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Dalton’s Atomic Theory (1808) 1. matter is made up of microscopic, indivisible particles (atoms) that cannot be created or destroyed Our Understanding Today 2. all atoms of an element are identical in mass, and have identical physical and chemical properties 3. atoms of different elements have different masses, physical properties, and chemical properties 4. atoms of different elements combine in simple whole numbers to form compounds 5. atoms of an element cannot be converted into atoms of other elements; chemical reactions involve reorganization of the atoms But Dalton’s Atomic Theory couldn’t explain it all— Why do atoms combine the ways they do? What about the electrically charged particles that were being observed? Observations Leading to the Nuclear Model of the Atom Cathode Ray Tube Experiments (Figure 2.2): J.J. Thomson, 1897 The particles discovered were called electrons. Thomson determined the mass/charge ratio of the electron in a related experiment. Another experiment was needed to determine the mass or charge. Millikan’s Oil Drop Experiment ( Figure 2.3), 1909 Control of oil drop by electric field allowed for calculation of charge of 1 electron: e– charge = –1.602 H 10–19 C Since the mass/charge ratio was already known, the mass could also be calculated: e– mass = 9.109 H 10–28 g So, where’s the positive charge? Speculation was that it was spread throughout the atom. Rutherford’s Alpha Particle Scattering Experiment ( Figure 2.6), 1910 To test this speculation, Rutherford shot α particles (positively charged and relatively massive) at gold foil. Most of the α particles passed through the foil relatively unaffected, but a few bounced back nearly right back at the source. Conclusion: Atoms have a dense, positively charged core called the nucleus. By 1932, Chadwick discovered the neutron (neutral, not electrically charged). Figure 2.6 2-2 The chemistry of an atom is determined primarily by its number and arrangement of electrons. Describing Atoms atomic number, Z: # of protons in the nucleus of an atom mass number, A: sum of # of protons + # of neutrons in the nucleus of an atom A write as Z X or AX or “atomic name–A” (where X = atomic symbol) e.g., isotopes: atoms with same Z but different A. (atoms of same element but with different masses) e.g., Atomic Mass standard: 12C atom mass = 12 atomic mass units (amu) exactly 1 the mass of a 12C atom [= 1 dalton (Da), often used in biochemistry and biology] ˆ1 amu = 12 Using mass spectrometry, we can measure the masses of other atoms relative to 12C. mass of 20 Ne atom = 1.66604 e.g., mass of 12 C atom mass of 20Ne atom = 1.66604 H 12 amu = 19.9924 amu Mass spectrometry also gives information about the relative abundances (fractions) of each isotope of the element. The weighted average of the masses of the isotopes is called the atomic mass. e.g., atomic mass of Ne = 0.904838 (19.9924 amu) + 0.002696 (20.9940 amu) + 0.092465 (21.9914 amu) 20.180 amu PROTIP atomic number (Z) 10 atomic symbol Ne atomic mass Neon 20.180 atomic name 2-3 The atomic mass is closest to the mass of the most abundant isotope of the element. The Periodic Table Mendeleev proposed the first periodic table in 1871, similar to today’s table 1A (1) 1 H 8A (18) 2 He 3 Li 2A (2) 4 Be 3A (13) 5 B 4A (14) 6 C 5A (15) 7 N 6A (16) 8 O 7A (17) 9 F 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 35 Br 36 Kr 1.008 6.941 9.012 10.81 4B (4) 22 Ti 5B (5) 23 V 6B (6) 24 Cr 7B (7) 25 Mn 12.01 14.01 16.00 19.00 4.003 10 Ne 20.18 19 K 20 Ca 3B (3) 21 Sc 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc (98) 44 Ru 101.07 45 Rh 102.91 46 Pd 106.42 47 Ag 107.87 48 Cd 112.41 49 In 114.82 50 Sn 118.71 51 Sb 121.76 52 Te 127.60 53 I 126.90 54 Xe 131.29 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 2B (12) 30 Zn 31 Ga 32 Ge 33 As 34 Se 87.62 88.91 91.22 55 Cs 56 Ba 71 Lu 72 Hf 132.91 137.33 174.97 178.49 180.95 183.84 186.21 190.23 192.22 195.08 197.00 200.59 204.38 207.2 208.98 (209) (210) (222) 87 Fr 88 Ra 103 Lr 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 Cn 113 Nh 114 Fl 115 Mc 116 Lv 117 Ts 118 Og (267) (268) (271) (270) (269) (278) (281) (281) (285) (286) (289) (289) (293) (294) (294) 57 La 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 138.91 140.12 140.91 144.24 (145) 150.36 151.96 157.25 158.93 162.50 164.93 167.26 168.93 173.04 89 Ac 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No [227] 232.04 231.04 238.03 [237] (244) (243) (247) (247) (251) (252) (257) (258) (259) [226] (262) 95.94 (10) 28 Ni 1B (11) 29 Cu 85.47 (223) 92.91 (8) 26 Fe 8B (9) 27 Co METALS, METALLOIDS, NONMETALS Columns: “Groups”; Rows: “Periods” Group Names to LEARN ELEMENTS IN THE O 1A(1): Alkali Metals SAME GROUP O 2A(2): Alkaline Earth Metals ARE USUALLY O 6A(16): Chalcogens CHEMICALLY SIMILAR O 7A(17): Halogens O 8A(18): Noble Gases (somewhat old-fashioned, but still in use) Know/be aware of the A and B group designations as shown above. While now not official, they are too popular and enduring to not be aware of. (They also have uses that we will take advantage of.) “A” group elements are “main group” or “representative” elements. “B” group elements are “transition metals.” The 14-column-wide section at the bottom would fit in between Groups 2A and 3B if there were space. Two numbering conventions are in use. American convention: “A” and “B” groups, as shown above [Europeans reverse the “A” and “B” designations for 3 to 8!] New International convention: Consecutive numbers 1 to 18 (future: 1 to 32?) 2-4