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Chapter 3 Notes History of Atomic Theory Democritus Greek philosopher - 460-370 B.C. Believed the world is made up of small indivisible particles called atoms and empty space. Ideas based on observations. Aristotle Greek philosopher - ~400 B.C. Matter is continuous with no indivisible particles. Aristotle “won” the debate with Democritus. His idea was accepted (vs Democritus) until the 1700s. Lavoisier (Antoine) French chemist in the 1780s Observed that the mass before and mass after a reaction remained the same. Law of Conservation of Matter (mass) In a normal chemical reaction, matter cannot be created or destroyed. Proust (Joseph) French chemist ~1800 Law of Definite Proportions Elements combine in a certain proportion or ratio (by mass) to form a compound. Example: H2O is always 11% hydrogen and 89% oxygen Dalton (John) English schoolteacher and chemist in the early 1800s. Law of Multiple Proportions. Proposed the atomic theory of matter (aka, the Modern Atomic Theory). Modern Atomic Theory 1. All elements are composed of small, indivisible particles called atoms. 2. Atoms of the same element are identical. Atoms of different elements are different. 3. Atoms combine in simple whole-number ratios to form compounds. 4. Chemical reactions occur, but atoms are never changed into atoms of another element. Changes: Atoms of the same element may be different : isotopes Atoms can be divisible. Berzelius (Jons Jakob) Swedish chemist – 1811 Established modern system of chemical symbols for elements. Thomson (JJ) British physicist – late 1800s/early 1900s. Using Crookes’ tubes (cathode ray tubes), he discovered electrons. Determined … (1) electrons are negatively charged particles. (2) all electrons are identical. (3) the charge to mass ratio using deflection of rays. Plum Pudding Model of Atoms The atom is a sphere of positive matter in which negative electrons are embedded. Millikan (Robert) American physicist - ~1909 Oil drop experiment. Discovered the charge of an electron. It has a charge of -1. Discovered the mass of an electron. It is 1/1840 of a proton. It has a mass of 0 amu. Goldstein (Eugen) German physicist – 1886 Discovered canal rays – also called anode rays. This led to the development of mass spectrometry. Wien (Wilhelm) German physicist – 1898 Discovered the proton. It has a +1 charge (equal and opposite to the electron). It has a mass of 1 amu. Rutherford (Ernest) British chemist/physicist – 1911 Performed the Gold Foil Experiment with the help of Geiger and Marsden. Alpha particles were “shot” at a thin sheet of gold foil. Results of Gold Foil: (1) Most particles (93%) went straight through the foil. (2) Some particles were slightly deflected. (3) About 1 in 8000 particles came straight back. Conclusions from Gold Foil: (1) The atom is mainly empty space. (2) The atom contains a small, dense positively charged core called the nucleus. Predicted the existence of neutrons. The mass of the protons in the nucleus was only about ½ of the known mass of the nucleus. Predicted a neutral particle must exist in the nucleus. Chadwick (James) English physicist – 1932 Discovered the neutron. It has a 0 charge. It has a mass of 1 amu. Worked with Geiger and Rutherford. Parts of the Atom Protons Positive (+) charge In the nucleus Contribute to atomic mass Determine the atomic number Never gained or lost Electrons Negative (-) charge Outside the nucleus In different energy levels; each row on the periodic table is a new energy level Do not contribute to atomic mass A proton’s mass is 1800 times greater than that of an electron Gained, lost or shared in bonding Neutrons Neutral charge (no charge) In the nucleus Contribute to atomic mass Never gained or lost Different isotopes of the same element have a different number of neutrons Atomic number Equal to the number of protons Identifies the element Is the whole number on the periodic table Examples: H-1, He-2, Li-3 At. # = # of p+ = # of e- (in the elemental state) Mass number (also called atomic mass number) Equal to the atomic number + the number of neutrons Equal to the number of protons + the number of neutrons Is a whole number; is not shown on the periodic table Mass # = At. # + # neut. Mass # = # p+ + # neut. Average atomic mass A weighted average of the masses of all the isotopes for a given element Examples: H-1.01; Cl-35.45 Isotopes Elements that contain the same number of protons but a different number of neutrons Example: C-12 and C-14; both have 6 p+ and 6e-; C-12 has 6 neutron and C-14 has 8 neutrons Determine the following for each atom: (1) element name; (2) element symbol; (3) atomic number; (4) atomic mass number; (5) # of protons; (6) # of electrons; (7) # of neutrons. 1. 2. 3. 28 protons 30 neutrons 1. Nickel 2. Ni 3. 28 4. 58 Tungsten – 186 1. Tungsten 2. W 3. 74 4. 186 Pb 124 neutrons 1. Lead 2. Pb 3. 82 4. 206 5. 28 6. 28 7. 30 5. 74 6. 74 7. 112 5. 82 6. 82 7. 124