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Chapter 4: The Structure of the Atom Early Ideas about Matter Name Date Democritus Aristotle John Dalton 400 BC atomos – smallest particle 350 BC hyle – continuous matter 1803 AD billiard ball – indivisible, smallest particle JJ Thomson 1897 AD plum pudding – electrons in positive pudding mass ( / charge) ratio e– 1910 AD discovered the charge of the electron 1911 AD nuclear model – small, dense nucleus with esurrounding empty space Robert Millikan Ernest Rutherford Theory Experiment none, philosophical argument none, philosophical argument laws of: conservation of matter definite proportion multiple proportion Cathode ray tube / e- beam (a Crookes tube) Oil drop experiment (gravity, e– charge, and charged plates) α – particle / gold foil Dalton’s Atomic Theory 1. Matter is composed of extremely small particles called atoms 2. Atoms are indivisible and indestructible 3. Atoms of a given element are identical in size, mass, and chemical properties 4. Atoms of a specific element are different from those of another element 5. Different atoms combine in simple whole-number ratios to form compounds 6. In a chemical reaction, atoms are separated, combined, or rearranged Atomic structure Atoms are made of three particles 1. protons (p+) 2. neutrons (n0) 3. electrons (e–) } nucleus Isotopes: atoms of the same element with different numbers of neutrons Examples: hydrogen isotopes are protium, deuterium, and tritium Hyphen notation: protium = hydrogen-1, deuterium = hydrogen-2, tritium = hydrogen-3 Nuclear symbols: protium = H , deuterium = H, tritium = H from Wikipedia: http://en.wikipedia.org/wiki/Isotope Calculating nucleons in isotopes Mass number = number of p+ + number of n0 Atomic number = number of p+ = mass number – atomic number Number of n0 (top number in a nuclear symbol) (bottom number in a nuclear symbol) Examples: (for neutral atoms) H Li n0 = 2 n0 = 4 p+ = 1 p+ = 3 e– = 1 e– = 3 Relative atomic masses The atomic mass unit (amu) is defined as of a carbon – 12 nuclide Mass defect explains why He = 4.002 603 254 15(6) but W = 181.948 204 2(9) + The more p squeezed into a nucleus, the more mass they must convert to binding energy (by Einstein’s famous equation E = mc2) Atomic masses on the Periodic Table are not whole numbers because they are the weighted average of all the naturally occurring nuclides of that isotope Example: magnesium has three nuclides: 24Mg at 78.99% and 23.985 041 7 amu, 25 Mg at 10.00% and 24.985 836 92 amu, and 26Mg at 11.01% and 25.982 592 93 amu. 24 Mg: 23.985 041 7 amu x 0.789 9 = 18.945 784 amu Mg: 24.985 836 92 amu x 0.100 0 = 2.498 584 amu 26 Mg: 25.982 592 93 amu x 0.110 1 = 2.860 683 amu 25 24.305 amu Unstable Nuclei and Radioactive Decay Three common types of radioactive decay Alpha (α) = He Beta (β) = e (or e–) Gamma (γ) = EMR (electromagnetic radiation or very high energy light) Examples: α: Ra → Rn + He or Ra → Rn + α β: C → N + e or Ra → Rn + e– Th + He + 2 γ or U→ Th + α + 2 γ γ: U→ Nuclear stability: The ratio of the number of p+ to n0 is the primary factor in causing the instability that results in nuclear decay Stable nuclei for small atomic numbers have a ratio of 1 p+ to 1 n0 Stable nuclei for high atomic numbers have a ratio of 1 p+ to 1.5 n0 Unstable nuclei will undergo a series of α and/or β decays until they reach a stable p+ to n0 ratio