Modern Chemistry - Chapter 3

Atoms: The Building Blocks of Matter

 

  1. The Atom: From Philosopical Idea to Scientific Theory
    1. Aristotle v Democritus - philosphy of each
    2. Foundations of Atomic Theory
      1. up to 1790 - assumptions and questions
      2. 1790 - quantitative approach
      3. led to
        1. Law of Conservation of Mass - statement
        2. Law of Definite Composition (Proportions) - statement
        3. Law of Multiple Proportions - statement
        4. for the following two compounds - the same two elements are present in each

          amt of carbon

          amt of oxygen

           

          carbon monoxide
          1 gram
          1.33 grams
          carbon dioxide
          1 gram
          2.66 grams

          Ration of oxygen in the two compounds is 1.33:2.66 or 1:2

      4. Chemical Reaction - definition
    3. Dalton's Atomic Theory
      1. 1808
      2. which laws it explained
      3. postulates of the atomic theory
      4. how this theory relates to the law of conservation of mass - see figure 3-2 page 67
      5. how this theory relates to the law of definite proportions - see figure 3-3 page 67
    4. Modern Atomic Theory
      1. which postulates have been modified
      2. process is to modify where possible as opposed to discarding

        Homework 3.1

  2. The Structure of the Atom
      1. two regions in the atom
      2. what particles are contained in each region
      3. meaning of subatomic particles
    1. Discovery of the Electron
      1. electricity and matter
      2. cathode ray tube (Crooke's Tube) - figure 3-4 page 70
    2. Cathode Rays and Electrons
      1. why name cathode ray tube
      2. link: http://ganymede.nmsu.edu/tharriso/ast301/class20.html
      3. link: Crooke's Tube Maltese Cross
      4. testing the hypothesis they found:
        a) an solid object placed between the cathode and anode cast a shadow on the glass;
        b) a paddle wheel placed on rails between the electrodes rolled along the rails from the cathode toward the anode (figure 3-5 page 71).
      5. conclusion from using the Maltese Cross and paddle wheel in cathode ray tube
      6. Further experiments showed:
        a) cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have a negative charge
        b) deflected by negative charged objects
      7. Conclusion of above experiment in item 6
      8. Joseph Thomson in 1897: back up conclusions and also measured the charge to mass ratio;
        ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube.
      9. Conclusion: all cathode ray particles are identical and negatively charged.
    3. Charge and Mass of the Electron
      1. Two major conclusions of Thomson's work
      2. Robert Millikan in 1909 determined the mass of the electron and later the charge of the electron
      3. Conclusions of Millikan's work.
      4. From Cathode ray experiments - atoms are divisible
      5. The basis for more experiments also came from this work:
        a) negative charge of electron must be balanced by some positive charge since atoms are neutral
        b) because of the small mass of electrons, other particles must contribute to mass of atom
    4. Discovery of the Atomic Nucleus
      1. 1911 - Ernest Rutherford: gold foil and alpha particles
      2. Figure 3-6 page 72
      3. Rutherford's experiment
      4. three paths of particles and conclusion for each path
      5. Further pieces to the puzzle: a) if nucleus were size of marble, the atom would be the size of a fottball field - i.e. volume of the nucleus is very small compared to the size of the atom; b) electrons surrounded the nucleus like planets around the sun; could not explain motion of electrons around the nucleus.
    5. Composition of the Atomic Nucleus
      1. Three subatomic particles in atoms:
        subatomic particles
        particle
        charge
        mass
        relative mass
        mass number
        proton
        +1
        1.673 x 10-27 kg 1.007 276 AMU

        1

        neutron
        none
        1.675 x 10-27 kg 1.008 665 AMU
        1
        electron
        -1
        9.109 x 10-31 kg 0.000 548 6 AMU
        0


        1. Elements are identified by the number of protons in their nucleus - the atomic number.
        2. Each element has a characteristic number of protons in the nucleus.
        3. Table 3-1 apge 74
        1. Forces in the Nucleus
          1. like charges repel; unlike charges attract
          2. exception is two protons close to each other in the nucleus
          3. have attraction between neutrons in nucleus
          4. nuclear forces - definition
    6. Sizes of Atoms
      1. electron cloud
      2. nucleus
      3. radius of atom - definition
      4. expressed in picometers -- 1 pm = 10-12m = 10-10 cm

