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History of Atomic Theory Scientists B.C. Democritus 400 BC Coined the term “atom” Aristotle 384-322 BC Believed matter is continuous Dalton’s Atomic Theory -early 1800’s- All matter is composed of tiny, indivisible particles called atoms. Atoms of the same element have the same properties (mass, size, etc.). In a chemical reaction, matter cannot be created or destroyed. (Law of Conservation of Mass) Compounds always contain elements in the same ratio by mass (Law of Definite Proportions) Atomic size A penny contains 2.4 x 1022 atoms Radius of one atom is around 2 x 10-10m or .2 nm Scanning tunneling microscope can generate images of individual atoms. Thomson’s Cathode Ray Tube -late 1800’s- Showed that electrons are negatively charged particles. Image from Addison Wesley Chemistry Thomson’s “plum pudding” model Rutherford’s gold foil exp. -early 1900’s- Conclusion: Most of an atom’s volume is empty space. Rutherford’s “planetary” model 5 Models of the Atom (a) Dalton's model (1803) (c) Rutherford's model (1909) (e) Electron-cloud model (present) (b) Thomson's model (1897) (d) Bohr's model (1913) © Prentice-Hall, Inc. Subatomic charge location mass Other feature particle proton + Nucleus 1 amu Defines the element -atomic no. neutron 0 Nucleus 1 amu Change no. to form isotopes electron - Electron ~0 cloud atom’s volume -dictates reactivity Nuclear Forces Short-range forces that hold the nuclear particles together. Isotopes Atoms of the same element that differ in mass Atomic no.=# protons #protons=#electrons Mass no.=#protons + # neutrons (nucleons) Num f neutrons Isotopes of Hydrogen Nuclide Protons Neutrons Protium 1 0 Mass Number 1 Deuterium 1 1 2 tritium 1 2 3 Isotopes can be written two ways 108 47 Ag 207 82 Pb 80 35 Br or bromine-80 Electrons Found in an electron cloud outside of the nucleus (but not in paths like the planets) 1st energy level holds 2 electrons 2nd energy level holds up to 8 3rd energy level holds up to 18 Periodic Table Arranged by increasing atomic number Rows are called periods Columns are called groups Average Atomic Mass An element’s atomic mass is the weighted average of its naturally occurring isotopes. Average Atomic Mass Multiply the mass of each isotope by its abundance to get the weighted average. (% x mass)+ (% x mass) + . . . 100 Ex.: Boron is 80.20% boron-11 (atomic mass 11.01 amu) and 19.80% boron-10 (atomic mass 10.01 amu). What is the average atomic mass of boron? (11.01amu)(80.20) + (10.01amu)(19.80) = 100 =10.81 amu Sample Problem ex.: Neon has 2 isotopes. Neon-20 has a mass of 19.992amu and neon-22 has a mass of 21.991amu. In an average sample of neon atoms, 90% will be neon-20 and 10% will be neon-22. Calculate the average atomic mass. (90 x 19.992amu)+(10 x 21.991amu) 100 = 20.192 amu