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Understanding the Mole Concept: From Dozens to Moles - Prof. Beverly A. Kelley, Study notes of Chemistry

This lecture outline from chem 1075 covers the concept of the mole as a unit of measure in chemistry. Topics include the comparison of dozens and moles, the definition and properties of the mole, mole calculations, and the relationship between atomic mass and molar mass. Slides include examples and practice problems.

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Pre 2010

Uploaded on 08/19/2009

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Chem 1075 Chapter 9 The Mole Concept Lecture Outline
Slide 2 Dozen compared to Mole
We use “dozen” to _________________________________; We know 1 dozen = _______________
We use “mole” to ___________________________________________________________________
Slide 3 The Mole
The mole (mol) is a unit of measure for the ________________________________________in a
chemical substance.
A ____________ is _____________________________________(abbreviated “N”) of particles or
_______________________________particles.
1 mol = Avogadro’s number = ___________________________ particles.
We can use the __________relationship to convert between the number of __________________ and
the _________________ of a substance.
Slide 4 How Big is a Mole?
The volume occupied by 1 mole of _________________________would be about the size of the
_______________________.
One mole of Olympic shotput balls has about the same _______________as the _________________.
Slide 5 Mole Calculations I
Use Avogadro’s number to convert between _________________________ and _______________:
__________________________particles = _________ mole
Slide 6 Mole Calculations I
How many sodium atoms are in 0.120 mol Na?
Slide 7 Mole Calculations I
How many moles of potassium are in 1.25 x 1021 atoms K?
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Download Understanding the Mole Concept: From Dozens to Moles - Prof. Beverly A. Kelley and more Study notes Chemistry in PDF only on Docsity!

Chem 1075 Chapter 9 The Mole Concept Lecture Outline Slide 2 Dozen compared to Mole We use “dozen” to _________________________________; We know 1 dozen = _______________ We use “mole” to ___________________________________________________________________ Slide 3 The Mole The mole (mol) is a unit of measure for the ________________________________________in a chemical substance. A ____________ is _____________________________________(abbreviated “N”) of particles or _______________________________particles. 1 mol = Avogadro’s number = ___________________________ particles. We can use the __________relationship to convert between the number of __________________ and the _________________ of a substance. Slide 4 How Big is a Mole? The volume occupied by 1 mole of _________________________would be about the size of the _______________________. One mole of Olympic shotput balls has about the same _______________as the _________________. Slide 5 Mole Calculations I Use Avogadro’s number to convert between _________________________ and _______________: __________________________particles = _________ mole

Slide 6 Mole Calculations I How many sodium atoms are in 0.120 mol Na?

Slide 7 Mole Calculations I How many moles of potassium are in 1.25 x 10^21 atoms K?

Slide 8 Practice

  • If you have 2.49 x 1024 molecules of CO2, then how many moles of carbon dioxide is that?
  • How many atoms of Fe are in 4.75 moles of iron?
  • How many moles would contain 3.56 x 1022 molecules of water?

Slide 9 Atomic Mass vs. Molar Mass

  • Atomic mass mass of iron is 55.85 amu. is the mass of an _________________ expressed in __________. The atomic
  • Molar mass is the mass of _______ ______________ of a substance expressed in ________. The molar mass of iron is 55.85 g.
  • Why are the numbers the same? Slide 10 Molar Mass
  • • Recall that 1 amu = ___________________g orCan show that 1mole Fe has a mass of 55.85g: ______________________amu = 1 g
  • Molar mass for iron is 55.85g
  • Can also think of as 55.85g = 1mole Fe or 55.85g/mole Slide 11 Calculating Molar Mass
  • The molar mass of a substance is the sum of the molar masses of each element in the formula.
  • What is the molar mass of magnesium nitrate, Mg(NO 3 ) 2? The sum of the molar masses is: 1 Mg x ___________ = ____________2 N x ____________ = ____________ 6 O x ____________ = ____________ Molar Mass = ___________ g / mole

Slide 18 Pressure and Units Pressure is defined as the ratio of ____________________per unit _________________ (i.e. psi)

Slide 19 Molar Volume of Gases ______mole gas = _____________________molecules gas = ___________L gas

Slide 20 Gas Density

  • The density of gases is much ___________than that of liquids.
  • • We can calculate the density of any gas _______________ easily.The formula for gas density at STP is:

Slide 21 Calculating Gas Density

  • What is the density of ammonia gas, NH 3 , at STP?
  • First we need the molar mass for ammonia;
  • The molar volume NH 3 at STP is 22.4 L/mol.
  • Density is mass/volume:

Slide 22 Molar Mass of a Gas

  • We can also use ______________________________ to calculate the ___________________ of an unknown gas.
  • • 1.96 g of an unknown gas occupies 1.00L at STP. What is the molar mass?We want g/mol, we have g/L.

Slide 23 Mole Unit Factors

  • We now have three interpretations for the mole: _____________ = _____________________ particles _____________ = _____________________ (in grams)
  • _____________ = _______________ at STP for a gasThis gives us 3 unit factors to use to convert between moles, particles, mass, and volume. Slide 24 Mole-Volume Calculation

Slide 25 Mole-Volume Calculation A sample of methane, CH 4 , occupies 4.50 L at STP. How many moles of methane are present?

Slide 26 Mass-Volume Calculation What is the mass of 3.36L of ozone gas, O 3 , at STP?

Slide 27 Molecule-Volume Calculation How many molecules of hydrogen gas, H 2 , occupy 0.500L at STP?

  • The molecular formula of octane is __________
    • The empirical formula of octane is ___________. Slide 34-35 Calculating Empirical Formulas (from experimental data) We can calculate the empirical formula of a compound from its experimental data. We can determine the mole ratio of each element from the mass to determine the formula of radium oxide, Ra?O?. A ________________ sample of radium metal was heated to produce ________________ of radium oxide. What is the empirical formula?
  • Radium + oxygen à radium oxide

The ____________________ of radium is ____________ g/mol and the ______________________ of oxygen is ________________ g/mol. (Comes from periodic table)

  • We get Ra______________ O _____________.
  • Simplify the mole ratio by dividing by the smallest number.
  • We get Ra ______________O______________ = RaO is the empirical formula. Slide 36-37 Empirical Formulas from % Composition We can also use percent composition data to calculate empirical formulas. Assume that you have ________________ of sample. Benzene is 92.2% carbon and 7.83% hydrogen, what is the empirical formula? If we assume ____________ of sample, we have _____________g C and __________g H. Calculate the moles of each element:

The ratio of elements in benzene is C_________H__________. Divide by the _________________________number to get the formula:

C ----------- H ------------- = C______ H _______ = _____________

Slide 38 Molecular Formulas The _____________________ formula for benzene is CH. This represents the __________________ ratio of C to H atoms of benzene. The actual molecular formula is some ___________________ of the empirical formula, ___________ Benzene has a molar mass of 78 g/mol. Find n to find the molecular formula.

Slide 39 Overview of Calculations