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Chapter 5

Chapter 5. Atoms, Molecules, and Ions. History. Greeks Democritus and Leucippus - atomos Aristotle- 4 elements. Alchemy-experimentation 1660 - Robert Boyle- experimental definition of element. Lavoisier- Father of modern chemistry. He wrote the book. Laws. Conservation of Mass.

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Chapter 5

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  1. Chapter 5 Atoms, Molecules, and Ions

  2. History • Greeks • Democritus and Leucippus - atomos • Aristotle- 4 elements. • Alchemy-experimentation • 1660 - Robert Boyle- experimental definition of element. • Lavoisier- Father of modern chemistry. • He wrote the book.

  3. Laws Conservation of Mass • In any chemical reaction the total mass of the reactants will equal the total mass of the products 36.04 g Reactants 36.04 g Products • 2H2 + O2 2H2O 4.04g + 32.00g = 36.04g

  4. Law of Definite Proportion- The law stating that a pure substance, e.g. H2O, will always have the same percent by weight, e.g. 11.2% H  and 88.8% O. elements will combine in specific ratios by mass.

  5. Law of Definite Proportions H20 H 2 x 1.01g = 2.02g O 1 x 16.00g = 16.00g 18.02g 2.02g H / 18.02 g H2O = 0.11 x 100 = 11% 16.00g H / 18.02 g H2O = 0.88 x 100 = 88 % What are the proportions in a CO2 molecule?

  6. Law of Multiple Proportions In chemistry, the law of multiple proportions is sometimes called Dalton's Law after its discoverer, the English chemist John Dalton. The law is: If two elements form more than one compound between them, then the ratios of the masses of the second element which combine with a fixed mass of the first element will be ratios of small whole numbers.

  7. For example carbon oxide: CO and CO2 100 grams of carbon may react with 133 grams of oxygen to produce carbon monoxide (CO) 100 grams of carbon may react with 266 grams of oxygen to produce carbon dioxide (CO2) The ratio of the masses of oxygen that can react with 100 grams of carbon is 266:133 ≈ 2:1, a ratio of small whole numbers.

  8. What?! • Compare Water H2O and Hydrogen Peroxide H2O2 • H2O - H 2 x 1.01g = 2.02 g O 1 x 16.00g = 16.00 g 16.00 g Oxygen / 2.02 g Hydrogen = 8gO/1gH • Water has 8 g of oxygen per 1 g of hydrogen.

  9. Law of Multiple Proportions H2O2 – H 2 x 1.01g = 2.02g O 2 x 16.00g = 32.00g 32.00 g Oxygen / 2.02g Hydrogen = 16gO/1gH Hydrogen peroxide has 16 g of oxygen per g of hydrogen. Compare the 2 samples 16g H2O2 / 8g H2 = 2 / 1 ratio Small whole number ratios.

  10. Dalton’s Atomic Theory • Elements are made up of atoms • Atoms of each element are identical. Atoms of different elements are different. • Compounds are formed when atoms combine. Each compound has a specific number and kinds of atom. • Chemical reactions are rearrangement of atoms. Atoms are not created or destroyed.

  11. Experiments to determine what an atom was • J. J. Thomson- used Cathode ray tubes

  12. Voltage source Thomson’s Experiment - +

  13. Voltage source Thomson’s Experiment - +

  14. Voltage source Thomson’s Experiment - + • Passing an electric current makes a beam appear to move from the negative to the positive end.

  15. Voltage source Thomson’s Experiment • By adding an electric field

  16. Voltage source Thomson’s Experiment + - • By adding an electric field, he found that the moving pieces were negative

  17. Thomsom’s Model • Found the electron. • Couldn’t find positive (for a while). • Said the atom was like plum pudding. • A bunch of positive stuff, with the electrons able to be removed.

  18. Rutherford’s Experiment • Used uranium to produce alpha particles. • Aimed alpha particles at gold foil by drilling hole in lead block. • Since the mass is evenly distributed in gold atoms alpha particles should go straight through. • Used gold foil because it could be made atoms thin.

  19. Florescent Screen Lead block Uranium Gold Foil

  20. What he expected

  21. Because

  22. Because, he thought the mass was evenly distributed in the atom.

  23. What he got

  24. + How he explained it • Atom is mostly empty • Small dense, positive pieceat center. • Alpha particlesare deflected by it if they get close enough.

  25. +

  26. Modern View • The atom is mostly empty space. • Two regions • Nucleus- protons and neutrons. • Electron cloud- region where you might find an electron.

  27. Sub-atomic Particles • Z - atomic number = number of protons determines type of atom. • A - mass number = number of protons + neutrons. • Number of protons = number of electrons if neutral.

  28. Symbols A X Z 23 Na (p+ and n0) Mass # 11 (p+) Atomic#

  29. Periodic Table Click Link

  30. Metals • Conductors • Lose electrons • Malleable and ductile

  31. Nonmetals • Brittle • Gain electrons • Covalent bonds

  32. Semi-metals or Metalloids

  33. Alkali Metals

  34. Alkaline Earth Metals

  35. Halogens

  36. Transition metals

  37. Noble Gases

  38. Inner Transition Metals

  39. +1 +2 -3 -2 -1

  40. Ions • Atoms or groups of atoms with a charge. • Cations- positive ions - metals losing electrons(s). • Anions- negative ions - nometals gaining electron(s). • Ionic bonding- held together by the opposite charges. • Ionic solids are called salts.

  41. Polyatomic Ions • Groups of atoms that have a charge. • No, you don’t have to memorize them. • List on page 97 or on the back of your periodic table.

  42. Naming compounds • Two types • Ionic - metal and non metal or polyatomics. • If the metal is in the 1st or 2nd column (representative element). NO (roman numeral) in the name • If the metal is a transition metal a (roman numeral) may be needed (check your periodic table) • Covalent- we will just learn the rules for 2 non-metals.

  43. Ionic compounds • If the cation is monoatomic- Name the metal (cation) just write the name. Mg 2+ Magnesium • If the cation is polyatomic- name it. NH4+ Ammonium • If the anion is monoatomic- name it but change the ending to –ide. Cl1- Chloride • If the anion is polyatomic- just name it PO43- Phosphate • Practice.

  44. Ionic Compounds • Have to know what ions they form • off table, polyatomic, or figure it out • Two types • Binary ionic – 2 elements (cation + anion) • Ternary ionic – 3 elements • (cation + Polyatomic ion) • CaS K2S • AlPO4 K2SO4 • FeS CoI3

  45. Ionic Compounds • Fe2(C2O4) • MgO • MnO • KMnO4 • NH4NO3 • Hg2Cl2 • Cr2O3

  46. Ionic Compounds • KClO4 • NaClO3 • YBrO2 • Cr(ClO)6

  47. Writing Formulas • Two sets of rules, ionic and covalent • To decide which to use, decide what the first wordor element is. • If it is a metal or polyatomic use ionic. • If it is a non-metal use covalent.

  48. Ionic Formulas • Charges must add up to zero. • Get charges from table, name of metal ion, or memorized from the list. • Use parenthesis to indicate multiple polyatomics.

  49. Ionic Formulas • Sodium nitride • sodium- Na is always +1 • nitride - ide tells you it comes from the table • nitride is N-3

  50. Ionic Formulas • Sodium nitride • sodium- Na is always +1 • Nitride - ide tells you it comes from the table • nitride is N-3 • Doesn’t add up to zero. Na+1 N-3

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