1. The Atom What to look for: Three Essential Questions
E. Q. #1 = What are the parts of an atom?
E. Q. #2 = In what manner can subatomic particles be differentiated?
E. Q. #3 = How does the atomic structure impact chemical change & reactivity?

2. The Building Block of Matter
400 BC Democritus (Greek Philosopher)
Atomos = “indivisible”
…like a brick is used to build a building; an atom is used to build all matter…

3. First Atomic Theory 1808 John Dalton (English Chemist)
Using experiments formulates the first atomic theory
All elements are composed of indivisible particles called atoms
Atoms of the same elements are exactly the same
Atoms of different elements are different
Two or more elements combine to form compounds

4. Subatomic Particles
The idea that the atoms was not a solid sphere, but a composition of small particles which came to be known as subatomic particles
Julius Plucker, William Crooks & J J Thomson
mid-1800’s

5. J J Thomson
1897 Discovers the electron

6. Ernest Rutherford (English Physicist)
Gold Foil Experiments
1908 Discovers that the atom is mostly empty space
Discovers a hard, dense center of the atom called the NUCLEUS

7. The Gold Foil Experiment

8. The atom is mostly empty space
Most of the particles passed straight through to gold foil.
First important discovery by Rutherford

9. The NUCLEUS is Discovered
1911 Rutherford proves that the particles that do not pass through are hitting a small, dense center of the atom called the nucleus.
Rutherford later goes on to discover the proton

10. 3
Name three major discoveries made by Ernest Rutherford & the Gold Foil Experiment:
1. _______________________________________
2. _______________________________________
3. _______________________________________
What are the two results of firing particles at atoms in the foil of the Gold Foil Experiments:
1. _______________________________________
2. _______________________________________ 2 1
Who was the first to attempt to form a complete atomic theory?
1. _______________________________________

11. Electron Energy Levels
1913 Niels Bohr (Danish physicist) discovers the electrons travel around the outside of the atom on different energy levels
The Bohr Model = “solar system” model

12. The Neutron is Found
1932 James Chadwick proves the existence of the neutron.
Only 4 years later the neutron is used to split the first atom
Why was the neutron’s existence more difficult to prove?
What evidence hinted of its existence? “the missing mass”
E. Q. #1 = What are the parts of an atom?

13. The Subatomic Particles
Location
Mass (AMU)
Charge
Electron
Electron Cloud
negative
Proton
Nucleus 1
positive
Neutron
neutral
E. Q. #2 = In what manner can subatomic particles be differentiated?

14. Finding the Number of Subatomic Particles in an Atom
Atomic Number = the number of protons in the nucleus of an atom
Identifies each element
Hydrogen is 1 and it has 1 proton in its nucleus (only hydrogen has 1 proton)
Helium is 2 and it has 2 protons in its nucleus (only helium has 2 protons)

15. Electrons
IF the atom is neutral (same number positive charges as negative charges) then the number of electrons will be the same as protons.
ION = a charged atom (unequal number of protons and electrons)
+1 charge = the atom LOST one electron
-1 charge = the atom GAINED one electron
+2 charge = LOST 2 electrons
-2 charge = GAINED 2 electrons and so on

16. Example of Ions Sodium will lose one electron when it bonds
Na+1 or Na+
Sodium’s atomic number is 11, so it has 11 protons.
The +1 charge means Na+1 has 10 electrons

17. Another Example Chlorine will gain an electron when it bonds
Cl-1 or Cl-
Chlorine’s atomic number is 17, so it has 17 protons
The -1 charge means Cl-1 has 18 electrons

18. Mass # - Atomic # = # of neutrons
Mass Number
The two subatomic particles that have mass?
Proton and neutron, both found in the nucleus
Mass Number = # of Protons + # of Neutrons
Mass # - Atomic # = # of neutrons

19. Practice Problems Mg 12 24 K 19 39 20 Na+1 11 23 +1 10 F-1 9 -1 Mg+2
Symbol
Atomic
Number
Mass
Charge
Protons
Neutrons
Electrons Mg 12 24 K 19 39 20
Na+1 11 23 +1 10
F-1 9 -1
Mg+2 +2
Do Parts of the Atom & Atomic Structure practice sheet now.

20. ISOTOPES
Atoms of the same element with a different number of neutrons.
Ex: carbon-12 & carbon-14
Carbon-12
6 protons
6 neutrons
6 electrons
Carbon-14
8 neutrons

21. Isotopes of Hydrogen Hydrogen-1 (protium) 0 neutrons
Hydrogen-2 (deuterium) 1 neutron
Hydrogen-3 (tritium) 2 neutrons

22. The ELECTRON CLOUD
The electron cloud is separated by different energy levels.
Electrons with less energy travel on the levels closest to the nucleus.
Energy levels are divided into sub-levels
Sub-levels are divided into orbitals
All of these levels, sub-levels & orbitals can overlap.

23. VALENCE ELECTRONS Valence electrons = the outermost electrons
The electrons involved in bonding
Determines the chemical properties of an atom
E. Q. #3 = How does the atomic structure impact chemical change & reactivity?
Do Practice sheets What Is This Atom Now? 1 & 2 now.

24. Maximum # of electrons on each energy level
More about Electrons
Maximum # of electrons on each energy level
Level 1 = 2 electrons
Level 2 = 8 electrons
Level 3 = 18 electrons
Level 4 = 32 electrons

25. “the Stable 8” gain electrons lose electrons share electrons
Due to the overlapping levels, sub-levels and orbitals the maximum number of outer most electrons is EIGHT.
An atom is becomes stable when it outer-most energy levels is full of electrons
This explains the chemical bonding properties of the atoms
Atoms will do one of three things in order to fill their outer most energy levels.
gain electrons
lose electrons
share electrons

26. Concepts Coming Up Next
The number of valence electrons an atom has determines the group it becomes to on the PERIODIC TABLE.
Groups or families of elements on the periodic table has similar chemical properties.

27. Concepts Coming Up Next
Valence electrons determine how atoms will bond with other atoms in order to complete their outer-most energy level.
The Stable 8
CHEMICAL BONDING
Ionic bonding (gaining/losing of electrons)
Covalent bonding (sharing of electrons)