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Chemical Quantities

Chemical Quantities. Or How I Learned to Love the Mole. A mole?. Not the type of mole we are talking about. Vocabulary. Mole Avogadro’s Number Representative Particle Molar Mass. How do you measure the amount of something?.

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Chemical Quantities

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  1. Chemical Quantities Or How I Learned to Love the Mole

  2. A mole? Not the type of mole we are talking about

  3. Vocabulary • Mole • Avogadro’s Number • Representative Particle • Molar Mass

  4. How do you measure the amount of something? • Look at the dozen things that you brought to class. Think about how you could measure the amount of matter in that stuff. • Brain storm 3 ways and compare your ways to your neighbor. • What ways did you come up with?

  5. Count it! • What are the advantages of this method? • What if your stuff is differently sized? • What if something is too small to count?

  6. Weight it! (take the mass) • Standardized method of measuring amount of matter! • What if it is too small to weigh? • Too big?

  7. Size it! (take the volume) • Volume is how much space it takes up. • Advantages of this system? • Disadvantages?

  8. Counting Matter What are some everyday ways we count matter? 1. 2. 3. DOZEN = 12 things 1 GROSS = 144 things What about molecules? Or atoms? BUSHEL of corn = 21.772 kg 1 MOLE = ???

  9. Problem! • Atoms are too small to count • Too small to put on a scale • Too small to take a volume • They vary in size • Small variations in size add up to big differences! • Can we figure out how many atoms are in some stuff?

  10. Solution! • Come up with a way to count atoms by taking the mass of the stuff! • Let’s try it with your dozen items…

  11. The Mole • It represents a counted number of things. • IN Chemistry the term MOLE represents the number of particles in a substance. In Chemistry is NOT this furry little animal or the spot on your face…

  12. Procedure • Find the mass of your dozen items together • Divide this total mass by 12 to find the average mass of one of these items. • Weigh one of the items individually and compare this number to the number you got in step 2. • Now estimate the mass of 8 of these items using the number from step 2. • Weight 8 of your items. How close was this? • Repeat this process with 5 items. • Now estimate the mass of your stuff if you had 6.02x10^23 items. This is a mole! • Let’s calculate the mass of a mole of pennies.

  13. How much is a mole? A mole of pennies is very large.

  14. Just how many is a mole? • One mole represents 6.02 x 1023 of things (units, molecules, compounds, formula units). This is called Avogadro’s number. • One mole of most elements contains 6.02 x 1023atoms. • 1 mole O2 = 6.02x1023molecules of O2 • 602 000 000 000 000 000 000 000

  15. Just how big is a mole? • Listen to The Mole Song! • An Avogadro's number of standard soft drink cans would cover the surface of the earth to a depth of over 200 miles • If you had Avogadro's number of un-popped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles. • If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.

  16. Counting stuff… • Count it! If it is something large like the number of beans in a bag, you just count them. • If, however, it is something exceedingly small like an atom or ion, how do you count it? • Counting individual particles is not possible, but you can “count” particles if you introduce a term that represents a specific number of particles. • A “dozen” represents 12 of that item. • and a mole represents 6.02 x 1023 representative particles of a substance.

  17. Solving the Problems Samples Required: dimensional analysis/factor label • How many molecules are in 3.00 moles of N2? • How many moles of Na are in 1.10 x 1023 atoms?

  18. Practice moles to particles • Determine the number of atoms in 2.50 mol Zn. • Given 3.25 mol AgNO3, determine the number of formula units. • Calculate the number of molecules in 11.5 mol H2O.

