Download Report

Chemistry L 3, Gas laws: Chapter 12: Name ________________ Page 1
pg. 326-355 and Notes: Keep your Forces handout. We will not use kilopascals for pressure on
worksheets or tests. Show your work on all worksheets!!
Temperature and Pressure Conversions
1. Convert these Temperatures to Kelvin: K = °C + 273 Or to Celsius °C = K – 273
25°C!
27°C!
100 K !
273 K
!
2. Convert these pressures to atm: 1 atm
760 mm Hg 800 mm Hg!
380 mm Hg !
Or to mm Hg, ! 760 mm Hg
1 atm
0.75 atm!
0.25 atm!
Boyle’s law : P1V1 = P2V2
Or P V = constant (Change P and V; don’t change T and moles, p. 333)
3. An average adult has a “deep breath” lung capacity of about 4.5 L. 4.50 L of air is measured at
740 mm Hg pressure. It is compressed until its volume is 3.50 L. What is its new pressure
(assuming unchanged temperature)?!!
(951 mm Hg)
4. A 200 mL sample of H2 gas is collected at a pressure of 1.05 atm. If the temperature remains
constant, what volume will the gas occupy at 600 mm Hg?!
(266 mL)
Gay-Lussac’s Law: P1 = P2 or P = Constant! Change P and T; don’t change V, moles
T1
T2
T!
T Must be in Kelvin, p. 338
5. Some CO2 gas at 760 mm Hg pressure and – 73.0°C in a sealed metal container is heated to
27.0° C. What is its new pressure?! !
(1140 mm Hg)
6. The pressure in an automobile tire is 2.00 atm at 27.0°C. At the end of a road trip the pressure
has risen to 2.20 atm. What is the temperature of the air in the tire, assuming the volume has
not changed? Calculate the final T in Kelvin, then change to °C.!
(330 K, 57°C)
Deriving the Volume and Temperature relationship: !
Page 2
7. A student measured the Volume and Temperature of a gas in a flexible container. The Pressure
and amount of particles did not change. Here is their data.
Volume(mL)! Temperature (K)!
Volume(mL)!
Temperature (K)!
!
100 !
273 !
117!
320
!
110!
300 !
124.5!
340
a. Is the Volume vs. Temperature relationship Direct or Inverse?
b. As the Temperature of a gas Increases, does the Volume increase or decrease?
c. Would a graph of this data be Linear or a Hyperbola?
d. Is V times T or V/T a constant? Calculate the constant, include units!
(0.366 mL/K)
e. Explain how Kinetic Theory explains this relationship.
f. The Temp of this gas is changed until its volume is 146.5 mL. Was the gas heated or cooled from
the final state above (124.5 mL at 340 K)? Explain why you think so.
g. Calculate the final Temperature of the gas in part f, and convert your answer to °C. !
!
!
(400 K, 127°C)
Charles’ law : V1 = V2
or ! V = Constant ! Change V, T; don’t change P, moles
!
T1
T2 T!
T MUST be in Kelvin!!!! p. 335-7)
8. A balloon is placed in a freezer and has a Volume of 500 ml at – 23.0°C. It is then removed from
the freezer. Its Pressure stays same. What will be its volume at + 27.0°C?
(600 mL)
Problems from Combined gas law: P1 V1 = P2 V2 Change P, V, T; don’t change mol
(p. 339-340)!
T1
T2
(T ... you should know by now)
9. The gas in a balloon occupies 100 mL at 27°C and 570 mm Hg. What will be its pressure if the
volume is decreased to 50 mL and temperature to – 48.0°C?!
(855 mm Hg)
10. The pressure of a 2.00 L sample of a gas is 3.50 atm, and its temperature is 127°C. The gas is
heated until it has a new volume of 3.00 L and a new pressure of 5.25 atm. What is its new
temperature, in °C?!
(900 K, 627°C)
Chemistry Level 3 : Gas Laws
Name _____________________ Page 3
Ideal gas law problems : PV = nRT P must be in atm ! V must be in Liters
(p. 341-3)!
n is the number of moles of the gas
!
Remember: STP = 0°C, 1 atm!
R = 0.0821 L atm /mol K !
T must be in K
1. What is the pressure of 30.0 moles of N2 gas in a 55.0 L storage tank at 27.0°C?! (13.4 atm)
2. What is the volume occupied by 4.00 grams of O2 gas at STP? (change g––>mol)! (2.80 L)
3. a. A gas in a 500 mL storage tank has a pressure of 1520 mm Hg at 27.0°C. How many moles of
a gas are in this tank? !
(0.0406 mol)
!
b. If this gas is methane, CH4, what is the mass of the CH4 gas in the tank?!
(0.650 g)
4. a. A gas has a volume of 400 mL at 1 atm pressure and 27.0°C. How many moles of gas are in
this sample?!
(0.0162 mol)
!
!
!
b. This gas has a mass of 1.04 grams. Calculate the experimental formula weight of the gas.
!
(64.0 g/mol)
c. A test reveals the gas is made of Sulfur and Oxygen. Is this gas SO2 or SO3?
Dalton’s law of partial pressures: Ptotal = P1 + P2+ P3 + . . .(can use mm Hg or atm)
(p. 350-1) 5. On a Humid summer day the partial pressure of the N2 in the air is 0.783 atm, of O2 =
160 mm Hg and of H2O vapor = 25 mm Hg. What is the total pressure? !
