2. Atoms, Molecules, and Ions Topic 1 in review book Chapter 5 (p107) in Textbook

3. HISTORY n Democritus 400 BC suggested that matter is made up of indivisible particles called atoms.

4. First Atomic Theory Dalton (1766-1844) n All elements are composed of indivisible atoms. n Atom of the same elements are identical. n Atoms of different elements can combine with one another to form compounds n Chemical reactions occur when atoms are separated, joined or rearranged. n Atoms of one element cannot change into atoms of another element as result of a chemical reaction.

5. The Law of Conservation of Mass n Established in 1789 by French Chemist Antoine Lavoisier n States that mass is neither created nor destroyed in any ordinary chemical reaction. n Or more simply, the mass of substances produced (products) by a chemical reaction is always equal to the mass of the reacting substances (reactants).

6. Experiments to determine what an atom was n J. J. Thomson (1897)- used Cathode ray tubes

7. Thomson’s Experiment Voltage source +-

8. Thomson’s Experiment Voltage source +-

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

10. Thomson’s Experiment n By adding an electric field

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

12. November 30 n Experiments that lead to find what is inside the atom n Thomsom cathode rays – electrons n Rutherford gold foil experiment – n Millikan n OBJECTIVE : WHAT IS INSIDE THE ATOM?

13. Thomsom’s conclusions n Thomsom named the cathodic rays electrons and concluded that they must be a part of atoms of all elements.

14. The Atom, circa 1900: n “Plum pudding” model, put forward by Thompson. n Positive sphere of matter with negative electrons imbedded in it.

15. Millikan (1868-1953) n Calculated the charge of an electron and its mass with a famous experiment.

16. Millikan Oil Drop Experiment Robert Millikan (University of Chicago) determined the charge on the electron in 1909.

17. Radioactivity: n The spontaneous emission of radiation by an atom. n First observed by Henri Becquerel. n Also studied by Marie and Pierre Curie.

18. Radioactivity n Three types of radiation were discovered by Ernest Rutherford: –  particles –  particles –  rays

19. The Gold foil experiment n In 1911 Ernest Rutherford designed an experiment using alpha particles and thin gold foil - “gold foil experiment”. Rutherford was very surprised by his finding and his conclusions led to which became known as the “planetary” model of the atom.

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

21. Lead block Uranium Gold Foil Florescent Screen

22. What he got

23. Discovery of the Nucleus Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

24. How he explained it + n Atom is mostly empty n Small dense, positive piece at center. n Alpha particles are deflected by it if they get close enough.

25. +

26. The Nuclear Atom Since some particles were deflected at large angles, Thompson’s model could not be correct.

27. Observations n Most of the alpha particles pass straight through the gold foil. n Some of the alpha particles get deflected by very small amounts. n A very few get deflected greatly. n Even fewer get bounced of the foil and back to the left.

28. Conclusions n The atom is 99.99% empty space. n The nucleus contains a positive charge and most of the mass of the atom. n The nucleus contains a positive charge and most of the mass of the atom. n The nucleus is approximately 100,000 times smaller than the atom.

29. Vocabulary n Deflected: curved or bent downward. n Scattered: to separate and drive off in various directions; disperse n Bounce:

30. Homework n From review book questions 1 to 8 topic 1 n (Page 4)

31. Other Subatomic Particles n Protons were discovered by Rutherford in 1919. n Neutrons were discovered by James Chadwick in 1932.

32. Modern Atomic Theory n The atom is mostly empty space. n Two regions n Nucleus- protons and neutrons. n Electron cloud- region where you might find an electron.

33. The Bohr Model n Niels Bohr proposed a model of the atom based on the solar system. n It was wrong but some of his ideas we still used. It was one step in the direction of the last model for the atom. n First let’s take a look at what is inside the atom.

34. n Atoms are neutral and contain same number of electrons than protons. In a chemical reactions atoms can lose, gain or share electrons. n When an atom loses an electron it becomes a positive ion. When it gains an electron it becomes a negative ion.

35. n The nucleus always remains the same in a chemical reaction. n The number of protons and neutrons never change.

36. Subatomic Particles n Protons and electrons are the only particles that have a charge. n Protons and neutrons have essentially the same mass. n The mass of an electron is so small we ignore it.

