AP Chemistry: Chapter 5-6 Student Notes
Objectives
Chapter 5: Gases
 Kinetic Molecular Theory of Gases
 Understand pressure and its units
 Boyles-Charles-Gay-Lussac-Avagadro’s Laws (Conceptual and
Calculations)
 Ideal Gas Law
 Gas Stoichiometry
 Daltons Law of Partial Pressures (and applications to Gas Stoichiometry)
 Grahams Law
 Real verses Ideal Gases
Chapter 6: Thermochemistry
 The First Law of Thermodynamics
 Energy in a chemical reaction (exothermic and endothermic)
 Enthalpy ∆H (4 ways to calculate it)
 Finding ∆H using calorimetry
 Hess’ Law version 1: Adding reactions
 Hess’ Law version 2:  Pr oducts   Re ac tan ts
 Different kinds of ∆H
Kinetic Molecular Theory of Gases
Kinetic = _____________________________
Molecular = _______________________
So the KMT is the ______________ ____________________ theory
All molecules _____________
Solids move ______________
Liquids ________________
Gases move in _____________- _______________ motion
How fast do they move?
Use KE = ½ mv2: Solve for velocity
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Example: Find the velocity of an oxygen molecule in this room
Effusion verses Diffusion
Effusion = (Draw a picture)
Diffusion = (Draw a picture)
Grahams Law Conceptual
Graham’s Law Equation
Example: Calculate the effusion rates of hydrogen gas and uranium hexafluoride.
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The Meaning of Temperature:
Temperatue is the ___________ ___________ _______________ of a molecule in a
sample.
KE = 3/2 RT
Understand Pressure and it’s units
Pressure is caused by ____________________
Pressure is measured with a __________________
Toricellian Barometer
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Units of Pressure
1.0 atm = 760 mmHg = 760 torr = 101.32 kPa
Pressure Conversion Practice
Boyle-Charles-Gay Lussac-Avagadro’s Law
Pressure verses Volume: Boyles Law
Graph and Equation
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Volume verses Temperature (Charles’ Law)
Graph and Equation
Pressure verses Temperature: Gay-Lussac Law
Graph and Equation
Combined Gas Law
Equation
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Gas Law Examples:
Wilson at the
beach
Wilson at the top of Pikes Peak
Ideal Gas Law
Actually a form of the combined Gas Law
Derive Ideal Gas Law
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Example: “Wilson” is filled with oxygen gas. He has a volume of 3.25 L and is at 22C
and 0.75atm. How many grams of oxygen are in “Wilson’s” head?
Gas Stoichiometry
Rules:
1. Balanced Equation
2. Label what you know and what you don’t know
3. Use PV=nRT to solve for either volume or moles of the gas in the equation.
4. Use dimensional analysis to determine answer
5. Step 3 and 4 are sometimes flipped.
Example 1 Gas Stoichiometry:
Solid potassium chlorate (KClO3) decomposes to produce solid potassium chloride and
oxygen gas. What volume of oxygen gas, measured at 40°C and 655 mmHg, will be
produced when 13.5 g of potassium chlorate is decomposed?
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Example 2: Gas Stoichiometry
When the following reaction occurs:
P(s)
+
O2(g) 
P2O5(s)
How many grams of P2O5 is produced when 82.54 mL of oxygen at 6000 K and 45 atm is
completely consumed?
Daltons Law of Partial Pressures
The total pressure is equal to the sum of the partial pressures
Equation
Application (Collecting a gas over water)—Draw Diagram
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Daltons Law Example 1:
When silicon dioxide reacts with carbon by heating, the following reaction occurs:
SiO2(s) +
3C(s) 
SiC(s)
+
2CO(g)
What will be the volume of carbon monoxide collected over water that will be produced
at 22.0˚C and 657mm when 96.25 grams of SiO2 completely reacts?
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Example 2:
A .500 L container contains nitrogen gas at 0.800 atm and 0°C. If the highest pressure
the container can withstand before exploding is 3.0 atm, what is the highest temperature
to which the gas can be heated? Assume the volume is constant. What is the original
temperature in Kelvin? Should T2 be bigger or smaller from the change in pressure?
Real verses Ideal Gases
No gas is “ideal,” but some are ___________ _______________
Real Gases have
1. ________________________________
2. ________________________________
Real Gas Equation:
2

