1. The Building Blocks of Matter
ATOMS
The Building Blocks of Matter

2. History of the Atom
Around 400 B.C. Democritus was the first person to use the term atom.
Atom is from the Greek word “indivisible”
1704 Isaac Newton proposed a mechanical universe with small solid masses in motion
1808 John Dalton proposed his atomic theory based upon measurable properties of mass

3. Properties of Mass
There are several basic laws which drive Atomic Theory
Law of Conservation of mass
Matter is never created nor destroyed during ordinary physical and chemical changes
Law of Definite Proportions
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
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

4. Dalton’s Atomic Theory
In 1808, an English schoolteacher named John Dalton proposed an explanation for the different properties of mass.
Reasoned that elements were composed of atoms and combine to form compounds

5. Daltons 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

6. Dalton’s WRONG!!!!
We now know that atoms can be divided into smaller particles.
An element can have atoms of different masses, called isotopes.

7. Structure of the Atom Nucleus Protons Neutrons Electrons 75047
Dense, positively charged center of the atom
Composed of protons and neutrons
Makes up the mass of an atom
Protons
Positively charged
Neutrons
No charge
Electrons
Negatively charged, located outside the nucleus
Makes up the size of the atom
75047

8. Electrons
In 1897, J.J. Thomson, using cathode ray tubes, found negatively charged particles called electrons
In 1909, Robert Millikan measured the charge of the electron
Oil drop experiment

9. Plum Pudding Model Proposed by Thomson
Negative electrons were spread throughout the positive charge of the atom
Like plums in plum pudding, thus the name…get it
He was wrong!!!

10. Glowing Matter
Around 1900, Max Planck explained the phenomenon of hot glowing matter.
Explained this using discrete units of energy he called “quanta” or “quantum”
This idea would later be used to “correctly” place electrons in an atom.

11. Behold…The Nucleus
In 1911, Ernest Rutherford found a dense, positively charged bundle of matter he called the nucleus
Used the Gold Foil Experiment to find the nucleus
Nuclear forces are short range forces that hold protons and neutrons very close together

12. Rutherford’s Model
Ernest Rutherford used his research to develop his own model of the atom
Dense positively charged center, nucleus
Electrons traveled around the nucleus in empty space
He was Wrong!!!!

13. To see if someone got the correct model of the atom!!!
STAY TUNED
To see if someone got the correct model of the atom!!!

14. The Power of the Proton
In 1914, H.G.J. Moseley using x-ray tubes, determined the charges on the nuclei of most atoms.
He wrote "The atomic number of an element is equal to the number of protons in the nucleus".
This work was used to reorganize the periodic table based upon atomic number instead of atomic mass.

15. The rest of the Nucleus
In 1932, James Chadwick, using alpha particles discovered a neutral atomic particle with a mass close to a proton.
Thus was discovered the neutron

16. Counting Atoms
Atomic number (Z) is the number of protons an atom contains
Change the number of protons, change the element
Isotopes are atoms with the same number of protons but with varying numbers of neutrons
Two was to symbolize isotopes, hyphen notation and nuclear symbol notation
Mass Number is the total number of protons and neutrons in an atom
Nuclide is a general term for a specific isotope of an element

17. Masses of Atoms
Relative atomic mass uses units called “atomic mass units (amu)”
An amu is defined as the 1/12th the mass of a carbon-12 atom
Since a Carbon-12 atom contains 6 protons and 6 neutrons, 1 amu equals the approximate mass of either a proton or neutron
Atomic Weight is the weighted average of atomic masses of all known isotopes
The most abundant isotopes contribute more to the overall mass of the element