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CHAPTER 3 – THE ATOM Read pgs. 107 - 110 I. History Democritus – Greek philosopher, 400 B.C., said all matter is made up of small, indivisible particles he called “atoms” (Greek for “indivisible”). He wasn’t believed until the late 1700’s John Dalton – 1803, realized that the Law of Conservation of Mass and the Law of Definite composition could only be explained if atoms existed. Wrote Dalton’s Atomic Theory, which was mostly right. 1. Matter is composed 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. Different atoms combine in simple whole- number ratios to form Compounds 5. In a chemical reaction, atoms are separated, combined or rearranged. Atom – the smallest particle of an element that can exist, either alone or in combination with other atoms. 1 As new info came out, Dalton’s theory was modified into the: MODERN ATOMIC THEORY 1. All matter is made of atoms. 2. All atoms of a given element have the same chemical properties, which differ from the chemical properties of other elements. 3. Atoms of a given element may not all have the same mass, but a sample of the element will have an average mass which is characteristic of that element. 4. Atoms are not subdivided during physical or chemical rxns. 5. Compounds are formed when 2 or more elements combine; each atom loses its characteristic properties as a result. II. The structure of the Atom A. Discovery of the electron --by Thomson in 1897 (pgs. 109 – 110 for background info) --Robert Millikan – “oil-drop experiment” in 1909 found that the mass of an e- is very small compared to the mass of the whole atom. Also discovered that e-‘s have a measureable amount of negative electric charge. Q. How could he have measured the mass of an e- when even today we don’t have a scale that will measure that? [Hint: How can you find the thickness of a page? (use a ruler)] --Ernest Rutherford – in 1908- 1909 bombarded thin metal foils with positively-charged particles. (see pg. 111 - 112 – know his exp.!) 2 Rutherford’s conclusion: The positively-charged particles that bounced back were repelled by a powerful positive charge which must occupy a very small area within the atom, since most of the particles passed through without being affected. (This also explains the deflected particles.) He named this area within an atom the nucleus – very small and extremely dense. Occupied by positively charged particles he called protons. (He concluded that most of the rest of the atom was empty space.) C. Protons – every neutral atom has the same number of protons as electrons, so the positive and negative charges will balance. D. Neutrons – discovered in 1932 by Chadwick. Are neutral (no electrical charge). Mass is virtually identical to the mass of a proton; both of these particles have a much larger mass than an electron has (see chart, pg. 111) E. Fission – in 1939, Lise Meitner and Otto Hahn discovered that an atom could be subdivided. They called this process “fission”. (Six years later we dropped a fission bomb on Hiroshima, Japan). III. Weighing Atoms A. Atomic Number (Z) – the number of protons in the nucleus of a given element’s atoms. Each element has a different # of protons in their nucleus. That is what makes elements different from each other. Elements are placed in order in the Per. Table from 1 118 by their atomic number. So: Atomic number of oxygen = ____8______ # of p in an oxygen atom = ____8______ 3 # of e- in an oxygen atom = _____8_____ Atomic number of hydrogen = ____1______ # of p in a neon atom = ___10_______ # of e- in an iodine atom = ____53______ B. Mass Number – tells the total number of protons plus neutrons in the nucleus (doesn’t count e-‘s since they virtually have no mass). To find the mass number of an element, round the “atomic mass” (the decimal number in the P.T.) to the nearest whole number. Ex: Na atomic mass 22.98977, mass number = __23_______ K atomic mass = __39.0983__, mass number = __39___ To find number of neutrons in an atom: # of n = mass# - atomic# 4 Element Na Atomic # 11 Atomic mass 22.989770 Mass# 23 #p 11 #n 12 #e 11 U 92 238.0289 238 92 146 92 Pb 82 207.2 207 82 125 82 Fe 26 55.847 56 26 30 26 Ag 47 107.868 108 47 61 47 P 15 30.97376 31 15 