        Homework 3.2

  3. Counting Atoms
      1. Use basic properties of atoms to count the number of atoms of an element in a sample whose mass you know.
      2. Mole (mol) is the unit we use
    1. Atomic Number
        1. Atoms of different elements vs elements of same element - number of protons
        2. atomic number - definition
        3. how to identify atomic number on periodic table
        4. used to identify element
    2. Isotopes
      1. isotopes - definition
      2. table 3-2 page 76
      3. most elements are a mixture of isotopes
    3. Mass Number
      1. mass number - definition
      2. table 3-2 page 76
    4. Designating Isotopes
      1. Isotopes of hydrogen vs istopes of most other elements
      2. Two notations for writing istopes
      3. calculating the number of neutrons
      4. nuclide - definition
      5. Table 3-3 page 77 five different nuclides
      6. Sample problem 3-1 page 77

        sample problem 3-1
    5. Relative Atomic Masses
      1. mass of atoms in grams in too small a number
      2. relative atomic mass
      3. relative implies a standard
      4. for atoms, the standard is carbon-12
      5. units for relative masses are atomic mass units
      6. AMU - definition
      7. standard of carbon-12 has a relative mass of exactly 12 AMU
      8. relative atomic masses are determined by comparing the mass of the nuclide with carbon-12
      9. table 3-4 page 80
      10. isotopes have different physical properties because they have different masses
      11. have similar chemical properties because they have the same number of electrons which have the same energy - chemical properties depend on the number and energy of electrons
      12. table 3-1 page 74 - relative masses of subatomic particles
      13. mass number v relative atomic mass
    6. Average Atomic Masses of Elements
      1. Average atomic mass - definition
      2. Calculating a weighted average:
      3. 25 marbles x 2.00 g = 50 g
      4. 75 marbles x 3.00 g = 225 g
      5. total mass = 275 g
      6. average: 275g / 100 marbles = 2.75 g/marble
    7. Calculating Average Atomic Mass
        1. Naturally occurring copper consists of 69.17% copper-63, atomic mass of 62.929 598 AMU, and 30.83% copper-65, atomic mass of 64.927 793 AMU Calculate the average atomic mass.
        2. Given: Copper-63, 69.17%, 62.929 598 AMU
        3. Copper-65, 30.83%, 64.927 793
        4. average atomic mass = ?
        5. (percent) x (atomic mass) + (percent) x (atomic mass)... = average atomic mass
        6. (.6917) x (62.929 598 AMU) + (.3083) x (64.927 793 AMU) = 63.55 AMU
        7. Author rounds atomic mass to two decimal places before using it in a calculation.
    8. Relating Mass to Numbers of Atoms
      1. mole & Avogadro's Number & molar mass: relate mass in grams to numbers of atoms
        1. The Mole
          1. SI unit
          2. mole - definition
          3. mole is a counting unit
        1. Avogadro's Number
          1. experimentally determined
          2. 6.022 x 1023
          3. exactly 12 g of carbon-12 contains 6.022 x 1023 atoms of carbon
          4. The Avogadro's Number - definition
        2. Molar Mass
          1. Molar mass - definition
          2. Numerically equal to the atomic mass for elements
          3. we use the unit grams in problems and calculations when we are dealing with moles
          4. one mole = molar mass = Avogadro Number of particles
        3. Gram/Mole Conversions
          1. molar mass is used as a conversion factor in calculations
          2. Figure 3-11 page 82
          3. Sample Problem 3-2 page 82

            sample problem 3-2
          4. Sample Problem 3-3 page 83

            sample problem 3-3


        4. Conversions with Avogadro's Number
          1. numer of moles ---> number of atoms
          2. number of atoms ---> number of moles
          3. grams ---> moles - using molar mass
          4. moles ---> grams - using molar mass
          5. Sample problem 3-4 page 84

            sample problem 3-4

          6. Sample Problem 3-5 page 84

            sample problem 3-5

            Homework: Section 3.3

            Definitions

            A chemical reaction is the transformation of a substance or substances into one or more new substances.

          7. Law of Conservation of Mass states that mass is neither destroyed nor created during ordinary chemical or physical reactions.

            Law of Definite Composition states that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.

            Law of Multiple Proportions states that if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.

            Dalton's Atomic Theory:
            All matter is composed of extremely small particles called atoms
            Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties;
            Atoms cannot be subdivided, created or destroyed;
            Atoms of different elements combine in simple whole-number ratios to form chemical compounds;
            In chemical reactions atoms are combined, separated, or rearranged.

            Nuclear Forces are short-range proton-neutron, proton-proton, and neutron-neutron forces holding the nuclear particles together.

            The radius of atom is distance from center of nucleus to outer portion of electron cloud.
            It is measured by taking 1/2 the distance between two adjacent nuclei of the same element.

            The Atomic Number (Z) of an element is the number of protons in the nucleus of each atom of that element.

            Isotopes are atoms of the same element that have the same atomic number but different masses.

            Mass number is the total number of protons and neutrons in the nucleus of an isotope.

            Nuclide is a general term for any isotope of any element.

            One atomic mass unit (AMU) is exactly 1/12 the mass of a carbon-12 atom, or 1.660 540 x 10-27kg.

            Average Atomic Mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.

            A mole is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12.

            Avogadro's Number is the number of particles in exactly one mole of a pure substance.

            The mass of one mole of a pure substance is called the molar mass of that substance.