AP Chemistry: Chapter 5-6 Student Notes

ObjectivesChapter 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: Different kinds of ∆H

Kinetic Molecular Theory of GasesKinetic = _____________________________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

Example: Find the velocity of an oxygen molecule in this room

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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 unitsPressure is caused by ____________________

Pressure is measured with a __________________

Toricellian Barometer

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Units of Pressure1.0 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 LawEquation

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Gas Law Examples:

Ideal Gas LawActually a form of the combined Gas Law

Derive Ideal Gas Law

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Wilson at the top of Pikes PeakWilson at the beach

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 StoichiometryRules:

1. Balanced Equation2. Label what you know and what you don’t know3. Use PV=nRT to solve for either volume or moles of the gas in the equation. 4. Use dimensional analysis to determine answer5. 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 StoichiometryWhen 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 PressuresThe 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:

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 ThermodynamicsEnergy cannot _________________

The sum total of energy in the universe ________________

Thermodynamics means (examine the word)

Energy in a Chemical ReactionEnergy can flow two ways in a reaction

Into: Endothermic

Out: Exothermic

Called Enthalpy: Symbol is ∆H

Four Ways to Calculate ∆H

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4 ways to calculate ?H4 ways to calculate ?H

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 Enthalpy3.25 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 capacityA 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 reactionsHow 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 __________________

: 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 for2. When you multiply a source reaction the value of ∆H for the reaction is

multiplied by the same factor3. when you reverse a reaction the sign of ∆H is switched4. 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 METHODExample 1Calculate 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 ∆H= -393.5 kJCO(g) + ½ O2 CO2 (g) ∆H = -283.0 kJ

Example 2Given the data:N2(g) + O2(g) 2NO(g) ∆H=+180.7 kJ2NO(g) + O2(g) 2NO2(g) ∆H = -113.1 kJ2N2O(g) 2N2(g) + O2(g) ∆H -163.2 kJ

use Hess’ law to calculate ∆H for the reactionN2O(g) + NO2(g) 3NO(g)

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Hess’ Law version 2: 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 coefficients3. Add up the reactants multiplying any values by the coefficients.4. Do the calculation:

Example 1: Finding the ∆H of a reaction using Hess’ Law method 2Find ∆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.

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