1 / 21

Chapter 12: Chemical Quantities

Chapter 12: Chemical Quantities. Section 12.1: Counting Particles of Matter. The macroscopic vs. the submicroscopic. Is it practical to count out 200 grains of rice, or 200 pieces of macaroni? No Way!

michaelkthompson
Télécharger la présentation

Chapter 12: Chemical Quantities

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. Chapter 12: Chemical Quantities Section 12.1: Counting Particles of Matter

  2. The macroscopic vs. the submicroscopic Is it practical to count out 200 grains of rice, or 200 pieces of macaroni? No Way! We sell these items by the pound. How would we be able to calculate how many items are in a container without counting?

  3. Is there a way to know the number without counting? There is! At your table, discuss a method to figure out how many items are in each bag, without counting. What information would you need?

  4. Counting by Weighing How many candy hearts are in the jar? What you need; the data: Mass of the empty jar (with stopper) = 209.624 g Mass of the full jar (with stopper)= 396.227 g Average mass of one candy heart= 1.92 g

  5. Counting by Weighing (cont) How many atoms are in 1.0079 g of Hydrogen if the average mass of one atom is 1.67425 x 10-24 g? What you need: • Total mass • Mass of one atom What number do you get? 6.02 x 1023 atoms of H

  6. STOICHIOMETRY The study of the relationships between the macroscopic quantities like mass and volume, and the submicroscopic quantities like number and mass of the atoms involved in a chemical reaction. Chemists cannot count individual items because the particles of matter are too small and too numerous- so we use this process.

  7. What is a mole? The group or unit of measure that we use in chemistry to count numbers of atoms, ions, molecules or formulas units of a substance (elements or compounds) Mole Map • http://www.youtube.com/watch?v=4QiGVqK7ASU • https://www.youtube.com/watch?v=TEl4jeETVmg&list=PLKEmXepzBsY9Zx4EHN27HHSQyvb4Xysb4&index=6

  8. What is Avogadro’s Number? 6.02 x 1023 • This HUGE quantity represents the number of particles in one mole of ANY substance. It is just a number! • A mole can be analogous to a dozen, a pair, a century…you know how many individual parts there are in each one of these measurements, right? 12, 2, 100…a mole is just A LOT more. • Video

  9. What is Molar Mass? The mass in grams of one mole of an element or a compound! (g/mol) For elements, it is the atomic mass in grams • One atom of magnesium = 24.31 u (atomic mass units ) • One mole of magnesium = 24.31 g/mol For compounds, it is the sum of all the masses of each element in the compound’s formula One mole MgSO4 = 24.31 + 32.07 + 4(16.0) = 120.4 g/mol

  10. So what about molecular mass and formula mass? These are not hard to calculate, because essentially the value of each is the same. It is the units that are different! • Formula Mass is a term used for IONIC compounds • NaCl has a FORMULA mass of 58.45 u and a MOLAR mass of 58.45 g/mol • ~ Same quantity, different unit Ex. CaCl2 Ca: 1 x 40.1 = 40.1 Cl: 2 x 35.5 = 71 Formula mass = 111.1 u Molar mass = 111.1 g/ mol

  11. So what about molecular mass and formula mass? (Cont) • Molecular Mass is a term used for COVALENT compounds • C2H6 has a MOLECULAR mass of 30.1 u and a MOLAR mass of 30.1 g/mol Ex: C2H6O C: 2 x 12 = 24 H: 6 x 1 = 6 O: 1 x 16 = 16 Molecular mass = 46 u Molar mass = 46 g/ mol

  12. When solving problems like the following, you must use DIMENSIONAL ANALYSIS (Factor Label Method). • Steps grams ↔ moles ↔ “stuff” Molar mass Avogadros #

  13. Practice Problems (pg 410) 68 g NH3 = ? mol 1 N = 1 x 14.007 = 14.007 g 3 H = 3 x 1.008 = 3.024 g 14.007 + 3.024= 17.031 g NH3 68 g NH3 x 1 mole NH3 = 3.99 mol NH3 17.031 g NH3

  14. Practice Problems (cont) 17.5 g CuO = ? mol 1 Cu = 1 x 63.546 = 63.546 g 1 O = 1 x 15.999 = 15.999 g 63.546 + 15.999 = 79.545 g CuO 17.5 g CuO x 1 mole CuO= 0.22 molCuO 79.545 g CuO

  15. Practice Problems (cont) 0.550 mol F2 = ? g 2 F = 2 x 18.998 = 37.996 g 0.550 mol F2 x 37.996 g F2= 20.90 g F2 1 mole F2

  16. Practice Problems (cont) 2.35 mol BaI2 = ? g 1 Ba = 1 x 137.33 = 137.33 g 2 I = 2 x 126.90 = 253.8 g 137.33 + 253.8 = 391.13 g 2.35 mol BaI2 x 391.13 g BaI2 = 919.16 g BaI2 1 mole BaI2

  17. Practice Problems (cont) 10.7 g Li = ? atoms 10.7 g Li x 1 mole Li X 6.02 x 1023 atoms Li 6.941 g Li 1 mole Li = 9.28 x 1023 atoms Li

  18. Practice Problems (cont) 144.6 g W = ? atoms 144.6 g W x 1 mole W X 6.02 x 1023 atoms W 183.85 g W 1 mole W = 4.73 x 1023 atoms W

  19. Practice Problems (cont) 1. Does 50.0 g S or 50.0 g Sn represent a greater # of atoms? 50.0 g S x 1 mole S X 6.02 x 1023 atoms S 32.066 g S 1 mole S = 9.39 x 1023 atoms S 50.0 g Snx 1 mole SnX 6.02 x 1023 atoms Sn 118.710 g Sn1 mole Sn = 2.54 x 1023 atoms Sn • Sulfur is smaller- has a greater number of atoms

  20. What mass of water must be weighed to obtain 7.50 mol of H2O • How many moles are in 17.5g copper (II) oxide • How many molecules are in each sample? • 89g nitrogen monoxide • 10.8g boron trifluoride • What is the molecular mass of aspirin (C9H8O4)

More Related