n 
 Pobs  a   x(V  nb)  nRT
 V  

Notice that it is similar to the ideal gas law with two new variables which are constants:
see table below.
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Chapter 6: Thermochemistry
First Law of Thermodynamics
Energy cannot _________________
The sum total of energy in the universe ________________
Thermodynamics means (examine the word)
Energy in a Chemical Reaction
Energy can flow two ways in a reaction
Into: Endothermic
Out: Exothermic
Called Enthalpy: Symbol is ∆H
Four Ways to Calculate ∆H
4 ways to calculate ?H
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Finding ∆H using calorimetry
Heat Equation
Specific Heat
 Amount of energy it takes to raise one gram of a pure substance one degree
Celsius
 The lower the specific heat the __________ the temperature change when heated
 The higher the specific heat the ___________ the temperature change when
heated.
Units on ∆H
Solving for ∆H
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Calorimetry Example 1: Calculating Enthalpy
3.25 grams of Mg is dropped into 125mL of hydrochloric acid. The initial temperature of
the calorimeter is 18.5˚C and the final temperature is 25.6˚C. Assume that the heat
capacity of the calorimeter is 4.86 J/g˚C. Calculate the enthalpy of the reaction.
Calormetry Example 2: Calculating heat capacity
A 46.2-g sample of copper is heated to 95.4˚C and then placed into a calorimeter
containing 75.0 g water at 19.6˚C. The final temperature of the metal and water is
21.8˚C. Calculate the heat capacity of the copper, assuming that all the heat lost by the
copper is gained by the water.
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Hess’ Law version 1: Adding reactions
How to find ∆H by adding reactions:
 All Elements in their standard state have a ∆H equal to _____________
 So we can write the formation of several compounds from their elements. This is
called the _______________ and is symbolized by __________________
H f : The f stands for _____________
The degree symbol stands for _______________________
Examples heats of formation reactions:
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To find ∆H from reactions you must:
1. Add the “source reactions” such that they cancel out to get the reaction that
you are looking for
2. When you multiply a source reaction the value of ∆H for the reaction is
multiplied by the same factor
3. when you reverse a reaction the sign of ∆H is switched
4. Sometimes it is easier to multiply the main reaction by a factor and then work
the problem. If you do this you must then divide the ∆H of the reaction by
that same factor at the end.
Examples of finding ∆H using the ADD REACTIONS METHOD
Example 1
Calculate the enthalpy for the combustion of C to CO:
C(s) + ½ O2(g)  CO(g) ∆H=?
Using these reactions:
C(s) + O2(g)  CO2
CO(g) + ½ O2  CO2 (g)
∆H= -393.5 kJ
∆H = -283.0 kJ
Example 2
Given the data:
N2(g) + O2(g)  2NO(g)
2NO(g) + O2(g)  2NO2(g)
2N2O(g)  2N2(g) + O2(g)
∆H=+180.7 kJ
∆H = -113.1 kJ
∆H -163.2 kJ
use Hess’ law to calculate ∆H for the reaction
N2O(g) + NO2(g)  3NO(g)
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Hess’ Law version 2:  Pr oducts   Re ac tan ts
This method is the simplest of them all.
1. Underneath the reaction write down the values that you find in the table that starts
on page A-21 in the back of your book
2. Add up the products multiplying the values by any coefficients
3. Add up the reactants multiplying any values by the coefficients.
4. Do the calculation:  Pr oducts   Re ac tan ts
Example 1: Finding the ∆H of a reaction using Hess’ Law method 2
Find ∆H for the following reactions:
NH3(g) + HCl(g)  NH4Cl(s)
C3H8(g) + O2(g)  CO2(g) + H2O(g)
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Different kinds of ∆H
∆H can sometimes be confusing because there are many different “kinds” of ∆H. This
table is a summary of the different kinds of ∆H.
H f
H fus
H vap
H solid
H cond
H sub
H comb
H sol

HCH
4
H rxn
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