1. Atomic Structure www.lab-initio.com

2. Standards Standards

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  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  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  The nucleus is small  The nucleus is dense  The nucleus is positively charged  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:

9. Quantum Mechanics

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

11. 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… Quantum Mechanical Model of the Atom

12. You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! “One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg Heisenberg Uncertainty Principle

13. Principle Quantum number 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: 2n 2 Electron Energy Level (Shell)

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

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

16. 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. p Orbital Shapes

17. 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”! d Orbital Shapes

18. f Orbital Shapes

19. Energy Level (n) Sublevels in main energy level (n sublevels) Number of orbitals per sublevel Number of Electrons per sublevel Number of Electrons per main energy level (2n 2 ) 1s122 2spsp 1313 2626 8 3spdspd 135135 2 6 10 18 4spdfspdf 13571357 2 6 10 14 32 Energy Levels, Sublevels, Electrons

20. Orbital Filling Table

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

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

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

24. ElementConfiguration notation Orbital notationNoble gas notation Lithium1s 2 2s 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 1 Beryllium1s 2 2s 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 Boron1s 2 2s 2 p 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 1 Carbon1s 2 2s 2 p 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 2 Nitrogen1s 2 2s 2 p 3 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 3 Oxygen1s 2 2s 2 p 4 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 4 Fluorine1s 2 2s 2 p 5 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 5 Neon1s 2 2s 2 p 6 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 6