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Chapter 2

Chapter 2. Atoms, Molecules, and Ions. Section 2.1. The early history of chemistry. History. Greeks Democritus and Leucippos - atomos Aristotle- elements Alchemy 1660 - Robert Boyle- experimental definition of element. First “chemist” Stahl – combustion and “phlogiston”

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Chapter 2

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  1. Chapter 2 Atoms, Molecules, and Ions

  2. Section 2.1 The early history of chemistry

  3. History • Greeks • Democritus and Leucippos - atomos • Aristotle- elements • Alchemy • 1660 - Robert Boyle- experimental definition of element. • First “chemist” • Stahl – combustion and “phlogiston” • Priestley- oxygen

  4. Section 2.2 Fundamental Chemical Laws

  5. Antoine Lavoisier • Law of conservation of mass • Wrote first modern textbook • Considered “father of chemistry” • Executed during French Revolution • Joseph Proust followed his lead • Law of definite proportions • John Dalton took all of their work and put it all together • Law of multiple proportions

  6. Laws • Conservation of Mass • Cannot be created/destroyed (transfer) • Law of Definite Proportion- compounds have a constant composition. • They react in specific ratios by mass. • Multiple Proportions • When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.

  7. Section 2.3 Dalton’s atomic theory

  8. 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.

  9. Atomic masses • Dalton put together table of masses for elements • Not always correct • Formulas were wrong • Gay-Lussac • under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume.

  10. A Helpful Observation • Avogadro - interpreted Gay-Lussac’s idea to mean: • at the same temperature and pressure, equal volumes of gas contain the same number of particles. • called Avogadro’s Hypothesis • Did not number the exact number yet • Berzelius • Developed modern symbols for elements

  11. Section 2.4 Early experiments to characterize the atom

  12. The Electron • J. J. Thomson • used Cathode ray tubes

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

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

  15. 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.

  16. Atomizer Oil droplets + - Oil Telescope Millikan’s Experiment

  17. Millikan’s Experiment X-rays X-rays give some electrons a charge.

  18. Millikan’s Experiment Some drops would hover From the mass of the drop and the charge on the plates, he calculated the mass of an electron

  19. Radioactivity • Discovered by accident • Bequerel • Three types • alpha- helium nucleus (+2 charge, large mass) • beta- high speed electron • gamma- high energy light

  20. 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.

  21. Fluorescent Screen Lead block Uranium Gold Foil

  22. What he expected

  23. Because

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

  25. What he got

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

  27. +

  28. Section 2.5 The modern view of atomic structure: an introduction

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

  30. 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.

  31. Symbols A X Z 23 Na 11

  32. Section 2.6 Molecules and ions

  33. Chemical Bonds • The forces that hold atoms together. • Covalent bonding - sharing electrons. • Makes molecules. • Chemical formula- the number and type of atoms in a molecule. • CH4 - 1 carbon atom, 4 hydrogen atoms, • Structural formula shows the connections, but not necessarily the shape.

  34. H • There are also other models that attempt to show three dimensional shape. • Ball and stick. • Space-filling H C H H Structural formula

  35. Ions • Atoms or groups of atoms with a charge. • Cations- positive ions - lose electrons(s). • Anions- negative ions - gain electron(s). • Ionic bonding- held together by the opposite charges. • Ionic solids are called salts. • Polyatomic ions - groups of atoms that have a charge.

  36. Ion List • Yes, you have to memorize them. • List from last year – good start!

  37. Section 2.7 An introduction to the periodic table

  38. PT terminology • Groups – columns – families (vertical) • Similar chemical properties • Periods – rows – series (horizontal)

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

  40. Nonmetals • Brittle • Gain electrons • Covalent bonds

  41. Semi-metals or Metalloids

  42. Alkali Metals

  43. Alkaline Earth Metals

  44. Halogens

  45. Transition metals

  46. Noble Gases

  47. Inner Transition Metals (or Rare Earth)

  48. +1 +2 -3 -2 -1

  49. Section 2.8 Naming simple compounds

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