16 15 F 9 18.998403 19 9 10 9 Ba 56 137.33 137 56 81 56 Al 13 26.98154 27 13 14 13 Li 3 6.941 7 3 4 3 C. Isotopes: Atoms of the same element, which have different masses because they have different numbers of neutrons in the nucleus. Ex: Most common isotope of H has: 1 proton 0 neutrons mass number = 1 Other isotopes of H: “Deuterium” 1 proton 1 neutron mass number = 2 “Tritium” 1 proton 2 neutron mass number = 3 5 Two methods of designating isotopes: 1. Hyphen notation “Name-mass #” ex: Hydrogen-1 Hydrogen-2 Hydrogen-3 (normal) (deuterium) (tritium) Ex: the isotope of uranium used in nuclear power plants is Uranium-235. Mass # = ___235_______ # of p = ____92_______ # of n = ____143_______ 2. Nuclear Symbol Method Mass # 3 H symbol (this is Hydrogen-3) Atomic # 1 Uranium-235 is shown as 35 Write the nuclear symbol for Chlorine-35 ___17 Cl_______ 206 Write the nuclear symbol for Lead-206 __82 Pb________ How many protons, neutrons, and electrons are in one atom of Chlorine-37? # of p = __17________ (atomic number, Z) # of e = __17________ (atomic number, Z) # of n = ___20_______ (mass # - atomic #) 6 Practice: 1. How many p, n, and e are in an atom of Bromine-80? _35____ p, _45____ n, __35___e 13 2. Write the nuclear symbol for Carbon-13. ____6 C______ 3. Write the hyphen-notation for the element that contains 15 p and 15 n. _______Phosphorus - 30________________________ D. Relative Atomic Mass The atomic mass of an element = the mass of 1 atom of that element. Since these atomic masses are so small (tiny fractions of a gram) they have a designated a new unit: “Atomic Mass Unit” (u) 1 u = 1/12th the mass of a Carbon-12 atom (everything is compared to Carbon-12) So, the mass of 1 atom of carbon-12 is 12 u. The atomic masses given on the Per. Table are averages for all isotopes of that element combined, given in “u”. Mass of 1 atom of He = __4.003 u__________ Mass of 1 atom of H = __1.00794 u_________ Atomic mass of Mg = ___24.3050 u________ 7 IV. Counting Atoms A. The Mole – a counting unit (abbrev. “mol”) “One Dozen” = 12 of anything “One Mole” = 6.022 x 1023 of anything Ex: one mole of cars = 6.022x1023 cars one mole of students = 6.022x1023 students one mole of copper atoms = 6.022x1023 atoms This number (6.022x1023) is called Avogadro’s Number How large is it? If every person on earth (5 billion +) counting 1 atom per second from a 1 mole sample of Cu, it would take 4 million years to count the atoms. B. Molar Mass – the mass, in grams, of 1 mole of a pure substance. Conveniently, this equals the atomic mass from the Per. Table, with grams as the unit. Ex: Mass of 1 atom of H = 1.00794 u Mass of 1 mole of H atoms = 1.00794 g Mass of 1 mole of He = ___4.003 g____________ Mass of 1 mole of Na = ____22.989770 g________________ Mass of 1 mole of Cu = ____63.546 g________________ B. Calculating with Moles **Use factor-label method **when using molar mass of Avogadro’s Number as a conversion factor, round it to the sig figs allowed in the problem before using it. **Moles should end up in every problem. 8 Ex: What is the mass, in grams, of 3.50 moles of Cu? 3.50 mol Cu x __63.5_______g Cu = _____222 g Cu___________ 1 mol Cu Ex: What is the mass, in grams, of 0.375 moles of K? 0.375 mol K x 39.1 g K = 14.7 g K 1 mol K Ex: You have 11.9 g of Al. How many moles of Al do you have? 11.9 g Al x 1 mol Al = 0.441 mol Al 27.0 g Al Ex: How many moles of Au are in 3.60 x 10-10 g Au? 3.60 x 10-10 g Au x 1 mol Au = 1.83 x 10-12 197 g Au 9 Ex: How many moles of Ag are in 3.01 x 1023 atoms of Ag? 3.01 x 1023 atoms Ag x 1 mol Ag__________ = 0.500 mol Ag 6.02 x 1023 atoms Ag Ex: 2500 atoms Sn = __?__ mol Sn? 2500 atoms Sn x _____1 mol Sn___ = 4.2 x 10-21 mol Sn 6.0 x 1023 atoms Sn Ex: How many atoms of Al are in 2.75 mol Al? 2.75 mol Al x 6.02 x 1023 atoms Al = 1.66 x1024 atoms Al 1 mol Al 10 Ex: How many atoms of Zn are in 4.93 mol of Zn? 4.93 mol Zn x 6.02 x 1023 atoms Zn = 2.97 x 1024 atoms Zn 1 mol Zn Ex: What is the mass in grams of 7.5 x 1015 atoms of Ni? 7.5 x 1015 atoms Ni x ___1 mol Ni_____ x 59 g Ni = 7.4 x 10-7 g Ni 6.0 x 1023 atoms Ni 1 mol Ni Ex: How many atoms of sulfur are contained in 4.00 g of sulfur? 4.00 g S x ___1 mol S x 6.02 x 1023 atoms S = 7.50 x 1022 atoms S 32.1 g S 1 mol S 11