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1 mole

1 mole. 1 mole marbles = covers Earth to depth of 50 miles. Counting by weighing. Purpose : Calculate the amount of pennies in the bag by weighing them. ( You may take the pennies out of the bag to weigh them, but do not count them) Data Calculations

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1 mole

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  1. 1 mole • 1 mole marbles = covers Earth to depth of 50 miles

  2. Counting by weighing Purpose: Calculate the amount of pennies in the bag by weighing them. ( You may take the pennies out of the bag to weigh them, but do not count them) Data Calculations • Atoms are too small to count so we weigh an amount (mole) and calculate the number

  3. Avogodro • The actual size of the molecule is not important • Avagadro concluded that equal volumes of gas have equal number of molecules • By the way, nobody really knew what a molecule was at the time

  4. Some words mean numbers • Pair • Dozen • Baker’s dozen • Gross • Ream

  5. Mole 1 mole = 6.02 x 1023 representative particles (atoms) (molecules) (formula units) ( the number of carbon-12 atoms in 12.00 g)

  6. Atomic mass units • We measure atoms in atomic mass units • Atomic mass unit = AMU = 1.66x10-27 Kg • It is defined as 12 for Carbon – 12 • Carbon 12 has six protons and six neutrons 1 proton=1 amu (1.0078amu) 1 neutron= 1 amu (1.0087 amu)

  7. Chemical Measurements • Atomic Mass The weighted average of all the mass numbers for all the isotopes of the atom (a.m.u.) • Formula Mass The sum of all the atomic masses for all atoms in the compound. (a.m.u.)

  8. Calculate the atomic mass or formula masses • Na • Cl • Br2 • NaCl • CO2 • Mg(OH)2 22.99 amu 35.45 amu 159.80 amu 58.45 amu 44.01 amu 58.33 amu

  9. Molar Mass = The mass of 1 mole ( 6.02 x 1023) of representative particles = The atomic mass in g = The formula mass in g

  10. Atom, Molecule and Ions • Representative particles – smallest particle of that substance SubstanceRepresentative Particle element atom molcular compound molecule ionic compound formula unit

  11. Calculate the Molar mass • Ca • H2 • KNO3 • (NH4)2S

  12. Remember 1 mole = 6.02 x 1023 r.p. 1 mole = molar mass(g)

  13. practice • Determine the number of FU in .866 moles of AgNO3? • Find the mass of 0.98 moles of CaCl2.

  14. Molar Volume The volume of 1mole of a gas at standard temperature – 0oC standard pressure – 1 atmosphere (pressure at sea level) 1 mole gas at STP = 22.4 L

  15. Remember 1 mole = 6.02 x 1023 r.p. 1 mole = molar mass(g) 1 mole = 22.4 L gas

  16. Multistep problems • How many fu in 18.9 g NaCl? • How many L in 4.5 x 1013molecules Ne?

  17. Solution= solute + solvent • % mass = g solute g of solution • PPM = parts of solute 1,000,000 parts of solution • Molarity = Moles solute • L of solution

  18. Calculate • 1. What is the molarity if 13.5 g NaCl in 451 ml solution? • 2. How many g of KOH are in 3.5 L of a .67 M solution? • 3. How many ml of solution would you need to have 17.5 g of NaOH in a .35 M solution?

  19. Percent Composition % element = g element x 100% g total compound

  20. Formulas • Empirical simplest whole number ratio • Molecular Formula actual number of atoms in the formula

  21. Practice: A compound has 13.5 g Ca 10.8 g O .675 g H What is its empirical formula?

  22. Try this one • What is the empirical formula that is • 25.9% N , 74.1% O

  23. Spark • 1. How many moles of HCl are there in 1.00 liters of 2.0 molar HCl? • 2. How many milliliters of 13.2 g of HCL of a 3.4 M solution? • 3. How many oxygen atoms are in 150.2 mL of 2.00 M H2SO4? • 4. How many sulfur atoms are in 158.2 g of aluminum thiosulfate , Al2 (S2O3)3?

  24. Molecular formula Set up table:

  25. A Chemical Reaction • Balance the following reaction • Hydrogen and oxygen react to produce water H2 + O2 H2O 2H2 + O2 2H2O

  26. Why Do We Balance an Equation / Reaction? • Makes it look pretty • I do what I’m told • Has something to do with conservation • Of mass

  27. Gas Has Mass • The balanced equation says I should • mix two hydrogens with one oxygen • And then the reaction should go well • Let’s See!

  28. So if……. • If the space between molecules is so large • and volume of gas at equal pressure • and equal temperature have the same number of molecules • How can we compare things that are not gases?

  29. Mass • We would weigh it. Of course. • But reactions go by numbers of molecules and atoms • Not by mass, which is in grams • I wonder if there is a conversion factor?

  30. How Can We Measure Atoms, the Darn Things are Too Small? • We can compare the mass of atoms • This was done back in the 1800’s • Hydrogen was found to be the smallest element • Everything was “relative” to Hydrogen • So we can compare elements by using hydrogen as a standard

  31. Consider This About Atoms • Each atom of hydrogen has 1.0 amu and uranium has 238 amu. • Which has more atoms, one gram of H or one gram of U? • Calcium has 40. amu per atom and helium has 4 amu per atom. • Approximately how many times more atoms are in one gram of helium than in one gram of calcium?

  32. Wow • So if we can compare masses of atoms to each other • Then we could weigh them to get the proper proportion of them in a reaction • Wouldn’t that be convenient? • If only there were a conversion factor!

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