1 / 35

Unit 4 – The Mole

Unit 4 – The Mole. Honors Chemistry Part 1. WHAT IS A MOLE?. 602214199000000000000000 6.02 x 10 23. Mole Facts. 6.02 X 10 23  Pennies: Would make at least 7 stacks that would reach the moon. 6.02 X 10 23  Watermelon Seeds: Would be found inside a melon slightly larger than the moon.

haracha
Download Presentation

Unit 4 – The Mole

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  1. Unit 4 – The Mole Honors Chemistry Part 1

  2. WHAT IS A MOLE? • 602214199000000000000000 • 6.02 x 1023

  3. Mole Facts • 6.02 X 1023  Pennies: Would make at least 7 stacks that would reach the moon. • 6.02 X 1023 Watermelon Seeds: Would be found inside a melon slightly larger than the moon. • 6.02 X 1023  Blood Cells: Would be more than the total number of blood cells found in every human on earth. • 1 Liter bottle of Water contains 55.5 moles H20

  4. Definition of Mole • The amount of atoms in 12.0 grams of Carbon 12 (6.02 x 1023 atoms known as Avogadro’s number). • A sample of any element with a mass equal to that element's atomic weight (in grams) will contain precisely one mole of atoms (6.02 x 1023 atoms).

  5. Molecular Mass • The sum of the masses of all the atoms in a molecule of a substance CaCO3 1 atom of Ca = 40.08 amu 1 atom of C = 12.00 amu 3 atoms of O = 48.00 amu 100.08 amu } Add these

  6. Formula Mass (Molar Mass) • The mass of 1 mole (in grams) • Equal to average atomic mass but the unit is grams 1 mole of C atoms = 12.01 g 1 mole of Na atoms = 22.99 g 1 mole of Cu atoms = 63.55 g

  7. Example Problem Find the mass of 1 mole of KAl(SO4)2● 12H2O 1 K = 39.10 1 Al = 26.98 2 (SO4) 2(32.06 + ((16.00 x 4))=192.12 12 H2O 12(2.02 + 16.00) =216.24 Mass of 1 mole = 474.44 g/mol

  8. Try These: • Find the molecular mass for these : • HNO3 • CO2 • Find the molar mass for these compounds: • C6H10O5 • H2SO4

  9. The Mole • 1 mole of gas always contains 6.02 x 1023 molecules of that gas • 1 mole Cl2 gas = 6.02 x 1023 molecules of Cl2 • 1 mole NO2 gas = 6.02 x 1023molecules of NO2 • 1 mole CO gas = 6.02 x 1023 molecules of CO • 1 mole CO2 gas = 6.02 x 1023molecules of CO2

  10. The Mole • Also applies to other particles! (not only molecules in a gas) • 1 mole C = 6.02 x 1023 C atoms • 1 mole H2O = 6.02 x 1023H2O molecules • 1 mole NaCl = 6.02 x 1023NaCl formula units • 1 mole of Na+ = 6.02 x 1023 Na+ ions • 1 mole of Cl- = 6.02 x 1023Cl– ions

  11. Avogadro’s Number • We can use Avogadro’s # as a conversion factor: 1 mole 6.02 x 1023 particles Or 6.02 x 1023 particles 1 mole • Note that a particle could be an atom, molecule, formula unit, or ion !

  12. Example Problems • How many molecules are in 3.5 moles of H2O? • How many moles are present in 4.65 molecules of NO2?

  13. Mass and Mole Relationship • 1 mole of any substance = the molar mass of that substance (in grams) • Find the number of moles present in 56.7 g of HNO3. 56.7 g HNO3 1 mole HNO3 63.01 g HNO3

  14. Example Problems • Find the number of grams present in 4.5 moles of C6H10O5. • Find the number of moles present in 12.31 g of H2SO4. • How many molecules are in 4.5 grams of NaCl?

  15. Gas Volumes and Molar Mass • Avogadro’s Law • Equal volumes of gases under the same conditions of temperature and pressure contain equal numbers of molecules • 1 mole of any gas at standard temperature and pressure (STP) occupies 22.4 liters • Standard temperature: 0ºC or 273K • Standard pressure: 1 atm or 101.325 kPa

  16. Gas Volumes and Molar Mass • 32.00 g O2 = 1 mole = 22.4 L • 2.02 g H2 = 1 mole = 22.4 L • 44.01 g CO2 = 1 mole = 22.4 L

  17. ExampleProblems • How many liters are present in 5.9 moles of O2? • How many liters are present in 3.67 moles of CO2? • How many atoms of O are present in 78.1 g of O2?

