1 / 106

Chapter 3

Chapter 3. Atoms: The Building Blocks of Matter. Objectives. State the postulates of Dalton’s atomic theory. Define the terms atomic mass and molar mass. Define the terms mole and Avogadro’s number. Write the name for common elements, given the symbol, or the symbol, given the name.

deepak
Télécharger la présentation

Chapter 3

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 3 Atoms: The Building Blocks of Matter

  2. Objectives • State the postulates of Dalton’s atomic theory. • Define the terms atomic mass and molar mass. • Define the terms mole and Avogadro’s number. • Write the name for common elements, given the symbol, or the symbol, given the name. • Calculate the molar mass of an element or compound, given its formula. • Calculate the mass of an element or compound given the number of moles, or the number of moles of a given mass of an element or compound. • Calculate the number of atoms or molecules of an element or compound given the number of moles, or the number of moles given the number of atoms or molecules.

  3. Chapter 3 Section 1 The Atom

  4. The modern definition of an element is a substance that cannot be further broken down by ordinary chemical means – H, C, O Elements also combine to form compounds that have different physical and chemical properties than those of the elements that form them – H2O. The transformation of a substance into one or more new substances is a chemical reaction.

  5. Law of Conservation of Mass Mass can not be created or destroyed in ordinary (not nuclear) chemical reactions. All the mass can be accounted for. Mass at the start = mass at end

  6. Dalton’s Atomic Theory John Dalton - 1808 • All matter is composed of extremely small particles called atoms. • Atoms of a given element are identical in size, mass and other properties; atoms of different elements differ in size, mass and other properties.

  7. 3. Atoms cannot be divided, created or destroyed. • 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. • In chemical reactions, atoms are combined, separated or rearranged. • (Law of Conservation of Matter)

  8. Relative Atomic Masses Masses of atoms expressed in grams are very small and not useful. It is more convenient to use relative atomic masses. Therefore the standard used to govern units of atomic mass is the carbon-12 atom. It has been assigned a mass of exactly 12 atomic mass units, or 12 amu.

  9. The atomic mass of any other atom is determined by comparing it with the mass of the carbon-12 atom. Examples: The hydrogen atom has an atomic mass of 1/12 that of the carbon-12 atom or 1 amu. Oxygen has an atomic mass of 16/12 the mass of a carbon-12 atom or 16 amu.

  10. Relating Mass to Numbers of Atoms • Introduction of three very important concepts: • The mole • Avogadro’s number • Molar mass

  11. The Mole Mole – the amount of a substance that contains as many particles are there are atoms in exactly 12 g of carbon. One mole of carbon weighs 12 grams. The mole is the SI unit for amount of substance. It is a counting unit. Mole is related to the counting term dozen.

  12. What is a counting unit? You’re already familiar with one counting unit…a “dozen” A dozen = 12 “Dozen” 12 A dozen doughnuts 12 doughnuts A dozen books 12 books A dozen cars 12 cars A dozen people 12 people

  13. A Mole of Particles Contains 6.02 x 1023 particles Avogadro’s Number 1 mole C = 6.02 x 1023 C atoms 1 mole H2O = 6.02 x 1023 H2O molecules 1 mole NaCl= 6.02 x 1023NaCl molecules

  14. Avogadro’s Number – 6.02 x 1023 is the number of particles in exactly one mole of a pure substance. 1 mole of gold = 6.02 x 1023 particles 1 mole of uranium = 6.02 x 1023 particles 1 mole of water = 6.02 x 1023 particles

  15. Objectives • Define the terms atomic mass and molar mass. • Define the terms mole and Avogadro’s number. • Calculate the molar mass of an element or compound, given its formula. • Calculate the mass of an element or compound given the number of moles, or the number of moles of a given mass of an element or compound. • Calculate the number of atoms or molecules of an element or compound given the number of moles, or the number of moles given the number of atoms or molecules.

  16. What does a “mole” count in? A mole = 6.02  1023 (called Avogadro’s number) 6.02  1023 = 602,000,000,000,000,000,000,000 “mole” 6.02  1023 1 mole of doughnuts 6.02  1023 doughnuts 1 mole of atoms 6.02  1023 atoms 1 mole of molecules 6.02  1023 molecules

  17. How big is a mole? Enough soft drink cans to cover the surface of the earth to a depth of over 200 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.

  18. Molar Mass Molar Mass – The mass (in grams) of one mole of a pure substance. Molar masses are written in units g/mol. The molar mass of an element is equal to the atomic mass of the element. Look on the periodic table for the atomic masses.

  19. Other terms commonly used for the same meaning of molar mass: Molecular Weight Molecular Mass Formula Weight Formula Mass

  20. Molar Mass Examples: Molar mass of lithium (Li) = 6.94 g/mol Molar mass of helium (He) = 4.00 g/mol Molar mass of mercury (Hg) = 200.6 g/mol

  21. Molar Mass A molar mass of an element contains one mole of atoms. 4.00 g helium = 1mole = 6.02 x 1023 atoms. 6.94 g lithium = 1mole = 6.02 x 1023 atoms. 200.6 g mercury = 1mole = 6.02 x 1023 atoms.

