1. Chapter 3
Atomic Theory

2. Laws Democritus-called basic particle –atom meaning indivisible
Law of Conservation of Mass – mass is neither created or destroyed
C + O2  CO2
Law of Definite Proportions – a chemical compound contain the same proportions by mass regardless the size of the sample or the source (a given compound is always composed of the same combination of atoms)

3. Laws (cont’d)
Law of Multiple Proportions- 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 is always a a ratio of small whole numbers
Example: CO2 and CO each compound has 1.0 gram of carbon while in carbon dioxide has a oxygen amount of 2.33 and CO has a 1.33 they have a ratio of 2 to 1.

4. Dalton’s Theory
All matter is composed of small particles called atoms.
Atoms of a given element are identical in size, mass, properties. (not always!)
Atoms cannot be subdivided, created or destroyed (not true  )
Atoms of different elements combine in simple ratios to form compounds.
In chemical rxns, atoms are combined, separated, or rearranged.

5. Structure of the Atom
Atom-smallest particle of an element that retains the chemical properties
Nucleus-very small center region
It contains the protons (positively charged particles) and neutrons (neutral).
Electrons are found outside the nucleus in a theorized electron cloud and are negatively charged.

6. Discovery of the Electron
Electrons were discovered by cathode-ray experiments by Thomson which founded their existence and negative charge. Millikan expts. showed the negative charge of electrons and that their mass was very small.
9.109 x kg or 1/1837 the mass of a hydrogen atom.
Two inferences were made about atom structure:
Since electrons are negative and atoms are neutral, there must be a positive charge.
Since electrons have such small mass, atoms must have other particles that account for the mass of an atom.

7. Thomson Experiment
Cathode Ray

8. Discovery of the Nucleus
Atomic nucleus Discovery
Rutherford—gold foil experiment in which he bombarded alpha particles at gold foil.
Some particles were deflected back and he concluded that a dense packed bundle caused it. He called it a nucleus.
Nucleus has:
Protons with mass x kg
Neutrons with mass of x kg

9. Rutherford Experiment
Rutherford Expt

10. Inside the Atom
Number of protons determine the identity of the element.
Nuclear forces hold the nucleus together
Atomic Number (Z)—number of protons of an element

11. Isotopes
Isotopes- atoms of the same element but have a different number of neutrons and therefore a different mass.
Three Isotopes of Hydrogen
Protium – one proton only (99.985%)
Deuterium- one proton and one neutron
Tritium-radioactive-one proton, two neutrons, and one electron

12. Writing Isotopes Tin (Sn) has the most stable isotopes, 10.
How can I (ID) an isotope?
By its mass number
Mass number-total number of protons and neutrons
Most are either wrote in:
Hyphen format Uranium-235
Nuclear format 235
92U
To Find the number of neutrons for any element:
Mass number-Atomic Number = number of neutrons
235-92= 143 neutrons
If you are looking for the number of electrons, the number of protons=number of electrons if the atom is neutral

13. Determining Atomic Mass
The standard use to govern units of atomic mass is carbon-12 which equals 1 amu.
Average Atomic Mass – weighted avg. of all the naturally occurring isotopes of an element
Calculated Average Atomic Mass depends on the mass and percentage of abundance.
Multiply percent (convert to decimal) X amu and add up all isotopes.

14. Example of Determining Atomic Mass
Copper has two isotopes
One Isotope is prevalent 68.17% in nature; the other 30.83%
For copper isotopes, x amu x amu = amu

15. Vocabulary
Mole-SI unit for the amt of substance. It is the amount of a substance that contains as many particles as there are in exactly 12 grams of carbon-12. (Counting unit)
Avagadro’s number-number of particles in exactly one mole of a pure substance.
Molar Mass-mass of one mole of a pure substance (g/mol)