Chapter 12- Gases LBS 172 Exam 1 Review I. What is a gas? a. Properties i. Non-definite volume, fills container, can flow, spread out, can be compressed b. Air is a gas composed of many gases i. Relatively stable gases are: 1. N 2 78.1% O 2 20.9% Ar 0.9% He 0.0005% H 2 0.00005% Kr Ne ii. Gases whose composition varies are: 1. CO 2 0.035% H 2 O 0-4% CO CH 4 0.0002% O 3 0.000004% II. What kinds of substances are gases? a. Nonmetals, small molecules i. Binary molecules: 1. H 2, N 2, O 2, F 2, Cl 2, O 3 (ozone) ii. Noble gases: 1. He, Ne, Ar, Kr, Xe, Rn b. Other gases i. CO 2 greenhouse gas (22x better than CH 4 but decays slower) ii. CH 4 greenhouse gas (cow burps release 400-500 L/day) iii. NH 3 iv. NO acid rain precursor v. NO 2 smog, acid rain precursor vi. N 2 O nitrous oxide is laughing gas used for dental work vii. H 2 S smell of rotten eggs viii. HCN hydrocyanic acid is a deadly poison that tastes like almonds ix. SO 2 acid rain precursor c. Any aroma is due to gaseous molecules III. Microscopic characteristics of gases a. Particles are randomly moving b. Particles have elastic collisions i. They bounce off of each other with no net loss of energy c. Particles exert pressure by bouncing off of container wall IV. Results in macroscopic world
a. Simple relationship between moles, pressure (P), volume (V), and temperature (T) b. Gases are compressible c. Gases have a much lower density than solids and liquids d. Gases will mix evenly and completely when put in the same container i. There is no such thing as immiscible gases V. Pressure a. Gases exert pressure by bouncing off of the walls of a container i. The more collisions, the more pressure ii. The higher the speed of the collisions, the more pressure b. Units of pressure i. Atmospheric pressure (at sea level) is 1.01x10 5 Pa = 101 kpa = 1 atm (atmosphere) c. How is pressure measured? i. The barometer was invented by Torricelli in the 1700s 1. It is a glass tube inverted in a dish of Hg and the atmospheric pressure keeps the Hg in the tube. Hg is used because it is the densest liquid at room temperature (1.0 atm) ii. 1.0 atm = 760 mm Hg = 29.92 in. Hg = 760 torr = 101 kpa = 1.0x10 5 Pa VI. Relationships between physical properties of gases a. Relationship between pressure (P) and volume (V) i. Boyle s Law was discovered in 1650 1. It states that V is inversely proportional to P a. P increases, V decreases 2. V α 1/P V = k (1/P) PV = k 3. P 1 V 1 = k = P 2 V 2 P 1 V 1 = P 2 V 2 b. Relationship between volume (V) and number of moles (n) i. Avogadro s Law 1. It states that number of moles is directly proportional to volume under the conditions of constant pressure and temperature a. V increases, n increases 2. V α n V = kn V/n = k 3. V 1 /n 1 = k = V 2 /n 2 V 1 /n 1 = V 2 /n 2 a. Note that this law is not dependent on the type of gas, only the number of moles of the gas. c. Relationship between volume (V) and temperature (T)
VII. VIII. IX. i. Charles Law 1. It states that temperature is directly proportional to volume under the conditions of constant pressure and number of moles of gas a. T increases, V increases 2. V α T V = kt V/T = k 3. V 1 /T 1 = k = V 2 /T 2 V 1 /T 1 = V 2 /T 2 a. Temperature must be converted to the Kelvin scale because it is necessary that the temperatures are absolute (positive numbers) d. Relationship between P, V, n, and T i. The Ideal Gas Law is a combination of the 3 laws 1. V α nt/p V = R (nt/p) PV = nrt a. R = 0.082057 L atm/mol K b. We often define standard temperature and pressure (STP) i. Standard T = 0º C = 273.15 K ii. Standard P = 1 atm Density of Gases (g/l) a. Gases have much lower densities than solids and liquids i. D = m/v 1. mass (m) = number of moles (n) x molar mass (M) a. m = n M n = m/m 2. PV = mrt/m m/v = PM/RT D = PM/RT a. Note that this law is only dependent upon pressure, molar mass, and temperature (NOT VOLUME!) Assumption behind calculations a. Dalton s Law of Partial Pressure (1801) i. It states that the total pressure of a mixture of gases is the sum of their individual pressures 1. P depends on moles of gas, not on chemical nature ii. P total = P A + P B + P C + iii. P total = X A P total +X B P total 1. P A = X A P total X A = P A /P total a. X A = mole fraction of a mixture of gases Kinetic Molecular Theory a. Large separations between molecules b. Random motion with collisions c. Average kinetic energy is proportionate to gas temperature
d. Energy Distribution e. i. From physical chemistry PV = (1/3)nMū 2, where n is the number of moles of gas, M is the molar mass of the gas, and ū 2 is the average of the square of speeds 1. Root Mean Square is the weighted average speed a. PV = nrt = (1/3)nMū 2 (ū 2 ) = (3RT/M) i. R = 8.314 J/mol K ii. Answer in m/s X. Gas Diffusion and Effusion a. Gas diffusion is the gradual mixing of molecules of one gas with molecules of another i. Graham s Law (1832)- At a constant pressure and temperature, the rate of diffusion of gases is proportional to the inverse of the square root of the molar masses XI. 1. Rate Rate 1 2 = M M 2 1 b. Gas effusion is the process by which gases escape out of a small opening in a container i. Gas effusion is also governed by Graham s Law 1. Furthermore, the rate of the escaping of gases is inversely proportional to time (slower rate = longer time) a. Rate Rate 1 2 = M M 2 1 = t 2 t1 Non-ideal gases a. PV = nrt assumes no interactions between molecules and assumes that molecules have no volume i. Non-ideal gas law: (P + a(n/v) 2 )(V-bn) = nrt