(780 mm Hg)
6. Underwater environments usually contain a lot of Helium gas because N2 and O2 gas are bad
for us at high pressures. For example, O2 gas is toxic if its Partial pressure is above about 1.4
atm. In an underwater environment, the Total pressure is 20.0 atm and the partial pressure of
the He is 13700 mm Hg and of N2 is 800 mm Hg. What is the Partial pressure of the O2 in this
environment? Is it safe? !
(700 mm Hg, 0.92 atm, yes)
Graham’s Law:(p.352-3) Lighter gases move (diffuse, effuse) faster than heavier gases
Page 4
Gas A = !
Gas B =
Distance A moves = √Formula Weight B =! Rate A
Lighter, !
Heavier,!
Distance
B
moves
√
Formula
Weight
A!
Rate
B
Faster! !
Slower
!
!
!
7. Which gas effuses faster, Chlorine (Cl2) or Hydrogen (H2)? How much faster? (Hint: Find the
Ratio Rate A/Rate B) !
(6 x)
8. In an experiment, N2 gas diffuses 43.0 cm in the same time an unknown gas diffused 25.0 cm.
What is the formula weight of the unknown gas? If it's a Noble gas, which gas is it?
!
!
(82.8 g/mol)
9. During an experiment, O2 gas traveled 30.0 cm in the same time an unknown gas traveled 41.0
cm. What is the formula weight of the unknown gas? Is it H2S, HF or NH3? (17.1 g/mol)
Real and Ideal gases
10. What would happen to the volume of a gas under the following conditions? (p. 330-2)
(gets smaller, gets larger, stays the same, or can't tell)
a. Increase Temp, keep pressure and amount constant! ____________________
b. Decrease pressure, keep temp and amount constant! ____________________
c. Increase pressure, add gas, temp constant!
____________________
d. Increase temp, decrease pressure, amount constant! ____________________
e. Add gas to a metal cylinder!
____________________
f. Double the Kelvin temp, double the pressure!
____________________
11. Identify the following as characteristic of an Ideal Gas (I) or a Real Gas (R) (p. 344-5)
____ a. Particles are very small!
____ b. Particles have no volume
____ c. Will have 0 volume at 0 K!
____ d. Would be a solid at 0 K
____ d. Have weak attractive forces!
____ e. Have no attractive forces
____ f. Obey gas laws at high T, low P! ____ g. Obey gas laws at all T and P
____ h. Would never become a liquid! ____ i. Will eventually become a liquid
12. Which real gas would act “most” like an Ideal gas? Why? SO2 Rn HCl (polar) Ar
13. If a real gas has strong attractive forces, would its Volume at Low Temperature be more or less
than an ideal gas? Why do you think so?
Chemistry Level 3 : Gas Laws Review Name _____________________ Page 5
1. In an experiment, you collect 6.00 L of butane gas at 27.0°C and 2.00 atm total pressure. The
mass of the butane gas you collected is 28.29 g. a. How many moles of gas did you collect?
b. Calculate your experimental formula weight of butane gas. !
!
!
!
c. What is the Formula for butane: C2H6
(0.487 mol, 58.1 g/ mol)
C3H8
or
C4H10?
d. Oops! Your lab partner was squeezing the soda bottle (making it smaller) while you were
collecting data with the pressure sensor. Is your Measured Pressure too high, too Low, or no
difference? Explain why you think so.
e. Oops! The Pressure sensor is wrong, and it’s reading too high (so the actual Pressure is lower
than 2.0 atm). Is your calculated moles of gas too big or too small? Explain why you think so.
2. A certain gas diffuses 50.0 cm in the same time O2 gas diffuses 12.5 cm. What is the formula
weight of the gas? What gas is it?!
(2.0 g/mol)
3.
!
!
!
!
Consider the Pressure and Temperature relationship (Gay-Lussac’s Law)
a. Is it Direct or Inverse?
b. Would a graph of this data be Linear or a Hyperbola?
c. Is P times T or P/T constant?
d. What unit do you need to use to express your Temperature for this relationship to hold?
Why?
!
e. Explain how Kinetic Theory explains this relationship.
!
f. If there are no sources of error, what should be the Temperature at 0 pressure? Why?
(1c. C4H10 , d. Too high, e. too big; 3. Direct, linear, P/T, Kelvin, high T = particles move faster= more collisions =
higher P, 0 Kelvin=absolute zero=no motion,no collisions)
Page 6
4. 100 mL of a gas has a pressure of 380 mm Hg at –73°C. What is its new pressure at –173°C if
the volume increases to 200 mL?!
(95 mm Hg)
5. A metal gas cylinder (constant volume) at a temperature of 127°C contains a gas at 200 atm
pressure. If the pressure is reduced to 100 atm, what is its new temperature in °C?
! !
(200 K, –73°C)
6. The total pressure of N2 and O2 in an oxygen tent is 1.10 atm. If the partial pressure of the N2 is
456 mm Hg, what is the partial pressure of the O2 in the oxygen tent?!
(0.50 atm)
7. Which gas diffuses faster, He or F2? How much faster?!
(3.1x)
8. Describe a practical application or example from everyday life of Boyle’s Law (the Pressure and
Volume relationship).
9. An 80.0 Liter helium tank has a pressure of 100 atm at 77.0°C. How many grams of Helium are
in the tank?!
(1115 g)
!
b. How many balloons can the helium in this tank fill? Assume the balloons measure 10.0 L at
1.00 atm pressure and 27.0°C. !
(686)
10. A balloon with a volume of 1.50 L at 27.0°C is cooled to –173°C. What is its new volume
assuming the pressure stays the same?!
(0.50 L)
11. A hot air balloon has a volume of 100 L at 1.00 atm pressure. If the temperature stays constant,
what is its new volume at 0.80 atm pressure?!
(125 L)