Pressure is the force per unit area. It is expressed in SI units as pascals, Pa (1 Pa = 1 N/m2). A related unit, the bar, equals 10^5 Pa. In chemistry, standard atmospheric pressure is used to define the atmosphere (atm) and the torr (also called the millimeter of mercury). One atmosphere of pressure equals 101.325 kPa, or 760 torr. A barometer is often used to measure the atmospheric pressure. A manometer can be used to measure the pressure of enclosed gases.Sections 10.3 & 10.4Studies have revealed several simple gas laws: For a constant quantity of gas at constant temperature, the volume of the gas is inversely proportional to the pressure (Boyle's Law). For a fixed quantity of gas at constant pressure, the volume is directly proportional to its absolute temperature (Charles's Law). Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules (Avogadro's hypothesis). For a gas at constant temperature and pressure, the volume of the gas is directly proportional to the number of moles of gas (Avogadro's Law). each of these gas laws is a special case of the ideal-gas equation.The ideal-gas equation, PV = nRT, is the equation of state for an ideal gas. The term R in this equation is the gas constant. We can use the ideal-gas equation to calculate variations in one variable when one or more of the others are changed. Most gases at pressures less than 10 atm and temperatures near 273 K and above obey the ideal-gas equation reasonably well. The conditions of 273 K and 1 atm are known as the standard temperature and pressure (STP). In all applications of the ideal-gas equation we must remember to convert temperature to the absolute-temperature scale (the Kelvin scale).Sections 10.5 & 10.6Using the ideal-gas equation , we can relate the density of a gas to its molar mass: M = dRT/P. We can also use the ideal-gas equation to solve problems involving gases as reactants or products in chemical reactions.In gas mixtures the total pressure is the sum of the partial pressures that each gas would exert if it were present alone under the same conditions (Dalton's law of partial pressures). the partial pressure of a component of a mixture is equal to its mole fraction times the total pressure: P1 = X1Pt. The mole fraction is the ratio of the moles of one component of a mixture to the total moles of all components. In calculating the quantity of a gas collected over water, correction must be made for the partial pressure of water vapor in the gas mixture.Sections 10.7The