1. Module 8 Gases

2. Substances that Exist as Gases At 25 o C and 760 torr (1 atm), the following substances exist as gases: Elements Compounds H2H2 HeHFNH 3 N2N2 NeHClNO 2 O2O2 ArHBrSO 2 O 3 (Ozone)KrHIH2SH2S F2F2 XeCOHCN Cl 2 RnCO 2 CH 4 (methane)

3. Properties of Gases Compressibility-Gases can be compresses into smaller volumes. Their densities can be increased by applying pressure. If too much pressure is applied they will become liquids or solids. Indefinite shape and volume-A gas can be made to fit the vessel in which it is placed. Expansion-Gases expand without limits so that gas samples completely and uniformly occupy the volume of a container. Miscibility or diffusion-Gases diffuse into each other completely. Low density-Because the density of a gas is small, it is usually measured in g/L. The average kinetic energy (KE) of gaseous molecules is directly proportional to their absolute temperature (KE  Provided that T is the same, KE of different gas molecules is equal. The amounts and properties of gases are described in terms of T (Temperature in K), P (pressure), V (Volume they occupy), and n (number of moles present).

4. Characteristics of Ideal Gases The equations that will be used apply only to ideal gases or those behaving very closely as ideal gases. An ideal gas is considered to have the following characteristics: a. each particle (atom or molecule) has a negligible volume b. has no attractive forces c. undergoes elastic collisions (they rebound after colliding and bounce back in the opposite direction with the same speed). Real gases can behave as ideal gases by avoiding: a. extremely low temperatures b. very small volumes c. very high pressures

5. Pressure Pressure is defined as force per unit area. The common units of pressure are: a. atmosphere (atm) b. torr c. mm of mercury (mm Hg) d. Pascal e. bar 1.00 atm = 760 torr = 760 mm Hg = 1.01 x 10 5 Pa = 1.01 bar. At sea level the pressure is 760 torr. Above sea level it is less than 760 torr. Below sea level it is greater than 760 torr.

6. Barometer The barometer is a device used to measure pressure: h  mm Hg) P atm mercury surface Schematic of a barometer P atm = h

7. Manometer A manometer is a partially filled glass tube of mercurey with one arm open to the atmosphere (side P 1 ) and the other arm (side P 2 ) attached to a gas tank. It can be used to measure the pressure of gas inside of a container. The difference in the height of mercury from one arm of a manometer to the other is 58 mm. If the atmospheric pressure is 1.02 atm, then what is the pressure (in torr) of the gas in the tank that is attached to the manometer? P gas = P atm +  h Pgas = 1.02 atm x (760 mm Hg/1 atm) + 63 mm = 839 mm Hg = 839 torr

8. Standard Temperature and Pressure (STP) and Standard Molar Volume STP = 0.00 o C (273.15K) and 1 atm (760 torr) At STP 1 mol of an ideal gas occupies 22.4 L.

9. Gas Laws Boyle’s Law: V  1/P (at constant T and n) Therefore, Volume is inversely related (inversely proportional) to pressure. If the pressure of a gas at constant temperature and number of moles is doubled, the volume will be cut in half.

10. Gas Laws Charles’ Law: V  T(K) (at constant P and n) For example, if the temperature (in K) of a gas at constant pressure and composition is tripled, the volume will triple.

11. Gas Laws Avogadro’s Law- V  n (constant T and P) Example: If the number of moles of a gas at constant temperature and pressure is quadrupled, its volume will quadruple. At the same T & P equal volumes of gases contain the same number of moles and molecules (at STP 1 mol of gas occupies a volume of 22.4 L).

12. Combined Gas Law Taking into account Boyle’s, Charles’, and Avogadro’s law we arrive at the combined Gas Law: P 1 V 1 = P 2 V 2 If n is constant: P 1 V 1 = P 2 V 2 n 1 T 1 n 2 T 2 T 1 T 2 Any unit of P and V as long as they are the same on both sides of the equation. T 1 and T 2 must be in K.

13. Ideal Gas Equation The ideal gas equation is as follows: PV = nRT R is the ideal gas constant. It can have the units:.0821 L x atm or 62.4 L x torr K x mol K x mol

14. Molar Mass and Density of Gases Since PV = nRT and n = g/MM and d = g/V, then: PV = g R T and V = g x R x T MM P x MM And: d = g = g x P x MM V g x R x T d = P x MM R x T

15. Dalton’s Law of Partial Pressures At constant volume and temperature, the total pressure extended by a mixture of ideal gases is the sum of the partial pressures of those gases. Gas A Gas B Gas A & B P A =200 torrP B =150 torrPT = PA + P B = 350 torr Place A and B in same container At constant T and V conditions.

16. Dalton’s Law of Partial Pressures P T = P A + P B + P C + …= Since PV = nRT then P= nRT/V P T = n A RT + n B RT + n C RT + … = (n A + n B + n C +…)(RT) V V V V There is a simple relation that exists between total pressure and individual pressures. Consider a sample containing gases A and B. Dividing P A by P T we obtain: n A RT P A = V = n A = X A Where X A is P T (n A + n B )RT n A + n B the mole V fraction of A. Mole fraction is a dimensionless quantity that expresses the ratio of moles of one component to the number of total moles. The mole fraction can be any number between 0 and 1. Note that the sum of all mole fractions must equal to one. If two components are present, then: X A + X B = 1 Rearranging the above equation ( P A /P T = X A ), we can express the partial pressure of A and B as: P A = X A P T P B = X B P T Therefore, in a sample containing two gases (A & B), the total pressure of the sample is given by: P T = P A + P B = X A P T + X B P T

17. Vapor Pressure of Water Gases that are insoluble and don’t react with water can be collected over water. The partial pressure exerted by the water vapor above the liquid is called its vapor pressure. Every liquid shows characteristic vapor pressures that vary with temperature. If a gas is collected over water it is saturated with water vapor, and at atmospheric pressure (P T = atmospheric pressure): P T = P gas + P water or P gas = P T – P water gas gas + water vapor Vapor pressure of water (P water ) in torr near room temperature: Temp ( o C) P water Temp ( o C) P water Temp ( o C) P water 20 17.5 23 21.1 26 25.2 21 18.7 24 22.4 27 26.7 22 19.8 25 23.8 28 28.4

18. Graham’s Law: Diffusion and Effusion of Gases Diffusion is the intermingling of gases. Effusion is the leaking out of gases through a small hole or orifice. Rate A /Rate B = (M B /M A ) 1/2 The rate is the rate of effusion: V/t mL / s or L / min, etc..

19. Gay-Lussac’s Law (Law of Combining Volumes) At constant temperature and pressure, the volumes of reacting gases can be expressed as the simple whole number ratio of the moles of gases that are reacting (V  M at constant T and P). If the reaction is carried out at STP, then 1 mol gas = 22.4 L H 2(g) + Cl (2) 2HCl (g 1 mol 1 mol 2 mol 1 volume 1 volume 2 volumes 22.4 L 22.4 L 44.8 L At STP