  22. One mole of carbon (12 grams) and one mole of copper (63.5 grams) Both contain 6.02 x 1023 atoms

  23. Molar Mass A molar mass of a compound is the sum of the molar masses of the elements. Example: Water, H2O: 2 H = 2 x 1g/mole = 2g/mole 1 O = 1 x 16g/mole = 16g/mole molar mass of H20 =18g/mol

  24. Molar Mass A molar mass of a compound is the sum of the molar masses of the elements. Example: methane, CH4: 4 H = 4 x 1g = 4g/mole 1 C = 1 x 12g = 12g/mole molar mass of CH4 =16g/mol

  25. Gram/Mole Conversions Molar masses can be used as a conversion factor in chemical calculations. Example: The molar mass of helium is 4.00 g/mol. To calculate how many grams of helium are in two moles of helium: amount of He in moles amount of He in grams 2.00 mol He x = 8.00 g He 4.00 g He 1 mol He

  26. Gram/Mole Conversions Review sample problem – page 82 Practice problems – Top of page 83, 1-4

  27. Gram/Mole Conversions A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced? mass of Al in grams amount of Al in moles 1 mol Al 27 g Al 11.9 g Al x = 0.44 mol Al

  28. Gram/Mole Conversions Practice problems – page 83 (bottom), 1-2

  29. Atoms / Gram Conversions

  30. Atoms/Molecules and Grams Since 6.02 X 1023 molecules = 1 mole AND1 mole = molar mass (grams) You can convert atoms/molecules to moles and then moles to grams! (Two step process) You can’t go directly from atoms to grams!!!! You MUST go thru MOLES.

  31. Calculations molar mass Avogadro’s numberGrams Moles atoms Everything must go through Moles!!!

  32. Atoms/Molecules and Grams How many atoms of Cu are present in 35.4 g of Cu? 35.4 g Cu 1 mol Cu 6.02 X 1023 atoms Cu 63.5 g Cu 1 mol Cu = 3.4 X 1023 atoms Cu

  33. Problem How many atoms of K are present in 78.4 g of K?

  34. Atoms/Molecules and Grams How many atoms of K are present in 78.4 g of K? 78.4 g K 1 mol K 6.02 X 1023 atoms K 39.0 g K 1 mol K = 12.1 X 1023 atoms K

  35. Problem What is the mass (in grams) of 1.20 X 1024 molecules of glucose (C6H12O6)?

  36. Problem What is the mass (in grams) of 1.20 X 1024 molecules of glucose (C6H12O6)? 1.20 x 1024 mol. 1 mole glucose 180 g glucose 1 mole glucose 6.02 x 1023 mol. = 359 grams glucose

  37. Molar Mass Worksheet: C.5: Molar Masses

  38. Homework Problems – page 87 and 88 Questions: 17 (a-d), 21, 24 (a, c, e and f), 28 (a-e)

  39. Chapter 7 Chemical Formulas and Chemical Compounds

  40. Chapter 7 Section 3 Chemical Formulas

  41. A chemical formula indicates which elements are in a molecule and how many of each element. Example: Water – H2O Subscripts indicate there are two atoms of hydrogen and one atom of oxygen in water.

  42. Example: Aluminum sulfate – Al2(SO4)3 Parenthesis are used to surround the polyatomic group to identify it as a unit. The subscript 3 refers to everything inside the parenthesis. Al – 2 atoms S – 3 atoms O – 12 atoms

  43. Molar Mass for Molecules The molar mass for a molecule = the sum of the molar masses of all the atoms in the molecule.

  44. Example: Molar Mass Example: Find the molar mass for H2O

  45. Example: Molar Mass 1 Count the number of each type of atom Example: Find the molar mass for H2O H 2 O 1

  46. Example: Molar Mass 2 Find the molar mass of each atom on the periodic table Example: Find the molar mass for H2O H 2 1.01 g/mole O 1 16.00 g/mole

  47. Example: Molar Mass 3 Multiple the # of atoms  molar mass for each atom Example: Find the molar mass for H2O  H 2 1.01 g/mole 2.02 g/mole =  16.00 g/mole O 1 16.00 g/mole =

  48. Example: Molar Mass 4 Find the sum of all the masses  H 1 1.01 g/mole 2.02 g/mole = Example: Find the molar mass for H2O  = + 16.00 g/mole O 2 16.00 g/mole 18.02 g/mole 1 mole of H2Omolecules would have a mass of 18.02 g = 18.0 g

  49. Example Example: Find the molar mass for CaBr2

  50. Example: Molar Mass 1 Count the number of each type of atom Example: Find the molar mass for CaBr2 Ca 1 Br 2

More Related