1 / 65

Atomic Structure

Atomic Structure. Unit 4 Expected test date: 28-Oct-08. The Atom. The smallest particle of an element that still retains the properties of the element. Atoms are much too small to see What does it look like?. Subatomic Particles. Atoms are made of smaller particles Protons Neutrons

maryawells
Download Presentation

Atomic Structure

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Atomic Structure Unit 4 Expected test date: 28-Oct-08

  2. The Atom • The smallest particle of an element that still retains the properties of the element. • Atoms are much too small to see • What does it look like?

  3. Subatomic Particles • Atoms are made of smaller particles • Protons • Neutrons • Electrons

  4. The Structure of Atoms

  5. The Philosophers • Ancient Greeks • Formed ideas of nature based on their experiences • Did no experiments, just thought

  6. The ancient idea • Everything made of : Fire Earth Air Water

  7. Democritus(460 – 370 B.C.) • Greek philosopher (not scientist) • Named the atom • Different kinds of atoms have different sizes and shapes. • Apparent changes in matter result from changes in the groupings of atoms and not from changes in the atoms themselves.

  8. John Dalton(1766-1844) • First atomic theory • All matter is composed of atoms. • All atoms of a given element are identical and different from those of any other element. • Atoms combine in simple whole number ratios to form compounds. • In a chemical reaction, atoms are separated, combined, or rearranged.

  9. Early Theories of Matter • Workbook: Page 19 • Numbers 1 through 12

  10. By the late 1800s… • Cathode ray tube invented

  11. JJ Thomson(1897) • Cathode rays made of particles smaller than atoms • First subatomic particles • Discovered the electron • Plum pudding model of the atom

  12. Electron • Charge of negative one (-1) • Almost no mass (1/1840 mass of proton) • Exists in the empty space around the nucleus

  13. Questions raised by Thomson • If electrons are particles smaller than atoms, are there other particles? • What makes up most of the mass of the atom? • If electrons are negatively charged, why are atoms neutral?

  14. Plum pudding model • Atom is a sphere • Positive charge is evenly distrubuted • Negatively charged electrons spread randomly through the sphere

  15. Let’s draw that • Workbook: Page 20 • Number 4

  16. Ernest Rutherford(1911) • Aimed a beam of “alpha particles” at a thin gold foil • The beam made a bright spot on a fluorescent screen

  17. Florescent Screen Lead block Uranium Gold Foil

  18. What Rutherford expected

  19. What Rutherford expected α α α α α α

  20. What Rutherford saw

  21. + What Rutherford saw α α α α α α

  22. Rutherford concluded • There had to be a heavy central core to the atom: the nucleus • Nucleus contains protons

  23. Nucleus • center of an atom • extremely small, positively charged, dense • contains protons, neutrons • surrounded by empty space where electrons move

  24. Proton • Subatomic particle • Exists in the nucleus • Has a positive charge (+1) • Has a mass of 1 atomic mass unit

  25. James Chadwick(1932) • Experimenting with radiation sources • Discovered a new particle with no charge : the neutron

  26. Neutron • subatomic particle • found in an atom’s nucleus • has a mass nearly equal to that of a proton (1 atomic mass unit) • Has a neutral charge (no charge)

  27. Nuclear Atomic Model • Neutrons (n0) • Protons (p+) • Electrons (e-) Workbook: Page 20, #5

  28. Electrons, Protons and Neutrons • Workbook: Page 20 • Numbers 5

  29. Structure of the atom • 99.97% of mass in nucleus • Most of the volume is empty space • Electrons in cloud

  30. How big is an atom? • Simulate the size of a hydrogen atom: • Nucleus : place a baseball on the 50 yard line of Reliant Stadium • Electron : put a grain of sand on the back row of the highest section

  31. Electrons, Protons and Neutrons • Workbook: Page 20 • Numbers 6, 7, 8

  32. Sub-subatomic particles • Protons, and neutrons are made of even smaller particles • 6 flavors of quarks • up, down, charm, strange, top, and bottom

  33. Differences in atoms • Different atoms - different numbers of protons and neutrons • The number of protons determines what the element is

  34. 6 C 12.0107 Wait… • Proton mass = 1 • Neutron mass = 1 • So where did those decimal places come from?

  35. 19 K 39.0983 Isotopes and atomic mass • Potassium has 19 protons • How many neutrons? 20 potassium atoms: 93.25% will have 20 neutrons, 6.7302% will have 22 neutrons, 0.0117% will have 21 neutrons

  36. Isotopes • Atoms of the same element with the same number of protons but different numbers of neutrons

  37. Video

  38. Atomic Number • The number of protons in an atom • Always a whole number • Number of electrons = number of protons Atomic number = protons = electrons

  39. 6 C 12.0107 Mass Number(whole number) • Total number of protons and neutrons in a given isotope Carbon-12 has 6 protons + 6 neutrons = 12 atomic mass units

  40. Carbon -12 is standard • Atomic Mass Unit (amu) is 1 / 12 the mass of a carbon – 12 atom • 1 amu is nearly (not exactly) equal to the mass of one proton or one neutron • Because of this, an atom’s mass is nearly equal to the number of protons and neutrons in its nucleus

  41. Isotopes • Atoms of the same element with the same number of protons but different numbers of neutrons

  42. 6 C 12.0107 Atomic Mass is an average(decimal) • An average of all known isotope mass numbers for an element • The atomic mass of an element is closest to the most common isotope found in nature • So, since Carbon’s atomic mass is closest to 12, carbon-12 would be the most common isotope found in nature

  43. Atoms review • Most atoms are neutral • Protons = electrons • Number of protons determines the element • Atomic number = number of protons

  44. Atoms on the periodic table

  45. Let’s Practice! • Complete 1-5 in study guide on page 21. • False • True • False (can be, but not always) • True • False

  46. More Practice! • Questions 6-12 on page 21. • 82 protons and 82 electrons • 8 protons • 30 • 85 • 104 protons and electrons • 84 protons and electrons • 102 protons and electrons

  47. Proton, neutron & electron relationships • Mass number = protons + neutrons • An isotope has an atomic number of 6 and a mass number of 13 • Number of protons, neutrons and electrons? Protons = 6 13 amu = 6 protons + _____ neutrons Electrons = 6 7

  48. Relationship Practice • On pages 21-22, questions 13-17. • Protons = 19, Electrons = 19, Neutrons = 20 • Protons = 14, Electrons = 14, Neutrons = 14 • Protons = 19, Electrons = 19, Neutrons = 21 • Protons = 51, Electrons = 51, Neutrons = 72 • 13 and 15 are both isotopes of Potassium

  49. There are two ways to represent elements: • Symbol: Mass Number =? _____________ Mass Number X # element symbol atomic # # (# of p+) Atomic Symbols # of p+ and n0 • Name: name of element followed by mass number. • Ex... Aluminum - 27 Atomic number =13 Nitrogen - 14 Atomic number =7 Atomic number =6 Carbon - 14 What are the Atomic Numbers for the above names?

  50. Symbol Name Nitrogen - 14 Nitrogen - 15 N N 14 15 7 7 For Example • There are two ways to represent nitrogen isotopes:

More Related