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Atomic Theory

This chapter explores the development of atomic theory, from the Greek belief in the elements to the discovery of the nucleus. It covers the contributions of Democritus, Aristotle, Robert Boyle, Antoine Lavoisier, John Dalton, Joseph John Thomson, William Thomson, and Ernest Rutherford.

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Atomic Theory

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  1. Atomic Theory Chapter 4

  2. History The Greeks thought that matter was composed of fire, earth, water, and air.

  3. Democritus and Leucippos • 400BC • Greek philosophers • Thought all matter was made up of atoms, which are the basic, indivisible particles of matter (atomos) • Alchemists-tried to turn cheap metal into gold

  4. Aristotle • Believed all matter was continuous-did not believe in atoms • His theory (explanation) was accepted for nearly 2000 years

  5. Robert Boyle • Performed the first quantitative experiments. • He studied the relationship between pressure and volume and came up with the definition of an element.

  6. Antoine Lavoisier • The father of modern chemistry. He wrote the first chemistry book. • Performed quantitative experiments where he weighed the reactants and products and verified the law of conservation of mass.

  7. Late 1700’s Scientists agreed that • Most natural materials are mixtures of pure substances. • Pure substances are elements or compounds. • Law of constant composition (Law of definite proportion or Proust’s Law)-a compound always has the same composition.

  8. John Dalton • English schoolteacher who tied all three laws together in his atomic theory • Dalton turned Democritus’s idea into a scientific theory that could be tested by experimentation.

  9. Dalton’s Atomic Theory • Elements are made of tiny particles called atoms. • All atoms of a given element are identical. Atoms of different elements are different. 3. Atoms of one element combine with others to form compounds. 4. Atoms are indivisible in chemical processes. A chemical reaction changes the way the atoms are grouped together.

  10. Dalton (cont) • He came up with the law of multiple proportions-when two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.

  11. What?! • Water (H2O) has 8 g of oxygen per 1 g of hydrogen. • Hydrogen peroxide (H2O2) has 16 g of oxygen per 1 g of hydrogen. • 16/8 = 2/1 • Small whole number ratios.

  12. Proof • Mercury has two oxides. One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass. • Show that these compounds follow the law of multiple proportion. • Speculate on the formula of the two oxides.

  13. A Helpful Observation • Gay-Lussac- compounds always react in whole number ratios by volume if they are under constant temperature and pressure. • Avogadro- interpreted that to mean that at the same temperature and pressure, equal volumes of gas contain the same number of particles (Avogadro’s Hypothesis)

  14. Figure 2.4 Gay-Lussac's Experimental Results

  15. Figure 2.5 A Representation of Combining Gasses at the Molecular Level

  16. Joseph John Thomson • Cathode-ray experiment: showed that the atoms of any element can be made to emit tiny negative particles • Determined the charge ratio of electrons

  17. Voltage source Thomson’s Experiment - +

  18. Voltage source Thomson’s Experiment - +

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

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

  21. William Thomson • Found the electron, but couldn’t find where the positive charge was located. • Plum pudding model-a bunch of positive stuff with the electrons able to be removed.

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

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

  24. 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.

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

  26. Rutherford, Geiger, Marsden-nucleus • Gold foil experiment, which led to the discovery of the nucleus. • Used uranium to produce alpha particles. • Aimed alpha particles at gold foil by drilling a hole in a lead block. • Like bullets through a tissue

  27. Florescent Screen Lead block Uranium Gold Foil

  28. What he expected

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

  30. What he got

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

  32. +

  33. Atomic Structure

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

  35. Table 2.1 The Mass and Charge of the Electron, Proton, and Neutron

  36. Protons • The number of protons in an atom determines the element’s identity • Nuclear forces hold the nuclear particles together • The atomic number equals the number of protons

  37. Electrons • Very small If the nucleus is a grape, the electrons would be about one mile away. • Have a negative charge • The arrangements of electrons determines the element’s chemical properties.

  38. Neutrons • Mass number= protons+neutrons • Neutrons=mass number-atomic number • Isotope-atoms that have the same number of protons and electrons but different numbers of neutrons (disproves point 2 of Dalton’s theory) • Nuclide-any isotope of any element

  39. Figure 2.15 Two Isotopes of Sodium

  40. 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).

  41. Symbols A X Z 23 Na 11

  42. Practice-Give the protons, neutrons, and electrons for each • Mercury • Sodium • Carbon • 13C 6

  43. Answers • Mercury 80p, 80e, 121n • Sodium 11p, 11e, 12n • Carbon 6p, 6e, 6n • 13C 6p, 6e, 7n 6

  44. Ions • An ion is formed when we remove or add an electron to a neutral atom. • Cation-a positive ion • Anion-a negative ion

  45. Ion Example Regular sodium has 11 electrons, 11 protons, and 12 neutrons. If we take away 1 electron, it would have 10 electrons (-), and 11 protons (+) so the charge would be +1. (Neutrons would stay the same).

  46. Periodic Table

  47. Mendeleev's Early Periodic Table, Published in 1872-arranged by mass number

  48. Organization Groups/families-vertical columns Periods-horizontal rows

  49. Metals Conductors Lose electrons (cations +) Malleable and ductile

  50. Nonmetals Brittle Gain electrons (anions -) Covalent bonds

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