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Chapter Two Atoms_ Molecules and Ions

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					Chapter Two: Atoms, Molecules and Ions
    • I.       Models of the Atom (Atomic Theory)
           – A. Greek (400 BC) Two models
               » 1.      All matter was continuous: Aristotle
               » 2.      Elementary particles: Atomos theory by Democritus
                 and Leucippos
           – B. Dalton Theory (1803)
               » 1.      Beginning of the modern ideas about atoms
               » 2.      Read about John Dalton on page 65 in text
               » 3.      Law of Multiple Proportions: Molecules are formed
                 by combination of atoms in a fixed proportion (ratio).
    •          CO       carbon monoxide
    •          CO2      carbon dioxide
    •          H2 O 2   hydrogen peroxide
               H2O      dihydrogen oxide
C. Thomson’s Model: J. J. Thomson
   1. Discovered the electron in 1897. (see Fig 2.3 in our text)

   2. Thomson reasoned that if an atom has a negative charge it
      must also have a positive charged particle.

   3. His model was a scoop of chocolate chip ice cream where the
      electrons are the chips and the ice cream is a sea of positive
      charge. (he called it a “Plum pudding” model)

   4. At this same time Becquerel discovered radioactivity. (Read
       about these in the first few pages of chapter 2 and read about
       determination of the charge of the electron by Millikan).
This is the experimental setup that Thomson used to discover the electron. The
next slide has his model of the atom.




                                                                          Fig. 2.3, p.62
D. Rutherford’s Model
 1. Discovered three types of radiation for radioactive decay of
     certain atoms. (alpha: -particles with a positive charge, beta: -
     particles with a negative charge, and gamma:  -particles with no
     charge). He showed that -particles were He2+ ions (a bare
     nucleus with two protons and no electrons).
 2. Conducted the now famous “Gold foil experiment”. It was
     designed to test Thomson’s model of the atom. (see a later slide).
     Rutherford bombarded a thin foil of gold with alpha particles and
     expected the sea of positive charge to repel most of the alpha
     particles backward. To his surprise most of the particles passed
     through the gold foil but a few alpha particles were deflected
     directly backwards.
   3. Rutherford proposed a model of the atom with a dense
      nucleus with all of the positive charge and the electrons
      very far from the nucleus. Solar system model of the atom.

   4. Based on the number of deflections directly backward
      versus the number of particles passing through the foil he
      calculated that the atom’s diameter is 104 - 105 times larger
      than the nucleus.

E. Later we will discuss the Bohr and the Quantum Mechanical
Models
Fig. 2.2b, p.61
Fig. 2.6, p.66
   II. Fundamental Atomic Particles and Atoms
       A. Particles
Name          Symbol        Mass        Charge           Discovery
1. Proton     p+       1.00728 amu      1+               1886
2. Electron e-        0.0005486 amu     1-               1897
3. Neutron n           1.00867 amu      0                1932
   Note: amu = atomic mass units.

      B. Size
         1. Atoms have a radius of 30-300 pm = 0.3-3 Angstroms
         2. Nucleus is 104 - 105 times smaller than the diameter of the
             atom.

      C. Atomic Number: Z = the number of protons, determines the
         identity of the atom.
      D. Mass number = A = the sum of the number of protons and
         neutrons
     E. Designation of nuclides: X = the symbol of the element.
        1. ZAXcharge
        2. Isotopes are nuclides with the same Z but different A
                 Hydrogen        1H = H
                                1
                 Deuterium        2
                                1 H=D
                 Tritium          3
                                1 H=T
        3. Alpha particle is 24He2+

     35Cl        17 p+ 18 n and 17 e- (natural abundance is 75%)
17
     37Cl        17 p+ 20 n and 17 e- (natural abundance is 25%)
17


Examples:        Protons        Neutrons       Electrons
     19F            9              10                9
   9
      41Ca          20             21                20
   20
     19 1-
   9 F              9              10                10
      41   2+
   20 Ca            20             21                18
 III. Atomic Masses
   A. The mass of an individual atom was too small for early chemist to
   measure, but they could measure relative masses. For example, burn
   a known mass of Mg in the presence of oxygen and find the final
   mass of MgO. By subtraction the mass of combined oxygen is
   known. It was found that
       Mass of magnesium/mass of oxygen = 1.51/1.0
   Repeat this experiment with hydrogen and oxygen and find that the
       Mass of oxygen /mass of hydrogen= 16.0/1.0
This shows us that Mass of magnesium /mass of hydrogen= 24.3/1.0
Repeat this experiment many times to find that hydrogen has the lowest
mass, so the early chemist assigned hydrogen a mass of 1.0. Now we
use 12C as the standard and it is assigned a value of 12.0 amu.

       B. Don’t use amu in the laboratory because we measure in grams
 so the number under the symbol on the periodic table is the atomic
 mass in grams.
C. Average Atomic Mass. The Average Atomic Mass is a weighted
    average of all nuclides of an elements (all isotopes) See page 72
    for an example using boron and page 73 for atomic bromine.
        1. Hydrogen consists of 99.9855% of 11H (1.007825 amu)
        and 0.0145% of 12H (2.014102 amu). Calculate the average
        atomic mass of hydrogen.      DO

Note: The fractional abundance would be 0.999855 and 0.000145.
   The fractional abundances must add to one.
ANS:
      2. Neon consists of 20Ne (19.992440 amu) and 22Ne (21.991386
amu). Calculate the fractional abundance for each isotope of Ne.

