10+Chemical+Quantities+and+the+Mole

=10 Chemical Quantities and the Mole•Objectives= 10.1 The Mole: A measurement of Matter 10.2 Mole-Mass and Mole-Volume Relationships
 * What are three methods for measuring the amount of something?
 * How is Avogadro's number related to a mole of any substance?
 * How is the atomic mass of an element related to the molar mass of an element?
 * how is the mass of a mole of a compound calculated?
 * How do you convert the mass of a substance to the number of moles of the substance?
 * What is the volume of a gas at STP?

Chapter 10.3
 * How do you calculate the percent my mass of an element in a compound?
 * What does the empirical formula of a compound show?
 * How does the moelcular formula of a compound compare with the empirical formula?

=Outlined Notes 10.1 The Mole: A Measurement of Matter=

=__Amadeo Avogadro__= Italian Lawyer- Turned Chemist Major Contributions Include: =__The Mole__= Relationship between the mole and the periodic table:
 * Established differences between molecules (Oxygen and Nitrogen exist as molecules)
 * Reconciled the work of Dalton and Guy-Lussac
 * Established Avogadro's Principle: Equal volumes of all gasses at the same temperature and pressure contain the same number of molecules.
 * ** He did not determine Avogadro's number nor the mole **
 * Honored because the molar volume of all gasses should be the same
 * Much of Avogadro's work was acknowledged after his death by Stanislao Cannizarro
 * A counting unit: A large unit used to describe large quantities such as the number of atoms
 * 6.02x10^23 is knows as Avogadro's Number (1 Mole= 6.02x10^23)
 * The atomic mass is the quantity (grams) of 1 mole of that element
 * The units of atomic mass are grams/mole
 * Mass is used by chemists as a way of "counting" the number of atoms/molecules of a substance
 * 6.02x10^23 protons= 1 gram
 * 6.02x10^23 neutrons= 1 gram

A. The Mole- Mass Relationship 1. Use the molar mass of an element or compound to convert between the mass of a substance and the moles of a substance. B.The Mole-Volume Relationship 1. Avogadro's Hypothesis: states that equal volumes of gases at the same temperature and pressure contain equal numbers of particles. 2. Standard Temperature and Pressure (STP): means of temperature of 0 degrees celsius and and a pressure of 101.3 kPa or 1 atmospheric pressure (atm). 3. AT STP, 1 mol or 6.02 x 10^23 representative particles, of any gas occupies a volume of 22.4 L. 4. Molar Volume: 22.4 L
 * __10.2 Mole-Mass and Mole- Volume Relationships__**
 * **Mole/Weight Relationship Examples Using Helium** ||
 * //Moles Helium// || //# Helium Atoms// || //Grams Helium// ||
 * 1/4 || 1.505 x 1023 || 1 g ||
 * 1/2 || 3.01 x 1023 || 2 g ||
 * 1 || 6.02 x 1023 || 4 g ||
 * 2 || 1.204 x 1024 || 8 g ||
 * 10 || 6.02 x 1024 || 40 g ||

=__10.3 Percent Composition and Chemical Formulas__= A. Percent Compostion from Mass Data B. Percent Composition from the Chemical Formula
 * 1) __**The Percent Composition of a Compound**__
 * __Percent Composition-__ The relative amounts of of the elements in a compound.
 * The percent composition of K2CrO4 is K = 40.3%, Cr = 26.8%, and O = 32.9%.
 * These percents MUST total 100%
 * The percent by mass of an element in a compond is the number of grams of the element divided by the mass in grams of the compound, multiplied by 100%.
 * [[image:Equation1.jpg]]
 * [[image:Equation2.jpg]]
 * [[image:Equation3.jpg]]
 * [[image:Equation4.jpg]]

C. Percent Composition as a Conversion factor
 * You can use percent composition to calculate the number of grams of any element in a specific mass of a compound. To do this, multiply the mass of the compound by a conversion factor based on the percent composition of the element in the compound.


 * __Empirical Formulas__**
 * __Empirical Formula-__ Gives the lowest whole-number ratio of the atoms of the elements in a compound.
 * The emipirical formula of a compound shows the smallest whole-number ratio of the atoms in the compound.


 * __Molecular Formulas__**
 * The molecular formula of a compound is either the same as its experimentally determined empirical formula, or it is a simple whole-number multiple of its empirical formula.

(Mr. Elford if he was a mole) > || 12 books = 1 dozen books > > > || 6.02 x 1023 atoms of aluminum = 1 mole of aluminum. > > > || 6.02 x 1023 molecules of water = 1 mole of water. > > > || ==
 * [[image:http://lc.brooklyn.cuny.edu/smarttutor/core3_22/images/12books.gif]] || [[image:http://lc.brooklyn.cuny.edu/smarttutor/core3_22/images/alumfoil.gif]] || [[image:http://lc.brooklyn.cuny.edu/smarttutor/core3_22/images/waterglass.gif]] ||

= = =Reference Pages=

=Practice Problems= 1. Convert 2.50 moles of CuSO4 to grams. In order to find the molar mass, you must add up all the molar masses in the compound given: Cu = 63.5 S = 32.1 O4 = 64.0 molar mass = 159.6 2.50 moles CuSO4 || 159.6 g CuSO4 ||
 * || 1 mole CuSO4 ||

= 399g CuSO4

1. Convert 0.954 moles of CuSO4 to grams. Cu = 63.5 S = 32.1 O4 = 64.0 molar mass = 160 0.954 x 160 = 152.64 (since the original problem shows only two significant digits we can show only two significant digits in our answer.) the final answer = 150g
 * To find the molar mass add up all the masses in the compound given.

1. How many moles of particles are contained in 34g of NH3(formula mass = 17g per mole)? N= 14.0 H3= 3.03 molar mass = 17.0 34/17 = 2.0 moles
 * To find the molar mass add up all the masses in the compound given.

1. How many molecules are there in 4.0 moles of H2O2? 4.0 x 6.02 x 1023 = 24 x 1023 = 2.4 x 1024
 * Remember: Avogadro's number is6.02 x 1023.



The math is: 11.2 ÷ 22.4 x 65.4 =

The math is: 11 x 22.4 = .69g CO2 • 1 mol CO2 • 2 mol CO • 28g CO 44g CO2 • 2 mol CO2 • 1 mol CO = .44g CO

.87g O2 • 1 mol O2 • 2 mol O3 • 48g O3 32g O2 • 3 mol O2 • 1 mol O3 = .87g O3

425g S • 1 mol S • 1 mol O2 • 22.4L O2 32.1g S • 1 mol S • 1 mol O2 = 297L O2

5.68L H2 • 1 mol H2 • 2 mol Na • 23g Na 22.4L H2 • 1 mol H2 • 1 mol Na = 11.6g Na

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=Sample Test= [] [] [] [] [] [] [] [] [] [] []
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= = =**Links**= [] [] [] [] [|http://genchem1.chem.okstate.edu/CCLI/Startup.html] [|__http://www.chemistry.co.nz/mole.htm__] [] [] []
 * 1) More Quizzes - http://lrc-srvr.mps.ohio-state.edu/under/chemed/qbank/quiz/bank2.htm
 * 2) 10 Chemical Quantities and the MoleLOTS of Mole Resources - []