  18. Percent Composition • Finding what percent of the total weight of a compound is made up of a particular element • Formula for calculating % composition: Total mass of the element in the compound Total formula mass X 100

  19. Example Problem: • Calculate the % composition of calcium in Ca(OH)2.

  20. Example Problems • Find the percentage composition of a compound that contains 1.45 g of carbon, 4.23 g of sulfur, and 1.00 g of hydrogen in a 6.68 g sample. • 21.7% carbon • 63.3% sulfur • 15.0% hydrogen

  21. Example Problem • A sample of an unknown compound with a mass of 5.00 grams is made up of 75% carbon and 25% hydrogen. What is the mass of each element? • 3.75 g of carbon • 1.25 g of hydrogen

  22. Formulas • Empirical Formula - expresses the smallest whole number ratio of atoms present • E.g. CH2O • Ionic formulas are always empirical formulas • Molecular Formula - states the actual number of each kind of atom found in one moleculeof the compound. • E.g. C6H12O6

  23. Empirical Formula • Determine mass in grams of each element • Calculate the number of molesof each • Divide each by the smallest number of moles to obtain the simplest whole number ratio • If whole numbers are not obtainedin step 3, multiply all by the smallest number that will give whole numbers

  24. Empirical Formula Remember this: • Percent to mass • Mass to mole • Divide by small • Multiply ‘till whole

  25. Example Problem • Given that a compound is composed of 60.0% Mg and 40.0% O, find the empirical formula.

  26. Example Problem • A compound is found to contain 68.5% carbon, 8.63% hydrogen, and 22.8% oxygen. The molecular weight of this compound is known to be approximately 140.00 g/mol. Find the empirical and molecular formulas.

  27. Hydrates • Ionic compounds • Water is bonded to the crystal structure • Ex: CuSO3 • 7H2O • The percentage of water in a hydrate can easily be calculated using the formula: % Water = Mass of water x 100 Mass of hydrate

  28. Example Problem • What is the percentage of water in CuSO3 • 7H2O? • A 3.5 g sample of a hydrate is heated and only 1.7 g of the anhydrous salt remain. What is the percentage of water?

  29. Law of Definite Proportions • Formulas give the numbers of atoms or moles of each element • Always a whole number ratio • 1 molecule NO2 : 2 atoms of O for every 1 atom of N • 1 mole of NO2 : 2 moles of O atoms to every 1 mole of N atoms

  30. Law of Multiple Proportions • When any two elements, A and B, combine to form more than one compound, the different masses of B that unite with a fixed mass of A have a small whole-number ratio • Example: • In H2O, the proportion of H:O = 2:16 or 1:8 • In H2O2, H:O is 2:32 or 1:16

  31. How Do We Determine Concentration? • Molarity • Molality

  32. How do we make solutions? M1 = m1/V1 rearrange to M1V1 = m1 M2 = m2/V2 rearrange to M2V2 = m2 If m1=m2 then, M1V1 = M2V2

  33. M1V1 = M2V2 • M1 = concentration of the first solutionV1 = volume of the first solutionM2 = concentration of the second solutionV2 = volume of the second solution • Let's consider a sample problem: • You have 1 L of a 0.125 M aqueous solution of table sugar.  You want to dilute the solution to 0.05 M.  What do you do?

  34. Dilution • To solve the problem, you simply plug in the three numbers you know: • (0.125 M) (1 L) = (0.05 M) V2 • 2.5 L = V2 • Using the equation, you determine that the volume of the diluted solution should be 2.5 L.  • So we simply add enough water to the first solution so that the solution's volume becomes 2.5 L.

  35. What is Saturation? • A solution is saturated if it contains as much solute as can possibly be dissolved under the existing conditions of temperature and pressure • Unsaturated: Has less than maximum amount of solute that can be dissolved • Supersaturated: Contains more than maximum (How can this happen?)

More Related