ANS:
D. Formula mass is the sum of the atomic masses for all
elements in the formula. Calculate the formula mass for
   1. CO2 (carbon dioxide)       DO
           ANS:
   2. BaCl2-2H2O (barium chloride dihydrate) DO
           ANS:
IV. Symbols and Formulas
  A. Symbol: short hand notation for an element. Either a capital
     letter or a capital letter followed by a lower case letter.

  B. Formulas: Indicates the kinds and number of atoms present in
      a compound. It also gives the stoichiometric ratio.
  1.      CO            1 : 1 stoichiometry
  2.      CO2           1 : 2 stoichiometry
  3.      H2SO4         2 : 1 : 4 stoichiometry

  C. Nature of Elements or Compounds at room temperature
     1. Elements may exits as
        a. Single atoms (He, Hg, Ar …)
        b. Aggregrates of atoms: metals and a few non-metals
                C, Si, Fe, Cu, Na, Rb, Se, etc.
        c. Two or more atoms
            H2, N2, O2, F2, Cl2, Br2, I2 (seven diatomic elements)
            P4, S8, etc.
2. Compounds may exist as

   a. Individual molecules (generally covalent bonded
   compounds which are two or more non-metals
   combined). CO2, C3H8, SO3, NH3, C2H5OH etc.

   b. Collections of ions (ionic bonded compounds
   which are one metal and one non-metal). NaCl,
   K3PO4, etc.
V. Mole Concept.
    A. The early chemists were able to measure relative
    masses and learned that H was the lightest element, an
    atom of oxygen was 16 times heavier, Mg was 24.3 times
    heavier, etc. Amedo Avogadro asked the questions
       How many Mg atoms in 24.3 g of Mg?
       How many O atoms in 16.0 g of atomic oxygen?
       How many H atoms in 1.00 g of atomic hydrogen?


    He never found the answer, but it is (his number) 6.022 x 1023
    atoms. Learn this number. He realized that this was the
    same number for all elements, their atomic mass in grams
    must contain 6.022 x 1023 atoms. This gives us many unit
    factors:       12.0107 g C = 6.022 x 1023 atoms of carbon
                   32.065 g S = 6.022 x 1023 atoms of sulfur
Note: He called this very large number a mole.

B. Molar Mass: The mass in grams of 1.00 mole of
molecules or formula units. Often called molecular weight.
               1.00 mole NaCl = 58.44 g
      1.00 mole C2H5OH = 46.07 g
      1.00 mole P4 = 123.9 g


 C. Conversion Factors can be generated from these
    relationships.
    1. Atoms:
        1.008 g H = 6.022 x 1023 H atoms = 1 mole H
        4.002 g He = 6.022 x 1023 He atoms = 1 mole He
        55.845 g Fe = 6.022 x 1023 Fe atoms = 1 mole Fe
 2. Molecules
    1 mole of H2O molecules = 6.022 x 1023 H2O
    molecules = 2 mole H atoms = 1 mole O atoms = 18.0
    g of H2O.
    1 mole C2H5OH = 2 mol C = 6 mol H = 1 mol O =
    6.022 x 1023 C2H5OH molecules = 46.07g of C2H5OH


D. Unit Factors: All of the above can be made into unit
factors. (Do example)
E. Sample calculations DO ALL CALCULATIONS
  1. What is the mass in grams of one Na atom?
ANS =
  2. Calculate the number of Na atoms in each of the following.
        a. 0.263 mol of Na
        b. 37.5 g of Na


               a. ANS
               b. ANS
3. How many moles of carbon in 2500. tons of coal?
       ANS:
4. How many atoms of C in 2500. tons of coal?
       ANS:
VI. Periodic Table: Classification of elements

       1. Based on general properties
           a. Metals: left side of the periodic chart (see front cover
                    of text)
           b. Nonmetals: right side of the chart (yellow)
           c. Metalloids: Properties intermediate between metals
                    and nonmetals (see front cover-green)
        2. Based on Periodic trends
           a. Main Group: IA –VIIA
           b. Transition: B Group
           c. Inner Transition
                    i. Lanthanide series
                    ii. Actinide series
        3. Periods-horizontal rows--periodic change in properties.
        4. Families or Groups-columns (similar properties).
p.80
p.79a
p.79b
VII. Overview of the elements. (read for content)
     A. Group 1A: Alkali metals: All are very reactive and when placed in water
from H2(g) and an aqueous solution of the metal hydroxide (basic solution).
Alkaline solution is an older name for a basic solutions, hence the name alkali
metals.
      B. Group IIA: Alkaline earth metals: Most react readily with water to form
H2(g) and an aqueous solution of the metal hydroxide. They are abundant in the
earths crust (Ca is #5 and Mg is #7) an calcium carbonate is limestone and is found
in sea shells, etc.
      C. Group IIIA: metals and metalloids (B). Al is the most abundant metal on
earth by mass.
      D. Group IVA: metals, nonmetals and metalloids are in this group. The most
important is carbon which is found in at least different three structural forms called
allotropes (graphite, diamond and C60--see the next slide).
     E. Read in the text the information about Groups V - VIIIA and the
Transitions Elements.
Allotropes of Carbon: Graphite,
diamond, and Buckyball = C60

				
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