  19. Practice Particles to Moles • How many moles contain each of the following? • 5.75 x 1024 atoms Al • 3.75 x 1024 molecules CO2 • 3.58 x 1023 formula units ZnCl2 • 2.50 x 1020 atoms Fe

  20. 10-2 Mass & the Mole 10-3 Moles of Compounds

  21. Molar Mass Defn: is the mass (think grams) of one mole of a substance • Atomic masses (from periodic table) represent molar mass. • Units g/mol • 1 mole of Carbon has 6.02 x 1023 atoms of C and they have a mass of 12.01 grams. • To calculate the molar mass of a compound, you add up the molar masses of all the elements in that compound

  22. Molar Mass Practice • What is the mass of 1.00 mole of Oxygen? Of Nitrogen? • Find the molar mass for: • SO3 • Na2SO4 1 mole O = 16.0 grams 1 mole N = 14.0 grams SO3= 80 g/mole 1 Mole = 142.043g Tutorial Site Molar Mass Calculator for homework help

  23. Molar Mass Practice • When you see 1.00 mole = _?_ g, think “g means GO to the PERIODIC TABLE” to find the molar mass. http://www.webelements.com/

  24. Practice Problems • Determine the molar mass of each of the following ionic compounds: • NaOH • CaCl2 • KC2H3O2 • HCN • CCl4 • H2O

  25. Grams-Mole Conversions • How many moles are in 56.8 g of HCl? • How many grams are in .05 moles Na2SO4?

  26. Practice Mole to Grams • Determine the mass in grams of each of the following. • 3.57 mol Al • 42.6 mol Si • 3.45 mol Co • 2.45 mol Zn

  27. Practice Gram to Mole • How many atoms are in each of the following samples? • 55.2 g Li • 0.230 g Pb • 11.5 g Hg • 45.6 g Si • 0.120 kg Ti

  28. Ex: Mass to Particle Conversion • Gold is one of a group of metals called the coinage metals (copper, silver and gold). How many atoms of gold (Au) are in a pure gold nugget having a mass of 25.0 g? Known:Unknown: Mass = 25.0 g Au Molar mass Au = 196.97 g/mol Au 25.0 g Au x 1 mole Au x 6.02 x 1023 atoms Au = 7.65 x 1022 atoms Au 196.97 g Au 1 mol Au Number of atoms = ? Atoms Au

  29. Practice Problems • How many atoms are in each of the following samples? • 55.2 g Li • 0.230 g Pb • 11.5 g Hg • 45.6 g Si • 0.120 kg Ti

  30. Ex: Particle to Mass Conversion • A party balloon contains 5.50 x 1022 atoms of helium (He) gas. What is the mass in grams of the helium? Known:Unknown: Number of atoms = 5.50 x 1022 atoms He Molar mass He = 4.00 g/mol He 5.50 x 1022 atoms He x 1 mol He x 4.00 g He = 0.366 g He 6.02 x 1023 atoms He 1 mol He Mass = ? G He

  31. Practice Problems • What is the mass in grams of each of the following? • 6.02 x 1024 atoms Bi • 1.00 x 1024 atoms Mn • 3.40 x 1022 atoms He • 1.50 x 1015 atoms N • 1.50 x 1015 atoms U

  32. Molar Volume • The volume of a gas is usually measured at standard temperature and pressure (_STP_) • Standard temp = ___0°_ C • Standard pressure = ___1___ atmosphere (atm) • 1 mole of any gas occupies __22.4__ L of space at STP

  33. Molar Volume Practice • How many moles would 45.0 L of He gas be? • How many liters of O2 would 3.8 moles occupy?

  34. Putting it all together • 1.0 mole = _6.02 x 1023___atoms or molecules • 1.0 mole = _?__ g(PT) • 1.0 mole = _22.4 L (at STP)

  35. Helpful Chart! MOLES Grams Volume in Liters Atoms or Molecules

  36. Chemical Formulas and the Mole • Chemical formula for a compound indicates the types of atoms and the number of each contained in one unit. • Ex. CCl2F2 - Freon • Ratio of carbon to chlorine to fluorine is 1:2:2 • Ratios can be written: or • In one mole of freon you would have 1 mole of carbon, 2 moles of chlorine and 2 moles of fluorine.

