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 = # of p in an oxygen atom = 3
# of e- in an oxygen atom = Atomic number of hydrogen = # of p in a neon atom = # of e- in an iodine atom = 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 = K atomic mass =, mass number = To find number of neutrons in an atom: # of n = mass# - atomic# Element Atomic # Atomic mass Mass# #p #n #e Na U Pb 4
Fe Ag P F Ba Al Li 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 Two methods of designating isotopes: 1. Hyphen notation Name-mass # ex: Hydrogen-1 (normal) Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 5
Ex: the isotope of uranium used in nuclear power plants is Uranium-235. Mass # = # of p = # of n = 2. Nuclear Symbol Method Mass # 3 H symbol Atomic # 1 (this is Hydrogen-3) Uranium-235 is shown as Write the nuclear symbol for Chlorine-35 Write the nuclear symbol for Lead-206 How many protons, neutrons, and electrons are in one atom of Chlorine-37? # of p = (atomic number, Z) # of e = (atomic number, Z) # of n = (mass # - atomic #) Practice: 1. How many p, n, and e are in an atom of Bromine-80? p, n, e 2. Write the nuclear symbol for Carbon-13. 6
3. Write the hyphen-notation for the element that contains 15 p and 15 n. 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 = 1/12 th 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 = Mass of 1 atom of H = Atomic mass of Mg = IV. Counting Atoms A. The Mole a counting unit (abbrev. mol ) One Dozen = 12 of anything One Mole = 6.022 x 10 23 of anything 7
Ex: one mole of cars = 6.022x10 23 cars one mole of students = 6.022x10 23 students one mole of copper atoms = 6.022x10 23 atoms This number (6.022x10 23 ) 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 = Mass of 1 mole of Na = Mass of 1 mole of Cu = 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. Ex: What is the mass, in grams, of 3.50 moles of Cu? 3.50 mol Cu x g Cu = mol Cu 8
Ex: What is the mass, in grams, of 0.375 moles of K? Ex: You have 11.9 g of Al. How many moles of Al do you have? Ex: How many moles of Au are in 3.60 x 10-10 g Au? Ex: How many moles of Ag are in 3.01 x 10 23 atoms of Ag? 9
Ex: 2500 atoms Sn =? mol Sn? Ex: How many atoms of Al are in 2.75 mol Al? Ex: How many atoms of Zn are in 4.93 mol of Zn? 10
Ex: What is the mass in grams of 7.5 x 10 15 atoms of Ni? Ex: How many atoms of sulfur are contained in 4.00 g of sulfur? 11