1. Unit 3, Part 1: Atomic Structure & The Periodic Table

2. 3.1 Early Models of the Atom
What are Atoms?
4th century B.C. Greek philosopher Democritus stated the universe was made of invisible units called atoms (atom “unable to be divided”).
Believed movements in atoms caused changes observed in matter.

3. Late 1700’s
French Chemist, Antoine Lavoisier established the Law of Conservation of Matter.
Matter can not be created or destroyed… only change forms
Joseph Proust, later established the Law of Constant Composition: the compound always contains the same elements in the same proportions by mass.

4. In 1803 John Dalton proposed an atomic theory:
Each element is composed of extremely small particles called atoms .
Atoms of the same element are exactly alike .
Every compound always has the same ratio and kinds of atoms .
A chemical rxn is a rearrangement of atoms; they are not created or destroyed .
There are some exceptions to Dalton’s theory; however they still are the basis for understanding Chemistry.

5. 3.2 Discovering Atomic Structure
Dalton and contemporaries thought the atom was like a marble… small, hard and round… but couldn’t explain why atoms from other element behaved differently.
The English Chemist, Michael Faraday, in 1839, showed atoms contain electrical charge.

6. Electricity
American, Benjamin Franklin, experimented with electricity (think kite and key experiment). Franklin was able to determine the 2 charges an object has and it was he who coined the names “positive” and “negative” charge. Opposite charges attract and like charges repel.
But where do these charges come from? What are they?

7. Cathode Rays and Electrons
Cathode Ray Tube (CRT) - A battery is connected to a tubing of partially evacuated glass. The glass is lined with fluorescent material, current flows to the ends of the tube. The end connected to the (-) terminal of the battery is called the cathode and the other is the anode (+). A stream of radiation flows from the cathode to the anode when the battery is turned on.

8. At the end of the 19th Century we knew:
The cathode ray could spin a small paddle wheel = actual a stream of particles.
Magnet deflects ray as expected for a (-) charge = ray made of (-) charged particles.

9. Electrons
English physicist, J.J. Thompson ( ), determined the ray WAS made (-) particles by allowing the ray to pass through a hole in the anode and then through a magnetic field. He determined that the negative particles emanate from the cathode, they had structure and he named them electrons.
American physicist, Robert Millikan ( ), was able to measure the charge of an electron. He sprayed oil and used X-rays to give the oil a negative charged. Then measured how different magnetic charges changed the rate the oil fell. He calculated the mass of the e- to be 9.11 X grams.

10. Radioactivity
French physicist, Henri Becquerel ( ) , accidentally placed uranium (U) on unexposed film. The U had produced an image suggesting it was emitting some type of radiation (the spontaneous emission of particles) and that elements could be radioactive.
His associates, Marie Sklodowska Curie and her husband Pierre, discovered two more radioactive elements, radium and polonium.
Radioactivity was key to understanding the atom, as a whole.

11. 3 Types of Radiation
New Zealander, Ernest Rutherford passed the cathode ray of a radioactive substance between two charged plates. The ray split- part of the beam was deflected towards the (-) plate (alpha radiation), another was deflected towards the (+) plate (beta radiation) and a third passed straight through undisturbed (gamma radiation).
Alpha particles 2+ charge
Beta particles 1- charge
Gamma rays have no charge
This experiment demonstrated that the atom was much more complex than previously thought.

12. The Nuclear Atom
Thompson showed us that atoms had electrons, but that doesn’t explain why atoms are electrically neutral (they don’t have a charge). If they have electrons they must have some type of (+) charges, too.

13. Atomic Model Battle!
Thompson suggested that the atom was like plum pudding. The e-s were spaced evenly throughout the atoms (+) interior, the way the plums were distributed through the pudding.
(You can also think of chocolate chip cookie dough- the dough is the positive interior while the chocolate chips are the e-s).

14. Alpha Scattering
1909 The Alpha-scattering Experiment by Rutherford- A beam of high-speed alpha particles bombarded a thin sheet of Au foil. Most of the alpha particles went straight through the foil but a small portion DID deflect. And they would scatter in every direction possible. Why??