  37. Ex: Mole Relationship from Chemical Formulas • Determine the moles of aluminum ions (Al3+) in 1.25 moles of aluminum oxide. 1.25 molAl2O3 x 2 mol Al3+ ions = 2.50 mol Al3+ ions 1 mole Al2O3

  38. Practice Problems • Determine the number of moles of chloride ions in 2.53 mol ZnCl2. • Calculate the number of moles of each element in 1.25 mol glucose (C6H12O6). • Determine the number of moles of sulfate ions present in 3.00 mol iron (III) sulfate (Fe(SO4)3). • How many moles of oxygen atoms are present in 5.00 mol diphosphoruspentoxide? • Calculate the number of moles of hydrogen atoms in 11.5 mol water.

  39. Practice Problems • A sample of silver chromate has a mass of 25.8 g. • How many Ag+ ions are present? • How many CrO42- ions are present? • What is the mass in grams of one unit of silver chromate? • What mass of sodium chloride contains 4.59 x 1024 units? • A sample of ethanol (C2H5OH) has a mass of 45.6 g. • How many carbon atoms does the sample contain? • How many hydrogen atoms are present? • How many oxygen atoms are present?

  40. Warm-up review You have 10.0 grams of ordinary sugar on a scale. Sugar is sucrose and has a formula of C12H22O11. • How many moles of sugar are on your scale? • How many molecules of sucrose are on your scale? • How many moles of carbon atoms are on your scale? • How many atoms of hydrogen are on your scale? • How many atoms of oxygen are on your scale? • What is the ratio of oxygen atoms to hydrogen atoms? • What is the ratio of carbon atoms to sugar molecules?

  41. Warm-up Review #2 • You have 30.0 g of barium hydroxide on a scale. • What is the chemical formula of this substance? • What is the molar mass? • How many moles do you have of the compound? • How many moles do you have of barium? • How many moles do you have of oxygen? • How many moles of (OH)- ions do you have? • How many ions of (OH)- do you have? • What is the ratio of (OH)- ions to Ba2+ ions?

  42. 10-3 Empirical & Molecular Formulas

  43. Percent Composition • the percentage by massof each element in a compound • The percent comp. is found by using the following formula:

  44. Percent Composition Example Ex. Compound XY is 55g element X and 45g element Y 55 g of element Xx 100 = 55 % element X 100 g of compound 45 g of element Y x 100 = 45 % element Y 100 g of compound

  45. Percent Composition from Chemical Formula • First find the molar mass of each element and the molar mass of the compound • Ex: what is the % composition of H in 1 mole of H2O? • Multiply the molar mass of the element by its subscript in the formula. • 1.01 g/mol H x 2 mol = 2.02 g H % by mass H = 2.02 g x 100 = 11.2% H 18.02 g H2O

  46. Percent Composition from Chemical Formula Example continued – • Molar mass of O for each mole of H2O? • 16.00 g/mol O x 1mol O = 16.00 g O 16.00 g x 100 = 88.8 % O 18.02 g

  47. % Composition Practice • What is the percent of C & H in C2H6? Hint… assume 1 mole! • What is the percent of each element in Na2SO3?

  48. Empirical Formulas • This is the LOWEST whole number ratio of the elements in a compound. For example, the empirical formula for • Molecular Formula C6H6 • Empirical Formula CH • What is the empirical formula for each? • C2H6 • C6H12O6

  49. 10-3: Calculating Empirical Formula • Steps for calculating Empirical Formula give mass or percent composition: • If given a percent sign (%), remove the sign & change to GRAMS. • You are assuming you have 100 g of the compound. • Convert grams ---> moles. 1. Select lowest number of moles 2. Divide each number of moles by this number. 3. If the number divides out evenly, these are the subscripts of the elements in the compound. 4. If any of the numbers have a .5, MULTIPLY them ALL by TWO & then place these numbers as the subscripts.

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