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Chemistry 100(02) Fall 2013

Chemistry 100(02) Fall 2013. Instructor: Dr. Upali Siriwardane e-mail : upali@coes.latech.edu Office : CTH 311 Phone 257-4941 Office Hours : M,W, 8:00-9:30 & 11:30-12:30 a.m Tu,Th,F 8 :00 - 10:00 a.m.   Or by appointment Test Dates :.

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Chemistry 100(02) Fall 2013

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  1. Chemistry 100(02) Fall 2013 Instructor: Dr. UpaliSiriwardane e-mail: upali@coes.latech.edu Office: CTH 311 Phone257-4941 Office Hours: M,W, 8:00-9:30 & 11:30-12:30 a.m Tu,Th,F8:00 - 10:00 a.m.   Or by appointment Test Dates: September 30, 2013 (Test 1): Chapter 1 & 2 October 21, 2013 (Test 2): Chapter 3 & 4 November 13, 2013 (Test 3) Chapter 5 & 6 November 14, 2013 (Make-up test) comprehensive: Chapters 1-6 9:30-10:45:15 AM, CTH 328

  2. REQUIRED: Textbook:Principles of Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro - Pearson Prentice Hall and also purchase the Mastering Chemistry Group Homework, Slides and Exam review guides and sample exam questions are available online: http://moodle.latech.edu/ and follow the course information links. OPTIONAL: Study Guide: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro 2nd Edition Student Solutions Manual: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro2nd Text Book & Resources

  3. Chapter 2. Atoms and Elements 2.1 Imaging and Moving Individual Atoms…………….. 43 2.2 Early Ideas about the Building Blocks of Matter……. 45 2.3 Modern Atomic Theory and the Laws That Led to It… 45 2.4 The Discovery of the Electron……………………….. 49 2.5 The Structure of the Atom……………………………. 51 2.6 Subatomic Particles: Protons, Neutrons, and Electrons in Atoms……………………………………………… 53 2.7 Finding Patterns: The Periodic Law and the Periodic Table…………………………………………. 58 2.8 Atomic Mass: The Average Mass Of an Element’s Atoms. 64 2.9 Molar Mass: Counting Atoms by Weighing Them……… 66

  4. Chapter 2. KEY CONCEPTS

  5. 1) What are following experimental techniques that are being used to image individual atoms?

  6. Scanning Tunneling Microscope

  7. Microscopy Microscopes 1) Optical Microscopes 2) SPB-Scanning Probe Microscopy a)STM-Scanning Tunneling Microscope b) Atomic Force Microscope 3) Electron beam Techniques SEM-Scanning Electron Microscope TEM-Transmission Electrum Microscope

  8. Alchemist: Discovery of Elements • Early scientist observed chemical changes of matter. They called these changes chemical reactions when there are changes in substances or the physical & chemical properties of the matter. They also observed a pattern or a repeatable observation during chemical reactions.

  9. Law of the Conservation of Matter Lavoisier proposed from his experimental evidence the following law: Matter is neither created nor destroyed in a chemical reaction. • Total mass of used reactants = Total mass of products produced • Total number of reactant atoms = Total number of product atoms

  10. Three Observed Chemical Laws: • Law of Conservation of Mass: • Law of Constant Proportions: • Law of Multiple Proportions:

  11. Early scientist observed changes of matter • They called these changes chemical reactions when there are changes in substances or the chemical properties of matter. They also observed a pattern or a repeatable observation in chemical reactions. They observed that Mass was neither destroyed nor created • (E.g. Hydrogen (4g) + oxygen (32g) gives water 36g • after the reaction), • and elements combine in Constant Proportions • (E.g. 36g of water contains 4g of hydrogen and 32g of oxygen) and in compounds in 1:8 • Multiple Proportion • (E.g. In CO, 1g C contains 1.33g of O and In CO21g C contains 2.66 g of O).

  12. Law of Multiple Proportions Law of multiple proportions: • Two elements A and X can form different compounds by combining in different proportions. • These combinations can be represented as a ratio. • For example: • A molecule of carbon dioxide (CO2) has a ratio of 1 C atom to every 2 atoms of oxygen, or 1:2. • A molecule of hydrogen peroxide (CO) has a ratio of 1 C atom to 1 atom of oxygen, or 1:1.

  13. 2) What are the patterns of observations in conducting chemical reactions by early chemists? • a) • b) • c)

  14.  3) What theory John Dalton came up with to explain the body of chemical observations and laws?

  15. Dalton’s atomic theory • All matter is composed of atoms -- the smallest particle of an element that takes part in a chemical reaction. • All atoms of an element are alike. • Compounds are combinations of atoms of one or more elements. The relative number of atoms each element is always the same. • Atoms cannot be created or destroyed by a chemical reaction. They only change how they combine with each other.

  16. 4) What are the postulates of Dalton’s atomic theory? • a) • b) • c) • d)

  17. 5) How was Dalton’s atomic theory modified based on new experimental observations leading to the discovery of the electron, nucleus, protons and neutrons? Reword Dalton’s postulates to accommodate new observations and particles. • a) • b) • c) • d)

  18. Radioactivity • Becquerel (1896) • Uranium ore emits rays that “fog” a photographic plate. • Marie and Pierre Curie (1898) • Isolated 2 new elements (Po and Ra) that did the same. • Marie Curie called the phenomenon radioactivity.

