1. ELECTRON CONFIGURATION

2. Electron Configuration  The way electrons are arranged around the nucleus.

3. Atomic Spectra and Bohr Bohr said classical view is wrong. Need a new theory — now called e- can only exist in certain discrete orbits e- is restricted to energy state (quanta = bundles of energy)

4. Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms. He developed the Solution gives set of math expressions called Each describes an allowed energy state of an e- E. Schrodinger 1887-1961 Quantum or Wave Mechanics

5. Heisenberg Uncertainty Principle Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the (= mv) of an electron. (= mv) of an electron. We define e- energy exactly but accept limitation that we do not know exact position. Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the (= mv) of an electron. (= mv) of an electron. We define e- energy exactly but accept limitation that we do not know exact position. W. Heisenberg 1901-1976

6. Quantum Mechanical Model  1920’s  Werner Heisenberg ( )  Louis de Broglie ( )  Erwin Schrodinger (mathematical equations using probability, numbers)

7. Principal Quantum Number, n  Indicates main n = 1, 2, 3, 4… positive integers As n increases, the electron’s from the nucleus and the electron’s energy increases.  Each main energy level has sub-levels  There are 4 sublevels

8.  The principle quantum number, n, determines the number of sublevels within the principle energy level.

9. Electron Configuration  Electrons always go into the lowest possible energy level (nearest the nucleus).  The atomic number is the number of protons, and hence the number of electrons.  The electrons are arranged in different levels, by filling up an inner level before filling an outer one. The energy sublevels are filled in a specific order as shown by the arrow diagram given below:

10. Orbital Quantum Number, ℓ (Angular Momentum Quantum Number)  This quantum number indicates the or type of orbital that corresponds to a particular sublevel.  ℓ = n-1 ℓ sublevel 0 s 1 p 2 d 3 f

11. Orbital  The space where there is a that it is occupied by a pair of electrons.  Orbitals are solutions of Schrodinger’s equations.

12. Orbitals in Sublevels Sublevel# Orbitals# electrons s1 p3 d5 f7

13. Magnetic Quantum Number, m (subset of the ℓ quantum number)  Also indicates the numbers and orientations of orbitals around the nucleus.  The value of m takes whole-number values, depending on the value of ℓ.  The number of orbitals includes 1 s orbital 3 p orbitals 5 d orbitals 7 f orbitals

14. Spin Quantum Number ( orientation of an electron’s magnetic field )  The quantum number is represented by: +1/2 or -1/2 ( or )  A single orbital can hold a maximum of electrons, which must have spins.  Remember: Opposites attract and Like

15. Arrangement of Electrons in Atoms Electrons in atoms are arranged as

16. QUANTUM NUMBERS The shape, size, and energy of each orbital is a function of 3 quantum numbers which describe the location of an electron within an atom or ion n (principal) ---> l (orbital) ---> m l (magnetic) ---> The fourth quantum number is not derived from the wave function s (spin) ---> of the electron (clockwise or counterclockwise: ½ or – ½)

17. Three rules are used to build the electron configuration:  Aufbau principle  Pauli Exclusion Principle  Hund’s Rule

18. Aufbau Principle  Electrons occupy orbitals of energy first.  So, the order in which the orbitals are filled matches the order of energies.

19. Aufbau Diagram

20. -Pauli Exclusion Principle (Wolfgang Pauli, Austria, 1900-1958) -Electron Spin Quantum Number  An orbital can hold only electrons and they must have spin.  Electron Spin Quantum Number (m s ): +1/2, -1/2

21. Pauli Exclusion Principle Two electrons can have the same value of n by being in the same main energy level. These two electrons can also have the same value of l by being in orbitals that have the same shape. These two electrons may also have the same value of m by being in the same orbital. But these two electrons have the same spin quantum number. If one electron has the value of 1/2, then the other electron must have the value of –1/2.

22. Hund’s Rule In a set of orbitals, the electrons will fill the orbitals in a way that would give the number of parallel spins (maximum number of unpaired electrons). Analogy: Students could fill each seat of a school bus, one person at a time, before doubling up.

23. Aufbau Diagram for Hydrogen

24. Aufbau Diagram for Helium

25. Aufbau Diagram for Lithium

26. Aufbau Diagram for Beryllium

27. Aufbau Diagram for Boron

28. Aufbau Diagram for Carbon

29. Aufbau Diagram for Nitrogen

30. Notations of Electron Configurations  Lets take a look at an example and write it in both standard and shorthand notation.  A Li atom has 3 electrons; the third one is added to the next lowers energy orbital, 2s and is unpaired.  Standard  We would write it as Li:  Shorthand  We would write it as Li:

31. Aufbau Diagram for Fluorine

32. Standard Notation of Fluorine Main Energy Level Numbers Sublevels Number of electrons in the sub level 1s 2 2s 2 2p 5

33. Shorthand Notation  Use the noble gas that is located in the periodic table right before the element.  Write the symbol of the noble gas in.  Write the configuration after the brackets.  Ex: Fluorine:

34. Blocks in the Periodic Table