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AP Notes Chapter 2 Atoms and Elements History of the atom Summed-up Greeks Democritus and Leucippus - atomos Aristotle- elements. Alchemy 1660 - Robert Boyle- experimental definition of element. Lavoisier- Father of modern chemistry. Dalton’s Atomic Theory 1) Elements are made up of atoms 2) Atoms of each element are identical. Atoms of different elements are different. 3) Compounds are formed when atoms combine. Each compound has a specific number and kinds of atom. 4) Chemical reactions are rearrangement of atoms. Atoms are not created or destroyed. The Atom Dalton (early 1800s) indivisible A Helpful Observation Gay-Lussac- under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. Avagadro- interpreted that to mean at the same temperature and pressure, equal volumes of gas contain the same number of particles. (called Avagadro’s Hypothesis) Experiments & theories to determine what an atom was John Dalton- atoms indivisible J. J. Thomson- Cathode ray tubes, electrons Marie Curie- radioactivity Robert Millikan- electron mass & charge Ernest Rutherford- protons James Chadwick- neutrons Thomson’s Experiment Voltage source - + Thomson’s Experiment Voltage source - + Thomson’s Experiment Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end. Thomson’s Experiment Voltage source By adding an electric field Thomson’s Experiment Voltage source + By adding an electric field, he found that the moving pieces were negative The Atom Thompson (~ 1900) cloud of (+) charge . .. . . . ... . electron (-) charge “plum pudding” model Thomsom’s Model Found the electron. Couldn’t find positive (for a while). Said the atom was like plum pudding. A bunch of positive stuff, with the electrons able to be removed. Millikan’s Experiment Atomizer Oil droplets + - Telescope Oil Millikan’s Experiment X-rays X-rays give some electrons a charge. Millikan’s Experiment •Some drops would hover From the mass of the drop and the charge on the plates, he calculated the mass of an electron Radioactivity Discovered by accident Henri Bequerel – photographic plates Marie Curie – studied & named it Three types alpha- helium nucleus (+2 charge, large mass) beta- high speed electron gamma- high energy light James Chadwick Neutrons Particles from radioactive polonium hit a beryllium target and produced particle no charge slightly greater mass than the proton Rutherford’s Experiment Used uranium to produce alpha particles. Aimed alpha particles at gold foil by drilling hole in lead block. Since the mass is evenly distributed in gold atoms alpha particles should go straight through. Used gold foil because it could be made atoms thin. Rutherford’s Experiment Lead block Uranium Florescent Screen Gold Foil Rutherford’s Experiment What he expected Rutherford’s Experiment Because Rutherford’s Experiment Because, he thought the mass was evenly distributed in the atom. Rutherford’s Experiment What he got Rutherford’s Experiment How he explained it Atom is mostly empty Small dense, positive piece at center. Alpha particles are deflected by it if they get close enough. + Rutherford’s Experiment + Gold Foil Experiment Rutherford (~1911) Nuclear Model .. . . . . . ... heavy central (+) nucleus e- “about” nucleus The Atom Rutherford (~ 1911) (e-) about nucleus . . .. .. . heavy central (+) nucleus Nuclear Model The Atom Bohr (~ 1913) central (+) nucleus n=3 n=2 n=1 . .... . . . . .. . e- in allowed orbits Planetary Model Modern View The atom is mostly empty space. Two regions Nucleus- protons and neutrons. Electron cloudregion where you might find an electron. The Atom Heisenberg, de Broglie, Schroedinger (mid 1920s) e- in regions defined by math functions . .. .. . .. . . . Quantum Mechanical Model Sub-atomic Particles Z - atomic number = number of protons determines type of atom. A - mass number = number of protons + neutrons. Number of protons = number of electrons if neutral. Nuclear Symbols & Notation Mass Number→ A X Z 23 24 Na Na 11 11 ←Element Symbol Atomic Number→ Isotopes of elements Isotopes are forms of an atom that differ by the number of neutrons Mass number is approximation of exact atomic mass of an isotope Atomic mass or atomic weight is the average mass of the isotopes of atoms Isotopic percent abundance or fractional abundance is a description of the proportion of an isotope in a sample of an element Atomic Mass Atoms are so small, it is difficult to discuss how much they weigh in grams. Use atomic mass units. an atomic mass unit (amu) is one twelth the mass of a carbon-12 atom. This gives us a basis for comparison. The decimal numbers on the table are atomic masses in amu. They are not whole numbers Because they are based on averages of atoms and of isotopes. can figure out the average atomic mass from the mass of the isotopes and their relative abundance. add up the percent as decimals times the masses of the isotopes. Isotopes of Hydrogen 1 1 H hydrogen 2 1 H deuterium 3 1 H tritium Examples There are two isotopes of carbon 12C with a mass of 12.00000 amu(98.892%), and 13C with a mass of 13.00335 amu (1.108%). There are two isotopes of nitrogen , one with an atomic mass of 14.0031 amu and one with a mass of 15.0001 amu. What is the percent abundance of each? Percent Abundance Percent abundance = number of atoms of a given isotope x 100% total number of atoms of all isotopes Fractional Abundance Fractional abundance = Percent Abundance 100% Atomic Weight = (abundance isotope 1)(weight isotope1) + (abundance isotope 2)(weight isotope2)… or %iso1 %iso 2 AW (Wiso1) (Wiso 2) .... 100 100 A portion of an atom’s mass of protons, neutrons and electrons is converted to energy that holds the atom together. Einstein gave us ΔE = (Δm)C2 The loss of this mass as the atom forms is called the mass defect. This missing mass is converted to “binding energy” (BE) Mass atom = BE + #pro. + #elec. + #neu. Allotrope Different forms of the same element that exist in the same physical state under the same conditions of Temperature & Pressure Carbon •Diamond •Graphite Graphite Diamonds Buckyballs Periodic Table Metals Conductors Lose electrons Malleable and ductile Nonmetals Brittle Gain electrons Covalent bonds Semi-metals or Metalloids Alkali Metals Alkaline Earth Metals Halogens Transition metals Noble Gases Inner Transition Metals Periodic Table 1A Families or Groups 2A 3A 4A 5A 6A 7A 3B 4B 5B 6B 7B 8B 1B 2B 8A Periodic Table Periods 1 2 3 4 5 6 7 8 Lanthanide Series Actinide Series Periods and Groups or Families Hydrogen The Hindenburg crash, May 1939. Shuttle main engines use H2 and O2 Group 1A: Alkali Metals Potassium Reaction of potassium + H2O Cutting sodium metal Group 2A: Alkaline Earth Metals Magnesium Magnesium Ablaze! Magnesium oxide Calcium Carbonate—Limestone The Appian Way, Italy Champagne cave carved into chalk in France Group 3A: B, Al, Ga, In, Tl Aluminum Boron halides BF3 & BI3 Gems & Minerals Sapphire: Al2O3 with Fe3+ or Ti3+ impurity gives blue whereas V3+ gives violet. Ruby: Al2O3 with Cr3+ impurity Transition Elements Lanthanides and actinides Iron in air gives iron(III) oxide Colors of Transition Metal Compounds Iron Cobalt Nickel Copper Zinc Group 4A: C, Si, Ge, Sn, Pb Quartz, SiO2 Diamond Group 5A: N, P, As, Sb, Bi White and red phosphorus Phosphorus Phosphorus first isolated by Brandt from urine, 1669 Group 6A: O, S, Se, Te, Po Sulfuric acid dripping from snot-tite in cave in Mexico Sulfur from a volcano Group 7A: F, Cl, Br, I, At Halogen Group 8A: He, Ne, Ar, Kr, Xe, Rn Lighter than air balloons “Neon” signs XeOF4