1. Atomic Structure

2. Standards
Students know how to relate the position of an element in the periodic table to its atomic number and atomic mass.
Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass.

3. Modern Atomic Theory All matter is composed of atoms
Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!
Atoms of an element have a characteristic average mass which is unique to that element.
Atoms of any one element differ in properties from atoms of another element

4. Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

5. Conclusions from the Study of the Electron
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.
Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons
Electrons have so little mass that atoms must contain other particles that account for most of the mass

6. Thomson’s Atomic Model
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

7. Rutherford’s Gold Foil Experiment
Alpha () particles are helium nuclei
Particles were fired at a thin sheet of gold foil
Particle hits on the detecting screen (film) are recorded

8. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

9. Quantum Mechanics

10. The Bohr Model of the Atom
I pictured electrons orbiting the nucleus much like planets orbiting the sun.
But I was wrong! They’re more like bees around a hive.
Neils Bohr

11. Quantum Mechanical Model of the Atom
Mathematical laws can identify the regions outside of the nucleus where electrons are most likely to be found.
These laws are beyond the scope of this class…

12. Electron Energy Level (Shell)
Generally symbolized by n, it denotes the probable distance of the electron from the nucleus. “n” is also known as the Principle Quantum number
Number of electrons that can fit in a shell:
2n2

13. Electron Orbitals
An orbital is a region within an energy level where there is a probability of finding an electron.
Orbital shapes are defined as the surface that contains 90% of the total electron probability.

14. s Orbital Shapes The s orbital has a spherical shape centered around
the origin of the three axes in space.

15. p Orbital Shapes
There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space.

16. d Orbital Shapes
Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells”
…and a “dumbell
with a donut”!

17. f Orbital Shapes

18. Energy Levels, Sublevels, Electrons
Sublevels in
main energy
level
(n sublevels)
Number of
orbitals per
sublevel
Electrons
per sublevel
Electrons per
level (2n2) 1 s 2 p 3 6 8 d 5 10 18 4 f 7 14 32

19. Orbital Filling Table

20. Electron Spin
Electron spin describes the behavior (direction of spin) of an electron within a magnetic field.
Possibilities for electron spin:

21. Pauli Exclusion Principle
Two electrons occupying the same orbital must have opposite spins
Wolfgang
Pauli

22. Electron Configurations of the elements of the first three series

23. Element Configuration notation Orbital notation Noble gas Lithium
1s22s1
____ ____ ____ ____ ____
1s s p
[He]2s1
Beryllium
1s22s2
[He]2s2
Boron
1s22s2p1
[He]2s2p1
Carbon
1s22s2p2
[He]2s2p2
Nitrogen
1s22s2p3
1s s p
[He]2s2p3
Oxygen
1s22s2p4
[He]2s2p4
Fluorine
1s22s2p5
[He]2s2p5
Neon
1s22s2p6
[He]2s2p6