Chapter 4, Section 1 - The Development of Atomic Theory                                                                                                                                           

Democritus - 4th Century BCE
  • Suggested that the universe was made of tiny, indestructible spheres called "atomos".
John Dalton - 1808
  • Said that all atoms of a given element are exactly alike.
  • Atoms of different elements can join to form compounds.
Law of Definite Proportions
  • A chemical compound always contains the same elements in exactly the same proportions (by weight or mass).
    • Example - A water molecule always has two hydrogen atoms and one oxygen atom.
      • By mass, water is always 11% hydrogen, 89% oxygen.
J.J. Thomson - 1897
  • Discovered negatively charged particles that came from inside atoms - electrons!
  • Plum-Pudding Model
    • Electrons are spread throughout the inside of the atom - like bits of plum in plum pudding.
Ernest Rutherford - 1909
  • Experiment - Used radioactive material to shoot alpha particles at a thin piece of gold foil.
  • Results...
    • Most passed straight through.
    • Some were slightly deflected.
    • Some bounced back!
  • Conclusions...
    • Most of the atom's mass is concentrated in a tiny nucleus at the center of the atom.
    • The electrons orbit the nucleus.
  • We understand that the nucleus is made of...
    • Protons (positively charged particles)
    • Neutrons (particles with no charge)

Chapter 4, Section 2 - The Structure of Atoms                                                                                                                                                               

Atoms are made of subatomic particles...

 Particle ChargeMass Location in the Atom
 Proton +11 amu   In the nucleus. 
 Neutron 01 amu   In the nucleus.
 Electron -1 1/2000 amu  Orbiting the nucleus.

Protons determine the element.  
  • Each element has a specific number of protons.  
    • Example - All carbon atoms have six protons.
Atoms have no overall charge. 
  • The number of protons in an atom equals the number of electrons, so the positives and negatives cancel out.
  • If an atom gains or loses electrons, it becomes charged and is now an ion (an atom with a charge).
The electric force holds the atom together.
  • Positive and negative charges attract each other - this is called the electric force.
  • The charges of protons and electrons hold the atom together.
Atoms of the same element always have the same number of protons, but can have different numbers of neutrons.
  • Atomic Number (Z) - The number of protons in the nucleus of an atom.
    • Each element has its own specific atomic number.  This number doesn't change!
  • Mass Number (A) - The total number of protons and neutrons in the nucleus of an atom.
    • The number of neutrons can vary from atom to atom of the same element.
    • Isotopes - Atoms of the same element (same number of protons) with different numbers of neutrons.
      • Some isotopes are more common than others...
        • Protium - 1H (hydrogen with one proton) - most common isotope of hydrogen.
        • Deuterium - 2H (1 proton, 1 neutron) - 1 out of 6,000 hydrogen atoms on Earth.
        • Tritium - 3H (1 proton, 2 neutrons) - very rare.  Unstable and decays over time.
          • Radioisotopes - give off radiation and decay into other isotopes over time.
  • To find the number of protons, neutrons or electrons in an atom...
    • # Protons = Atomic Number
    • # Neutrons = Mass Number - Atomic Number
    • # Electrons = Atomic Number
Atoms have mass - but their mass is very small.
  • Unified Atomic Mass Unit (u) - A unit of mass that is 1/12 the mass of a carbon-12 atom (with 6 protons and 6 neutrons).
    • Also called atomic mass unit (amu).
    • The mass of one proton or one neutron is about 1u.  (Electrons are too small to affect atomic mass significantly.)
  • Average Atomic Mass - The average mass of the different isotopes of an element, as they are found in nature.
    • Chlorine has two isotopes: 35Cl and 37Cl.  The 35Cl isotope is more common, so the average atomic mass is 35.453u.

Mole (mol) - The SI (metric) unit used to measure the amount of a substance of very small particles.
  • 1 mol = 602,213,670,000,000,000,000,000 particles
    Avogadro's Number
    • Avogadro's number = 6.022 x 1023
      • The number of atoms in 12.00g of carbon-12.
  • Molar Mass - The mass (g) of 1 mole of a substance.
    • An element's molar mass (g/mol) = its average atomic mass (u).
      • 1 mol of Cl = 35.435g  (The average atomic mass of Cl is 35.453u.)
    • A compound's molar mass = the sum of the molar masses of the atoms in the compound.
      • H2O = 2 hydrogen atoms, 1 oxygen atom
        • Oxygen's atomic mass = 16.00u     (Molar mass = 16.00 g/mol.)
        • Hydrogen's atomic mass = 1.01u     (Molar Mass = 1.01 g/mol.)
          • H2O's molar mass = (2 x 1.01 g/mol) + 16.00 g/mol = 18.02 g/mol

    Page 127 #2, 4, 5, 8-11, 13-16.

Chapter 4, Section 3 - Modern Atomic Theory                                                                                                                                                               

Modern Models of the Atom
  • Electrons exist only in specific energy levels.  (Not between levels.) - Niels Bohr (1913)
  • Electrons act more like waves than like particles.
  • It is impossible to know both the position and the speed/direction of an electron.
    • We can know the area around the nucleus where electrons are most likely to be.  These areas are called orbitals.

Electron Energy Levels
  • Electrons can have different amounts of energy, so they can exist in different energy levels.
  • The number of energy levels that are filled depends on the number of electrons.
  • Valence Electrons - Electrons in the outer energy level.
    • Valence electrons determine an atom's chemical properties, because they are available to interact with other atoms.
  • Each energy level contains orbitals.  There are four kinds of orbitals: s, p, d, f.

    • Each orbital can hold two electrons.

Electron Transitions
  • Electrons do not exist between energy levels, but they can jump from one level to another.
    • They do this by gaining or losing energy.
      • Gain energy to move to a higher energy level.
      • Lose energy when moving to a lower energy level.
    • Ground State - An electron's lowest energy state.
    • Excited State - A higher energy level that electrons move to when they absorb photons.
    • Photon - A unit (or particle) of light.
      • Electrons absorb photons to move to higher levels.
      • Electrons give off photons when moving to lower levels.
    • Photons of different colors of light have different amounts of energy.  The photon's energy determines which energy level the electron will jump to.  (See the Bohr Atom Interactive link below for a demonstration.)
      • Since different elements have different structures, they absorb and emit (give off) different photons of light.
        • So, each element gives off its own unique set of colors.

  • Bohr Atom Interactive - This is an activity from my Astronomy class.  It shows how different wavelengths of light make the electrons jump to different levels.  Once you are there, you'll need to click on "The Bohr Atom (51.0K)" on the top to start this interactive.
  • Rutherford's Gold Foil Experiment Animation - This let's you see what happened at the molecular level, as well as at our scale.