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Chapter Four

Chapter Four. The Structure of the Atom. 4.1 Early Theories of Matter. Philosophers believed matter was made of earth, water, air, and fire. Democritus (460-370 BC ) was first to propose that matter was made up of atomos , which could not be further divided

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Chapter Four

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  1. Chapter Four The Structure of the Atom

  2. 4.1 Early Theories of Matter • Philosophers believed matter was made of earth, water, air, and fire.

  3. Democritus (460-370 BC) • was first to propose that matter was made up of atomos, which could not be further divided • Atoms have different sizes and shapes giving them different properties

  4. John Dalton (1766-1844) Dalton’s Atomic Theory- • Referred to the atom as a “hard sphere”

  5. All matter is made up of atoms • Atoms of the same element are identical • Atoms cannot be created, divided, or destroyed • Atoms combine in certain reactions to form compounds • In chemical reactions atoms are separated, combined, or rearranged

  6. 4.2 Subatomic Particles and the Nuclear Atom • 1879-Crookes invented the cathode ray • Cathode Ray- a ray of radiation that originates from the cathode and travels to the anode of a cathode ray tube

  7. Led to invention of television

  8. By end of the 1800s scientists concluded that cathode rays were a stream of charged particles • Particle carried a negative charge • Electron- negatively charged particle

  9. Evolution of the Atom • J.J. Thomson (1856-1940) determined mass-to-charge ratio of the electron • Determined that charged particle mass was less than that of smallest element, hydrogen

  10. Meant that atoms were made of smaller particles, disproving Dalton’s theory • Created Plum-pudding model of the atom • Robert Millikan (1868-1953) determined that an electron has a negative charge

  11. Thomson’s Plum-pudding or chocolate-chip cookie dough model of the atom Proposed that negatively charged electrons (chips) were distributed through a “dough” of positive charge.

  12. The Nuclear Atom • Ernest Rutherford (1871-1937) used a gold foil experiment to discover existence of nucleus • Nucleus- dense region in center of atom which is positively charged and contains virtually all of its mass

  13. Used a gold foil experiment to see if positive alpha particles would be deflected by the electrons in the atom.

  14. Since the positive charge was thought to be spread out, he thought it would not alter the path of the alpha particles.

  15. Amazingly some were deflected at large angles which meant there must be a concentrated positive area.

  16. Completing the Atom- The Discovery of Protons and Neutrons • Rutherford refined concept of nucleus to include protons and neutrons

  17. Bohr (1913) • Electrons orbit the nucleus • Orbitals have a set size and energy level • Lowest energy level is the smallest orbit

  18. Defining the Atom • Atom- the smallest particle of an element that retains the properties of the element

  19. How big is an atom? consider this: world population in 2000: 6 000 000 000 # of atoms in a single copper penny: 29 000 000 000 000 000 000 000

  20. Proton- subatomic particle carrying a positive charge • Electron -subatomic particle carrying a negativecharge • Neutron- has a mass nearly equal to a proton, but carries no charge (neutral)

  21. Protons, Electrons and Neutrons

  22. Mass

  23. 4.3 How Atoms Differ • Atomic Number- number of protons in an atom • Atomic number= # protons = # electrons

  24. Mass of Individual Atoms • Atoms have extremely small masses which are hard to work with, so scientists use a standard for comparison • Standard used is a Carbon-12 atom

  25. Carbon-12 atom has mass of 12 atomic mass units • 1 atomic mass unit (amu) is nearly equal to mass of 1 proton or 1 neutron • Atomic mass is average mass of isotopes of element

  26. Symbolic Notation atomic number = number of protons = number of electrons Symbol Mass number

  27. Isotope identification Uranium – 238 Element name mass #

  28. Mass Number= # protons + # neutrons • Isotope- atoms with same number of protons but different numbers of neutrons

  29. Percent Abundance • The chance that it will be found in nature

  30. Average Atomic Mass • The average atomic mass is equivalent to the most abundant isotope

  31. Calculating Atomic Mass • Isotopes of elements exist in nature in varying amounts • Atomic Mass = sum of % abundance x atomic mass for each isotope

  32. Calculate the atomic mass unit of chlorine, whose percent abundance is 75% of chlorine-35, and 25% of chlorine -37. • Ex: Chlorine: • Isotopes • Chlorine-35 x (75%) = 26.25 amu • Chlorine-37 (25%) = 9.25 amu • Atomic Mass = 35.5 amu (26.25 + 9.25 amu)

  33. STOP!!!

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