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The MOLE

The MOLE. It’s a beauty mark… It’s a small furry garden pest… No, wait… its how we count ATOMS!. Molar Mass. The mass of 1 mole of a substance. Calculate the molar mass . Calculate the molar mass of 1 mole of magnesium chloride. (first you need a formula). FORMULAS - review. MgCl 2

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The MOLE

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  1. The MOLE It’s a beauty mark… It’s a small furry garden pest… No, wait… its how we count ATOMS!

  2. Molar Mass The mass of 1 moleof a substance.

  3. Calculate the molar mass • Calculate the molar mass of 1 mole of magnesium chloride. (first you need a formula)

  4. FORMULAS - review MgCl2 • The subscript is the number at the bottom of a formula. • There is 1- Mg & 2 - Cl

  5. How to calculate molar mass • Identify the # of atoms of each element • Multiply # atoms by the atomic mass of that element. (round to 2 #’s after the decimal) • Add them all together • Grams (g) is the unit

  6. Molar Mass MgCl2 Mg – 1 (24.31) = 24.31 Cl – 2 (35.45) = 70.90 or 95.21 g/mol 1 mole MgCl2 = 95.21 g MgCl2

  7. Molar Mass • Molar mass allows scientist a way to measure moles in a lab. • Molar mass of Fe = • Molar mass of O2= • Molar mass of Cu(OH)2=

  8. Calculate the molar mass: (you MUST write the formula correctly before answering) • 1 mole of Ammonium phosphide • 1 mole of Trinitrogenpentachloride

  9. Percent Composition • The percent BY MASS of each element in a compound – divide the element’s total mass (part) by the molar mass (whole) then multiple by 100 to get the percent. • Ex: % composition of MgCl2 Mg – 1 (24.31) = 24.31 / 95.21 x 100 = 25.53% Mg Cl– 2 (35.45) = 70.90 / 95.21 x 100 = 74.47% Cl Molar mass = 95.21 g/mol (PART) (WHOLE) (PART) (WHOLE)

  10. Practice - Calculate the % comp of KMnO4: K – 1 (39.10) = 39.10 / 158.04 x 100 = 24.74% K Mn – 1 (54.94) = 54.94 / 158.04 x 100 = 34.76% Mn O – 4 (16.00) = 64.00 / 158.04 x 100 = 40.50% O molar mass KMnO4 = 158.04

  11. Calculating the amount of an element in a sample • Find the % comp of the element in the compound • Change the % to a decimal (move decimal 2 times to the left or divide by 100) • Multiply that decimal by the amount (g) of the sample. • Ex: Calculate amount of chlorine in 203.5 grams of MgCl2. (use the % we found earlier) • 74.47% Cl = .7447 x 203.5 = 151.5 grams Cl

  12. Practice: Calculate the amount of oxygen in 15.75 grams of water. • H – 2(1.01) = 2.02 • O – 1(16.00) = 16.00 • Molar mass H2O = 18.02 / 18.02 x100 = 88.79% O 88.79% O = .8879 x 15.75 = 13.98 grams O Complete % comp worksheet

  13. Empirical Formula • The lowest whole number ratio (subscripts) of elements in a compound. • Cannot be reduced!!! not empirical empirical • Ex: C6H12O6 CH2O

  14. Molecular Formula • Actual number of atoms in a chemical compound molecular • EX: C12H24O12 • Molecular Formulas can be reduced to Empirical Formulas molecular empirical • EX: C12H24O12 CH2O • Different molecular formulas can have similar empirical formulas • molecular empirical • EX: N3O9  N12O36  Molecular formula: C76H52O46 Empirical formula: ___________ NO3

  15. PRACTICE:1. Identify each as empirical (can’t be reduced) or molecular (can be reduced) 2. If its molecular – write the empirical • C2H4 • NO3 • S9Cl12 • C3Cl9 • N4S9 CH2- empirical molecular empirical S3Cl4 - empirical molecular CCl3 - empirical molecular empirical

  16. Finding Empirical Formula from Percent Composition • Ex: A compound was found to be 54.53% Carbon, 9.15% Hydrogen, and 36.32% Oxygen. Find its Empirical Formula. Steps: • Assume a 100g sample (change %  g) • Use molar mass to find moles of each • Divide all moles by the smallest number of moles • Round each to the nearest whole # (sometimes you have to multiply to get a whole number - special) • The resulting whole #are the subscripts for that element in the empirical formula

  17. Calculating Empirical Formula • 63.5% Silver 8.2% Nitrogen 28.3% Oxygen • 63.5 g Ag 8.2 g N 28.3 g O • 107.87 14.01 16.00 • .589 mole Ag .59 mole N 1.77 mole O • .589 .589 .589 • 1 1 3 • AgNO3

  18. Calculating Empirical Formula (special) • 60.00%C 4.48%H 35.53%O • 60.00g C 4.48g H 35.53g O • 12.01 1.01 16.00 • 4.996 mole C 4.44 mole H 2.221 mole N • 2.221 2.221 2.221 • 2.249 2 1 x4 x4 x4 • 9 8 4 • C9H8O4

  19. Calculating Molecular Formula • Find the empirical formula • Calculate the molar mass of your empirical formula • Identify the molar mass of your molecular (GIVEN in the problem everytime!) • Divide the molecular mass / empirical mass • Round to the nearest whole # • Multiply the whole # by the subscripts in the Empirical formula

  20. Practice • If a compound has an empirical formula of NO3 and a molecular mass of 186g – what is the molecular formula? • Empirical formula: NO3 molar mass: 62.01g • Molecular mass (given) 186g • empirical mass 62.01 • 3 x NO3 = N3O9

  21. What is a mole? • For counting matter (quantity) in chemistry we use the mole • Moles are used to correctly measure chemicals for reactions in a lab. • Problem – no way to “physically” measure a mole. (no lab equipment measures moles) • Solution – molar mass (balance)

  22. Using Molar Mass • Remember – Molar mass is the mass (grams) of 1 mole • 1 mole Fe = _________grams Fe • 2.5 mole Fe = ________ grams Fe • 113.5 grams Fe = _______ moles Fe • Mass to mole = divide by molar mass • Mole to mass = moletiply by mole mass 

  23. Using molar mass • How many grams are in 15.7 mole MgCl2 • How many moles are in 0.75 grams of silver? • What is the mass of 30.7 mole water?

  24. Using molar mass • In the lab, Mrs. Mathieson needs 2.57 moles of NaCl to do an experiment. How many grams would be needed to equal 2.57 moles of NaCl? • After doing the experiment, Mrs. Mathieson has 1.02 moles of NaCl remaining – how many grams does that equal?

  25. What is a mole? • For counting matter (quantity) in chemistry we use the mole • 1 mole = 6.02 x 1023 representative particles (particles are verytiny) • This value is called Avogadro’s Number

  26. What are REPRESENTATIVE PARTICLES? • ie. the smallest particle that retains chemical and physical properties • 3 types depending on the compound: • Atoms: Single element • Molecules: covalent compound • Formula Units: Ionic Compounds or ions

  27. Using Avagadro’s Number • Remember there are 6.02 x 1023 particles (atoms, molecule, f.u) in 1 mole • How many atoms are in 1 mole? • How many atoms are in 2.10 moles of Copper? • How many moles are in 4.21 x 1026 atoms of aluminum?

  28. What is a mole? • For counting matter (quantity) in chemistry we use the mole • 1 mole = 22.4 L of any gas • This value is called Molar Volume

  29. Using molar volume • How many liters are in 1 mole of any gas? • How many liters are in 12.95 L of oxygen gas? • How many moles are in 0.758 moles of nitrogen gas?

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