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Chemistry Chapter 4

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  1. Chemistry Chapter 4 The Structure of the Atoms

  2. Top Ten

  3. Table 3.1

  4. Table 3.3

  5. History of Chemistry • 400 B.C. Greeks proposed 4 elements • Earth • Fire • Water • Air • Next 2000 years—alchemy • During this period discoveries were made • Hg, S, Sb • prepared acids

  6. Controversial Greek Thinking! Democritus (460-370 B.C): -Matter is composed of “atomos” (now atoms) -Atoms were homogeneous & indivisible -Could not answer what holds atoms together Aristotle (384 B.C.-322B.C.): -Matter was continuous and indefinitely divisible (did not believe in atoms) -Matter made of earth, fire, air, & water -Idea was accepted for nearly 2000 years!

  7. Indivisible or Divisible?Democritis vs. Aristotle

  8. Late 1700’s • Most chemists accepted element definition • Understood elements combined to form compounds with various properties • Disagreed whether compounds are always in the same ratio

  9. What happened in 1790? • Study of matter was revolutionized by new emphasis on Quantitative Analysis • Aided by improved balances • Measurements were actually ACCURATE!!!

  10. Robert Boyle • Founder of Modern Chemistry (1627-1691) • Took the “Al” out of Alchemy (although he started as one) • First scientist to understand the importance of careful measurement • Insisted science be based on experiments • Famous for P=1/V

  11. Antoine Lavoisier • Fatherof Modern Chemistry (1743-1794) • Recognized and named hydrogen and oxygen • Introduced the metric system • Wrote first list of elements and revised nomenclature • Because of prominence in pre-revolutionary government, was beheaded at the height of the French revolution

  12. John Dalton—Beginning of Modern Atomic Theory • Englishman from a Quaker family (1766-1844) • Revolutionized chemistry by emphasizing that atoms can have weights and weights can be measured (quantitative) • Opened a school at age 12 • Color blind/researched • Interested in botany • Theory not accepted until 1905 Albert Einstein paper

  13. Dalton’s Atomic Theory (1808) • Matter is composed of extremely small particles called atoms • Atoms are indivisible and indestructible. • Atoms of a given element are identical in size, mass, and chemical properties. • Atoms of specific element are different from those of another element. • Different atoms combine in simple whole-number ratios to form compounds. • In chemical reactions, atoms are separated, combined, or rearranged

  14. Dalton vs. Today

  15. Law of Conservation of Mass Mass is neither created nor destroyed during chemical or physical reactions. Total mass of reactants = Total mass of products Antoine Lavoisier

  16. Law of Multiple Proportions • If two or more different compounds are composed of the same two elements, then the ratio is always small whole numbers. (CO, CO2)

  17. What does this mean? (Law of Definite Composition) • 50.0 g sample of pure H2O decomposed into its elements • would find 5.6 g H and 44.4 g oxygen • % mass would be: mass H = 5.60 g x 100 = 11.2% H total mass 50.0 g mass 0 = 44.4 g x 100 = 88.8% O total mass 50.0 g

  18. Law of Definite (or Constant) Composition • The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound.

  19. Figure 3.2: Representation of NO, NO2, and N2O.

  20. What does an atom look like? (Sketch it on your paper!)

  21. This is The Modern Atomic Model • Atom: The smallest particle of an element that retains the properties of the element • Only seen by STM (Scanning Tunneling microscope)

  22. Subatomic Particles

  23. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

  24. Figure 3.7: Schematic of a cathode ray tube.

  25. Some ModernCathode Ray Tubes

  26. Mass of the Electron 1909 – Robert Millikan determines the mass of the electron. The oil drop apparatus Mass of the electron is 9.1 x 10-28 g

  27. Conclusions from the Study of the Electron • Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. • Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons • Electrons have so little mass that atoms must contain other particles that account for most of the mass

  28. Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. Based on the following facts: (1) atoms contain small, negatively charged particles called electrons and (2) the atoms of the element behave as if they have no charge at all

  29. Ernest Rutherford • 1871-1937 • Learned physics in JJ Thomson’s lab • Did much work with alpha particles (+ charged part with mass) • Most famous for his GOLD FOIL EXPERIMENT

  30. Figure 3.5: Rutherford’s experiment.

  31. Try it Yourself! In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?

  32. The Answers Target #1 Target #2

  33. Figure 3.3: Plum Pudding model of an atom.

  34. Figure 3.6: Results of foil experiment if Plum Pudding model had been correct.

  35. Figure 3.6: Actual results.

  36. Most of the particles passed right through • A few particles were deflected • VERY FEW were greatly deflected Rutherford’s Findings “Like howitzer shells bouncing off of tissue paper!” Conclusions: • The nucleus is small • The nucleus is dense • The nucleus is positively charged

  37. Disbelievers…. • Albert Einstein when to his grave not totally believing it • According to classical physics, the electron would have collapsed into the nucleus • 1910-1930 began the Quantum Physics Revolution (the physics of atomic and subatomic particles)

  38. The Atomic Scale • Most of the mass of the atom is in the nucleus (protons and neutrons) • Electrons are found outside of the nucleus (the electron cloud) • Most of the volume of the atom is empty space “q” is a particle called a “quark”

  39. The Quark… Oops…wrong Quark!

  40. About Quarks… Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons”

  41. Figure 3.9:A nuclear atom viewed in cross section.

  42. Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.

  43. Isotopes Elements occur in nature as mixtures of isotopes. Isotopes are atoms of the same element that differ in the number of neutrons

  44. Figure 3.10: Two isotopes of sodium.

  45. Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p+ + n0 18 8 8 18 Arsenic 75 33 75 Phosphorus 16 15 31

  46. Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011

  47. Isotopes…Again (must be on the test) Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.

  48. Chlorine Practice Problem • Chlorine exists as 2 isotopes in nature. Cl-35 (atomic mass 34.969 amu) has a 75.77% relative abundance. Cl-37 has an atomic mass 36.966 amu. • What is the % abundance of the Cl-37 isotope?

  49. Chlorine Practice Problem • Chlorine exists as 2 isotopes in nature. Cl-35 (atomic mass 34.969 amu) has a 75.77% relative abundance. Cl-37 has an atomic mass 36.966 amu Calculate the atomic mass of Chlorine.