  19. 6) What is radioactive decay?

  20. + β γ − α Beam of α, β, and γ Radioactive material Electrically Charged plates screen Radioactivity Types of Radiation Alpha ray α (positive charge) Beta ray β (negative charge) Gamma ray γ (no charge) Electrical behavior: + attracted to - (opposites attract) (like charges repel)

  21. 7) What are following radiation? a) a b) b c) g

  22. 8) Which of the following radiation, a,b, and g is most harmful?

  23. – high voltage + Electrons • Thomson (1897) studied cathode rays and discovered the electron: fluorescent screen cathode ray • Beam travels from the cathode (-) to the anode (+). • the beam flies through a ring anode and hits a fluorescent screen. • The cathode rays come from the cathode metal. • They are negative particles – electrons (e−).

  24. + – – high voltage + Electrons • Thomson showed that electric and magnetic fields deflect the beam. From the deflections, Thomson calculated the mass/charge ratio for an e-: = −5.60 x 10-9 g/C (Coulomb (C) = the SI unit of charge)

  25. The Discovery of the Electron, Nucleus • Subatomic Particles: Protons, Neutrons, and Electrons in Atoms • 9) How did Thompson know that every element has electrons? • 10) How did Thompson know that an electron has a negative charge?

  26. 11) In the Millikan’s oil drop experiment, how did he remove electrons from atoms? • 12) In the Millikan’s oil drop experiment, where did some of the electrons removed from atoms ended up? • 13) In the Millikan’s oil drop experiment, why was some oil drops had multiples (1,2,3 of −1.60 x 10-19 ) of charges?

  27. 14) Thomson calculated the mass/charge (m/e) ratio for an e- to be = −5.60 x 10-9 g/C. and then Millikan found the charge on an e- to be −1.60 x 10-19 C. What is the mass on an electron?

  28. 15) In the Rutherford’s experiment, what caused a few α’s were deflected through large angles and some came almost straight back!

  29. Electronic Charge Robert Millikan (1911) studied electrically-charged oil drops. • For a single charged drop, he measured: • the time to fall a fixed distance, and • to rise the same distance in an electric field. • He showed that each drop had a charge that was an integer multiple of −1.60 x 10-19 C. • (The charge of an electron. ) • The modern value is −1.602176462 x 10-19 C. • (Often written in “atomic units” as charge = −1).

  30. Millikan’s Experiment

  31. mass charge me = charge x Mass of an Electron The experiments by Thomson and Millikan gave the mass/charge ratio and charge of an e−. = (−1.60 x 10-19 C)/(−5.60 x 10-9 g/C) = 8.96 x 10-28 g The modern value is: me =9.10938188 x 10-28g

  32. Atomic Structure: Plum-Pudding Model • J. J. Thomson (plum-pudding model) • - The atom is composed of a positive cloud of matter in which electrons are embedded. • Explains the positive (+), negative (-) charged behavior of matter

  33. Nuclear Atom • Thompson thought it was a ball of uniform positive charge, with small negative dots (e-) stuck in it. However “plum pudding” model was short lived and was changed to Nuclear model.

  34. Rutherford’s Gold Foil Experiment Setup Gold foil experiment: Could not explain Thomson’s plum-pudding atom model. Led to the discovery of the atom’s nucleus.

  35. Rutherford & the Nucleus: Gold Foil Experiment From the gold foil experiment, the following conclusions were proposed: • The atom contains a tiny, dense center called the nucleus. • The nucleus has essentially the entire mass of the atom. • The electrons weigh so little they give practically no mass to the atom. • The nucleus is positively charged. • The amount of positive charge balances the negative charge of the electrons. • The electrons are dispersed in the empty space of the atom surrounding the nucleus.

  36. Determination of nuclear charge • Rutherford estimated that the charge of the nucleus of an atom was about one half of the atomic mass. • Moseley, while working for Rutherford, developed a more accurate measurement. • While working with cathode rays on metal targets, he measured the wavelength of the X-rays produced. • He found that a direct relationship exists between the metal’s atomic number and the square root of the frequency.

  37. Discovery of Protons and Atomic Number • Moseley, Henry & Gwyn Jeffreys • 1887–1915, English physicist. • studied the relations among x-ray spectra of different elements. • concluded that the atomic number is equal to the charge on the nucleus based on the x-ray spectra emitted by the element. • explained discrepancies in Mendeleev’s Periodic Law.

  38. Determination of nuclear charge Moseley concluded that the charge of the nucleus was an integer. Further, it was the same as the number of electrical units (electrons) but of opposite charge. Atomic number X-Ray Frequency1/2

  39. Summary of Subatomic Particles Remember: Atoms are usually electrically neutral, Indicating equal numbers of protons and electrons!

  40. Atomic number, Z • The number of protons in the nucleus • The number of electrons in a neutral atom • The integer on the periodic table for each element

  41. Relative size of atom and atomic nucleus

  42. Ions Charged single atom Charged cluster of atoms • Cations: positive ions • Anions: negative ions Ionic compounds: combination of cations and anions with zero net charge

  43. Nuclear Notation X= atomic symbol A= mass number Z= atomic number C-12, carbon-12 XA C12 ZXA 6C12

  44. Mass Number, A • integer representing the approximate mass of an atom • equal to the sum of the number of protons and neutrons in the nucleus

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