Modern Atomic Theory

Daltons Atomic Theory, circa 1808

In 1808 an English schoolteacher proposed the following explanation of matter. Since then we have learned more about the atom and now have a slightly different 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.
  • In chemical reactions, atoms are combined, separated, or rearranged.
  • QUESTION: Which of the above premises have changed? How should the wording above be modified to explain our new information?

    JJ Thomsons Cathode Ray Experiment, 1897

    In the late 1800s many scientists were experimenting with electricity to determine its nature and what it was.

    The experiment: Thomson demonstrated that electricity in the cathode ray tube could be deflected by positive and negative charges (towards (+) and away from (-)). He also showed that the ray would actually spin and move a paddlewheel inside the tube.

     

    The conclusions:

    Electricity is composed of particles of matter (electrons).

    Electrons are negatively charged.

    Electrons are very small.

    Robert A. Millikans Oil Drop Experiment, 1909

    Millikan showed that the mass of an electron is in fact 1/2000 of the simplest type of hydrogen atom. His experiment also showed that the charge to mass ratio was very high. This means that although the mass of an electron is much smaller than that of a Hydrogen atom it has about the same size charge as a negative ion of hydrogen.

     

    Enerst Rutherfords Gold Foil Experiment, 1911

    The Experiment: Rutherford shot alpha particles at a thin gold foil. As he expected most went right through the foil. Unexpectedly however, some were deflected slightly, and a few even bounced back. This was startling information.

    The Conclusions:

    The atom is composed mostly of empty space.

    There is a small but very dense central core known as a nucleus.

    The Nucleus has positive charge.

    Neils Bohr's Light Experiment, 1913

    Bohr experimented with how gases absorb and re-emit light in only specific colors (energies). His experiment showed that the electrons orbit the nucleus in definite orbitals of specific energy.

     

    Modern Atomic Theory - Quantum Theory.

    Many scientists have contributed to the atomic theory since those listed above. Some notable names are Einstein, DeBroglie, Schrodinger, and Heisenberg. Quantum theory has shown us that the electrons although they are particles also exhibit properties of waves. Now we think of the atom as a nucleus that is surrounded by probability clouds. The clouds represent the most probable locations of electrons. We still refer to these clouds as orbitals. The size and shapes of the orbitals may be calculated mathematically by using the equations for the waves.

     

    Anatomy of an Atom

    Atomic Structure

    The following table summarizes the parts of the atom and their arrangement within the atom.

    Particle Symbol Mass Charge Location
    neutron no ~1 amu. (0) nucleus
    proton p+ ~1 amu. (+) nucleus
    Electron e- ~ 1/2000 amu. (-) electron cloud

     

    Atomic Number (Z)

    The atomic number is the number of protons in an elements atoms. The atomic number is the sole factor that determines the identity of the atom.

    Atomic Mass and Mass Number

    The atomic mass is the mass of a single atom. Because the mass of an atom is primarily due to the nucleons (protons and neutrons) we assign a mass number to an atom. The mass number is simply the sum of the protons and neutrons and therefore an approximation of the mass of the atom. Mass numbers are always whole numbers.

    Isotopes are atoms of the same element that differ in number of neutrons. For example, Carbon-12 and Carbon-14 are two isotopes of carbon. They differ in that C-12 has 6 neutrons and C-14 has 8 neutrons. The number of protons and electrons in each are identical.

    Isotopes are symbolized in two ways

    Carbon-12, the name followed by the mass number of the isotope.

    , the symbol of the element with the mass number and atomic number preceding it. This symbol may also be shortened to the mass number and the symbol 12C. This is OK because the atomic number is unique to the element and does not need to be stated.

    Ions are atoms of an element that are charged due to the loss or gain of electrons.

    Ions are symbolized by there symbol followed by their charge written as a superscript. I.e. H+, Mg2+, Br-.

    Relative Atomic Mass

    Because the actual mass of an atom is very small (An atom of Carbon-12 has a mass of 1.660x10-27kg.) it is easier to express the mass of an atom relative to other atoms. The carbon-12 atom has been chosen as the arbitrary standard and is given the exact mass of 12 atomic mass units (amu).

    The average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element in their normal distribution.

    The average atomic mass is calculated by taking a naturally occurring sample of an element and averaging the weights of all the atoms in the sample.

    Here is an example of how to calculate the average atomic mass of carbon:

    To do these problems you need some information: the exact atomic weight for each naturally-occurring stable isotope and its percent abundance. These values can be looked up in a standard reference book such as the "Handbook of Chemistry and Physics."

    Example #1: Carbon

    mass number of isotopes

    exact weight (amu)

    percent abundance

    12

    12.000000

    98.90

    13

    13.003355

    1.10

    To calculate the average atomic weight, each exact atomic weight is multiplied by its percent abundance. Then, add the results together, divide by the number (100 atoms in this case) and round off to an appropriate number of significant figures.

    This is the solution for a sample of 100 carbon atoms:

    (12.000000 amu) (98.90) + (13.003355 amu) (1.10) = 1201.1

    divide by 100 to get the average = 12.011 amu

     

     

    Mole Concept: Counting Atoms

    Definition of Mole, Avogadros Number

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

    This number is known as Avogadro's number and is equal to 6.022 1367x1023. For most purposes Avogadro's number is rounded to 6.02x1023.

    Avogadros number is used as a conversion factor between the number of particles and the number of moles.

    Molar Mass

    Molar mass is defined as the mass of exactly one mole of a substance.

    The molar mass in grams is numerically equivalent to the average atomic mass for any element.

    The molar mass of a compound is equal to the sum of the molar masses for each element in the element.

    The molar mass of a substance is used as a conversion factor between the mass of a substance and the number of moles.

    Atom/Mole Conversions

    To convert from moles to number of atoms:

    # moles x 6.02x1023 atoms = #atoms
      1 mole  

    Example:

    23.5 moles x 6.02x1023 atoms = 1.41x1025atoms
      1 mole  

    Notice how the units cancel to leave you with the correct units.

    To convert from number of atoms to moles:

    # atoms x 1 mole = #moles
      6.02x1023 atoms  

    Example:

    3.50x1024atoms x 1 mole = 5.81 moles
      6.02x1023 atoms  

    Notice how the units cancel to leave you with the correct units.

     

    Mass/ Mole Conversions

    To convert from moles to mass:

    # moles x molar mass (g) = mass (g)
      1 mole  

    Example 1:

    15.9 moles C x 12.011 g. C = 191 g. C
      1 mole C  

    Example 2:

    3.60 moles Mg x 24.305 g. Mg = 87.5 g. Mg
      1 mole  

    Notice how the units cancel to leave you with the correct units.

    To convert from mass to moles:

    Mass (g) x 1 mole = mole
      Molar mass(g)  

    Example 2:

    43.7 g. O x 1 mole O = 2.73 mole O
      15.9994 g. O  

    Example 2:

    19.2 g. Br x 1 mole Br = 0.240 mole Br
      79.904 g. Br  

    Notice how the units cancel to leave you with the correct units.