XII. 1. a(n/v) 2 corrects for intermolecular force assumption (stickiness) 2. bn corrects for molecular volume assumption ii. This formula becomes an issue at high pressures and low temperatures. 1. High pressures- less space so V becomes more important 2. Low temperatures- slower motion so it increases stickiness Problems involving gases a. PV = nrt can be used to get moles of gas from P, V, and T. Once you have moles, then it is just a normal limiting reagent or stoichiometry or titration b. Gas law shortcuts all revolve around the fact that in gases, V and P do not depend on the nature of the gas i. Implications 1. The ratio of n:v or n:p is constant (at a given T and P or V) a. So 2 moles of gas C takes up twice the volume of 1 mole of gas A under the same conditions Chapter 13- Intermolecular Forces, Liquids, and Solids I. Intermolecular forces in order of increasing strength a. Induced dipole-induced dipole, London dispersion forces, or Van der Waals forces (example- oil and oil) i. Two non-polar molecules interact with each other ii. London dispersion forces are present in all molecular interactions b. Dipole-induced dipole forces (example- oil and water) i. One polar and one non-polar molecule interact with each other 1. The process of inducing a dipole is called polarization a. The greater the surface area or molar mass of the molecules involved, the greater the polarizability ii. When a partial negative approaches an electron cloud, the electrons repel and the electrons around the partial positive shift, causing a slight positive charge on one side and a slight negative charge on the other side c. Dipole-dipole forces (example- methanol and methanol) i. Two polar molecules interact with each other
II. III. 1. It is important to be careful with molecules that are nonpolar but have polar bonds and partial charges (example- CO 2 ) d. Hydrogen bonding (example- water and water) i. An H atom is bound to an F, O, or N and that H atom is attracted to an F, O, or N on another molecule 1. F, O, and N are the three highest electronegative atoms, so when they bond with hydrogen, it creates a highly polar bond 2. F, O, and N are the smallest highly electronegative atom, so their charge per surface area is greater than that of, say, Cl ii. Hydrogen bonding is what allows ice to float 1. The density of water as a liquid is greater than the density of ice because hydrogen bonding in ice creates holes. e. Ion dipole forces (example- Na + and water) i. An ion charge creates a large dipole force Relationship between polarity and solubility a. Intermolecular forces are critical to solubility i. If the IMFs between 2 molecules are similar in strength and type, they will dissolve in each other Properties of liquids a. In liquids, molecules are closer together than in gases, therefore, IMFs are very important, whereas in gases they are not present i. To go from liquid to gas, IMFs must be overcome 1. In general, as IMFs increase, boiling point increases a. Hydrogen bonding causes unexpectedly high boiling points 2. H vap > 0 ALWAYS for transition from liquid to gas a. It will always take energy to go from liquid to gas, whereas energy will always be given off when going from gas to liquid ( H vap < 0)
b. As temperature increases, the percent of molecules with enough energy to turn into gas increases. If you look to the right of the red line, the area under the T 2 curve is greater than that of the T 1 curve b. If we put a liquid into a sealed container, then the liquid and the vapor will come into equilibrium (liquid gas rate = gas liquid rate) i. Pressure of gas is called vapor pressure 1. This is dependent on temperature (temperature increases, vapor pressure increases) c. Viscosity i. Resistance of a liquid to flow (high viscosity = slow flow) 1. As IMFs or molecular size increase, viscosity increases d. Surface tension i. Energy required to break the surface of a liquid 1. Surface molecules experience fewer molecular forces because there is nothing above them e. Capillary action i. Water will climb up a small tube due to attract to the walls of the tube by adhesive forces (between 2 different substances) and other water molecules are pulled along by cohesive forces (between water molecules) f. Paper chromatography i. Separation of molecules on paper 1. Mobile phase- water because it is more polar 2. Stationary phase- paper because it is non-polar IV. Phase diagrams (pressure vs. temperature) a. Depending on the conditions of T and P, a substance can exist as a gas, a liquid or a solid, or two or even three states can coexist in equilibrium i. Lines are equilibriums between two phases ii. Triple point is the only point where all three phases are at equilibrium iii. Above the critical point are the conditions under which a supercritical fluid can exist 1. This is the state where the difference between a gas and a liquid can not be deciphered V. Solids (4 types) a. Molecular solids (examples- ice, dry ice, sugar) i. Solids made of molecules that are held together by intermolecular forces
ii. Relatively low melting points because they are being held together by relatively weak forces iii. Non-conducting and brittle b. Ionic solids (example- salts such as NaCl) i. Solids made of ions that are held together by electrostatic attractions ii. High melting points because it requires a lot of energy to overcome their attractions iii. Non conducting and very brittle c. Covalent solids (examples- graphite (2D), diamond (3D)) i. Held together by covalent bonds in infinite 2D or 3D structure ii. Very high melting points due to the strength of their bonds iii. Usually used as insulators or semiconductors d. Metallic solids (examples- iron, silver, copper, etc.) i. Metal atoms held together by metallic forces and electrostatic attractions among metal ions and electrons ii. Wide range of melting points iii. Good electric conductivity and good heat conductivity