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Many people have asked for some additional stoichiometry questions.  So here you go.  Note that the exam will definitely ask some questions on converting between moles and grams.  In addition, the exam will likely have one or two somewhat more difficult stoichiometry calculations such as those below.   Remember, there is a lot more to know for this exam than just stoichiometry questions so don't over focus on these.  

How many moles of hydrogen are produced when 0.1 mol of sodium reacts with 0.1 mol of water?

2 Na(s) + 2H2O (l) --> H2(g) + 2 NaOH(aq)

answer 0.05 mol

How many moles of oxygen are required to completely burn 4.8 moles of butane?

2 C4H10 + 13 O2 --> 8 CO2 + 10 H2O

Answer: 31 moles

Consider the following reaction:

PbCl2 + Na2CrO4 -> PbCrO4 + 2 NaCl

How many grams of NaCl are produced when 0.5 moles of Na2CrO4 react.

answer: 58.5 g

How many moles of carbon dioxide are produced when 88.0 g of oxygen are reacted with an excess of butane, C4H10?

2 C4H10 + 13 O2 ---> 8 CO2 + 10 H2O

answer: 74.5 moles.

 Chapter 7 - States of Matter

Today - Gases:

I. About gases

A. some gases you might recognize

Atmosphere: 78% N2, 21% O2, 1% Ar

CO2 0.025% (1990), 0.036% today - increasing

Other gases you might be familiar with:

CO (carbon monoxide)

Chlorofluorocarbons (CFCs): CF2Cl2 and CFCl3

Cl2, H2, F2 

noble gases

B. Properties of gases

1. Gases are a collection of individual particles with rapid, random motion.

2. Energy of gases is kinetic energy (KE), which is energy due to motion.

molecules move faster at higher temperatures so the kinetic energy of a gas increases with temperature

3. Distance between gas atoms or molecules is large compared to their sizes

1 mole H2O(l) = 0.02L

1 mole H2(g) = 22.4 L

4. Gases are characterized by a pressure.  Pressure is a force per unit area.  Pressing a sharp object into your hand hurts more than pressing a dull object into your hand.  

The collisions of gas molecules with the container they are in exerts a pressure on it.

P = F / A where F = force and A = area

common units: atmosphere (atm), mm Hg = torr

1 atm = 760 mm Hg = pressure which can support a column of mercury 76cm high.  Alternatively, the pressure exerted by a column of mercury 76cm high.

1 torr = 1 mm Hg

1 atm = 760 torr

At sea level, the atmosphere exerts a pressure of 1 atm = 760 mmHg

Barometric pressure often measured in inches of Hg.  1 atm = 30 in Hg

5. In comparing gases and gas behavior we use STP - "standard temperature and pressure"

For STP:  T= 273  K, P = 1 atm

Under these conditions, 1 mole of gas occupies 22.4L, independent of what the gas actually is. 

Avogadro's hypothesis - equal volumes of gases at the same temperature and pressure contain equal numbers of molecules or particles. 

Because of this, we can interpret the coefficients of chemical reactions involving gases in terms of volumes!

For instance,


3 H2 (g)




usual interpretation

3 moles H2

1 moles N2


2 moles NH3

since moles are proportional to volume, the coefficients can also be interpreted as

3 liters H2

1 liter N2


2 liters NH3

II. Gas Laws describe how gases behave - Interrelating temperature, pressure and volume - NOTE: FOR QUANTITATIVE CALCULATION NEED TO USE TEMPERATURE IN KELVIN

1. Boyles law, pressure-volume relation, for a given mass of gas at constant T, the volume varies inversely with pressure.

T constant, 

increase pressure, volume goes down
decrease pressure, volume goes up

Mathematically:  P1V1 = P2V2 for constant temperature

Ex: A gas initially has a volume of 22.4L at a pressure of 1 atm.  It is compressed to a final volume of 11.2 L, what is the final pressure of the gas.

P1 = 1 atm, V1 = 22.4 L
P2 = ?, V2 = 11.2 L

Solve for P2 and plug in:

P2 = P1V1 / V2 = (1 atm)(22.4 L) / (11.2 L) = 2 atm.  Does this make sense??

2. Charles law, Volume-Temperature relation, for a pressure constant, the volume of a fixed mass of gas varies in direct proportion with temperature

P constant

increase T, volume goes up
decrease T, volume goes down

V1 / T1 = V2 / T2 for constant pressure

3. Temperature-pressure relation

V constant

increase T, P increases
decrease T, P decreases

4. Combined gas law - puts the three laws above together.

P1V1 / T1 = P2V2 / T2

III. Ideal gas law - interrelates the pressure, volume, temperature, and number of moles of a gas.  Above laws always dealt with the same amount of gas.

PV = nRT

n = number of moles,  R = 0.0821 L atm mol-1 K-1