15. Atomic Model Battle! Cont.
Rutherford’s data suggested this was not true. He determined that there must be an (+) core in the atom. He named this the nucleus. The particles which went straight through suggested the atom was mostly empty space.

16. 3.3 Modern Atomic Theory The Structure of the Atom
We now know that atoms can be divided into many different subatomic particles. For example:
Nucleus- the center of the atom, very dense, contains protons and neutrons, has an overall positive (+) charge.
Protons- positively (+) charged subatomic particle.
Neutrons- neutral (not charged) subatomic particle.
Electrons- negatively (-) charged subatomic particle that has very little mass. Creates a “cloud” that encircles the nucleus.
Rutherford thought of it as like a mini solar system: the nucleus in the middle encircled by electrons.

17. Atomic Numbers
Englishman, Henry Mosley ( ) discovered that each element had a unique amount of positive charge. This concept leads to the understanding of why atoms of different elements are unique. The identity of the atom comes from the number of protons contained in the nucleus.

18. Atomic #
Atomic # (Z)- the # of protons in the nucleus of an atom. (Always a whole #)
Ex. O-8 = 8 protons in the nucleus
C-6 = 6 protons in the nucleus
Pb-82= 82 protons in the nucleus
Now, individual atoms are electrically neutral… meaning they have an equal # of protons as electrons.
O has 8 protons (we know this by its atomic number) and 8 electrons.
C has 6 protons and 6 electrons
Pb has 82 protons and 82 electrons.
WHAT HAPPENS WHEN THEY ARE NOT EQUAL??

19. Ions Charge on Ion = # of protons - # of electrons
The atom can gain or lose electrons. This is an ion.
(NOT PROTONS- if it change the # of protons you change the element/atom.)
If you gain an electron you get a negative ion.
If you lose an electron you have a positive ion.
Charge on Ion = # of protons - # of electrons
Magnesium (Mg) can loose 2 electrons making it a positive ion- written as:Mg2+
O will gain 2 electrons making it a negative ion. It is written as such: O2-
Atoms that lose electrons, like Mg, are called cations. Atoms that gain electrons, like O, are called anions.

20. Atomic Mass #
Atomic mass # (A) - the # of protons + neutrons in an atom.
Because the bulk of the atom’s mass is provided by the protons and neutrons, we only consider their masses when calculating the atomic mass #.
Ex: O has 8 protons and 8 neutrons, so A= 16.
Dalton said that every atom of the same element is exactly alike… NOT SO. They do have the same # of protons; however, they do not necessarily have the same # of neutrons!

21. Isotopes
This means that different atoms of the same element may/will have different masses.
isotopes (atoms having the same # of protons but different # of neutrons).
Some isotopes are more common than others.
3 isotopes of H (why A = ).
Protium (only has one proton in the nucleus) A=1
Deuterium (1 proton + 1 neutron) A=2
Tritium (1 proton + 2 neutrons) A=3

23. Determining The # of Neutrons
Calculating the # of neutrons in an atom:
General Formula:
A – Z= # of neutrons.
Ex: calculate the # of neutrons in the isotopes of Cl (Cl 35 and Cl 37)

24. The Mass of an Atom
The mass of a single element is extremely small (in the one trillionth of a billionth range) and is very difficult to work with. So instead we express the mass of atoms in atomic mass units (amu).
One amu is equal to 1/12th of the mass of a carbon-12 atom. This isotope has exactly 6 protons and 6 neutrons so the mass of each has to be about 1.0 amu.
1amu = 1/12 (mass of 126C atom) = 1.66x10-24 g

25. Average Atomic Mass
The atomic mass of an element is often listed as the average atomic mass as found in nature. This is a weighted average of the isotopes for that particular element. The more commonly found isotopes have a greater effect on the averages mass than the more rare isotopes.
Ex: Cl 24% Cl-37 and 76% Cl-35, thus the average atomic mass (35.45 amu) is much closer to 35 than 37.