37. Atomic Number (Z) n The number of protons in the nucleus. It identifies the element. In the neutral atom n # protons = # electrons

38. Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z)

39. Mass Number (A) * It is not in the periodic table n Protons plus neutrons in the atom n A= # protons + # neutrons n *In the table we can find the Atomic Mass. n Number of Neutrons =Mass number-Atomic number Number of Neutrons Number of Neutrons

40. Atomic Mass Unit (amu) is one-twelfth the mass of a carbon-12 atom.

41. Mass Number The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

42. December 7 n n How to calculate the Atomic Mass of an element n n Objective: Isotopes n n Atomic Mass n n Hw : Finish worksheet n n Page 11 Review Book n n Questions 13 to 26

43. Mass Number (A) * It is not in the periodic table n Protons plus neutrons in the atom n A= # protons + # neutrons n *In the table we can find the Atomic Mass. n Number of Neutrons =Mass number-Atomic number Number of Neutrons Number of Neutrons

44. Isotopes n atoms with same number of protons and different number of neutrons. Same atomic number different mass number. n SAME ELEMENT with different # of neutrons!

45. Isotopes, Atomic Numbers, and Mass Numbers Atomic number (Z) = number of protons in the nucleus. Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). By convention, for element X, we write Z A X. Isotopes have the same Z but different A. We find Z on the periodic table.

46. Isotopes: n Atoms of the same element with different masses. n Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C

47. Atomic Mass n is the weighted average of the mass of all the isotopes of one element. n It takes into consideration the mass of the different isotopes of the element and their natural abundance

48. Nuclear Charge n n Is the number of protons!

49. Ion n n A charged particle n n When an atom loses or gains electrons it becames an ION. n n Remember protons can not be touched! n n If an atom loses electron the ion will be positive (more protons than electrons) n n If an atom gains electrons it will became negative (more electrons than protons)

50. Bohr’s Atom electrons in orbits nucleus

51. Planetary Model n Neils Bohr 1913 n Electrons are arranged in orbits around the nucleus. n He proposed that electrons in a particular orbit have a fixed energy. The electrons cannot fall into the nucleus.

52. Energy level n is the region around the nucleus where the electron is moving. n As the electrons are further away from the nucleus they have more energy. The closest orbits or energy levels to the nucleus have low energy.

53. n He named the orbits with letters beginning with k,l,m,n… n The first energy level could hold 2 electrons. n Second energy level could hold up to 8 electrons

54. Electron Configurations n Indicates how the electrons are located in the atom. n Niels Bohr proposed that electrons are located in energy levels at different distances from the nucleus. n 2-8-3 total of electrons=13 n 3 energy levels

55. HELIUM ATOM + N N + - - proton electron neutron Shell What do these particles consist of?

56. ATOMIC STRUCTURE Particle proton neutron electron Charge + ve charge -ve charge No charge 1 1 nil Mass

57. ATOMIC STRUCTURE MASS NUMBER the number of protons and neutrons in an atom the number of protons in an atom He 2 4 number of electrons = number of protons Atomic number

58. ATOMIC STRUCTURE Electrons are arranged in Energy Levels or Shells around the nucleus of an atom. first shella maximum of 2 electrons second shella maximum of 8 electrons third shella maximum of 8 electrons

59. ATOMIC STRUCTURE There are two ways to represent the atomic structure of an element or compound; 1.Electronic Configuration 2.Dot & Cross Diagrams

60. SUMMARY 1. The Atomic Number of an atom = number of protons in the nucleus. 2. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. 3. The number of Protons = Number of Electrons. 4. Electrons orbit the nucleus in shells. 5. Each shell can only carry a set number of electrons.

61. Modern Theory

62. The Nature of Energy Another mystery involved the emission spectra observed from energy emitted by atoms and molecules. When gases at low pressure were placed in a tube and were subjected to high voltage, light of different colors appeared

63. Continuous Spectra Radiation composed of only one wavelength is called monochromatic. Radiation that spans a whole array of different wavelengths is called continuous. White light can be separated into a continuous spectrum of colors. Note that there are no dark spots on the continuous spectrum that would correspond to different lines. Line Spectra and the Bohr Model

65. Line Spectra If high voltage is applied to atoms in gas phase at low pressure light is emitted from the gas. If the light is analyzed the spectrum obtained is not continuous. SPECTROSCOPE

66. Line Spectra. When the light from a discharge tube is analyzed only some bright lines appeared.

67. Bohr’s Model n Niels Bohr adopted Planck’s assumption about energy and explained the hydrogen spectrum this way: 1. Only orbits of certain radii corresponding to certain definite energies are permitted for the electron in the hydrogen atom.

68. Bohr Model n 2 An electron in a permitted orbit has a specific energy an is in an “allowed” energy state. It will not spiral into the nucleus n 3 Energy is emitted or absorbed by the electron only as the electron changes from one allowed state to other