1. Quantum Model and Electron Configurations

2. Atomic Models:  Old version = Bohr’s  Also known as the planetary atomic model  Describes electron paths as perfect orbits with definite diameters  Good for a visual  New version = Quantum Theory  Most accepted  Diagrams electrons of a atom based on probability of location at any one time

3. Bohr’s model:  Nucleus is in the center of an atom(like the sun) and the electrons orbit the nucleus similar to the planets.  Orbits are called shells  1 st shell = 2 electrons  2 nd shell = 8 electrons  3 rd shell = 18 electrons  4 th shell = 32 electrons Last slide

4. QOD VOCAB  What is the approximate mass of an electron?  WHICH IS A TRUE STATEMENT?  Compounds can be broken down (decomposed) by chemical means  Compounds can be decomposed by physical means 0.000549 amu

5. Quantum THEORY & Mechanics Study of how light interacts with matter

6. Quantum Theory:  To better the description of the atomic structure, atoms were exposed to energy (heat) which made the electrons go into what is called the excited state (normal = ground state).  When electrons returned to ground state they emitted energy in the form of light.

7. Quantum Theory:  This method of study is called spectroscopy (spectrum)  Visible light = part of the electromagnetic spectrum between 400-700 nm

8. Electromagnetic Spectrum

9. Quantum Theory:  Electromagnetic Spectrum  From crest to crest = frequency which is measured in hertz.  This therefore can be used to identify elements (absorption of energy and color emitted is a fingerprint of an element) Kind of like wearing your team colors. Team Oxygen Team Carbon

10. Continuous spectrum of white light  When you pass sunlight through a prism, you get a continuous spectrum of colors like a rainbow. Line-Emission Spectrum  However, when light from Hydrogen & Helium gases were passed through a prism, they found a dark background with discrete lines.  WHY? This lead to the quantum theory. H He

11. A scientist, Bohr suggested that electrons must exist in Electron Orbitals (shows the most probable area to find an electron of a certain energy.) So whenever an excited hydrogen atom falls to its ground state or lower energy level, it emits a photon of light, which means that energy levels must be fixed. Quantum Theory:

12.  Video Video

13. Quantum Theory: Electron Configuration  Electrons (e-) of atoms are the basis for every chemical reaction.  In quantum theory, electrons exist in orbitals based on probabilities and these orbitals are arranged within energy levels. Notice… these orbitals look different from Bohr’s. This diagram is more correct.

14. Quantum Theory: Electron Configuration Quantum Numbers  Quantum numbers specify the properties of atomic orbitals and the properties of electrons in those orbitals  We will define these numbers & letters. Example of Quantum #: 3s 2

15. Quantum Theory: Electron Configuration Principle Quantum Number (n)  Is equal to the number of the energy level ( n ).  The principle quantum # corresponds to the energy levels 1-7 which is the period number (row) on the periodic table. Example of Quantum #: 3s 2

16. P E R I  O D S 1-7 Blocks and Sublevels d (n-1) 12345671234567 4545

17. Maximum e- in Energy Levels  The maximum number of e- in any one level is given by the equation 2 n 2  Calculate the maximum number of electrons that can occupy the 4 th principal quantum number (period 4).  Solve: Use 2n 2 2(4) 2 32 electrons total Quantum Theory: Electron Configuration N=1, 2e N=2, 8e N=3, 18e N=4, 32e

18. Sublevels and Orbitals  An energy level in made up of many energy states called sublevels.  The number of sublevels for each energy level is equal to the value of the principal quantum number. EX: one sublevel in energy level one (period 1) two sublevels in level two (period 2) three sublevels in level three (period 3) *now lets find out what those sublevels are called… Quantum Theory: Electron Configuration Example of Quantum #: 3s 2

19. Sublevels and Orbitals  There are 4 sublevels: spdf  Energy levels and sublevels work together to form an e- cloud.  e- are repelled by one another and move as far apart as possible.  e- clouds take on characteristic shapes called orbitals. Quantum Theory: Electron Configuration

20. Sublevels and Orbitals (notice the shapes)

21. Orbital Shapes s orbitals are spherical. This diagram represents an s orbital. p orbitals are “dumbbell” shaped. This diagram represents 1 of the 3 types of p orbitals. d orbitals contains 5 possible orbital shapes. f orbitals contain 7 possible orbital shapes.

22. Electrons & Orbitals  Orbitals overlap and change shape as electrons are added.  Each orbital can only hold 2 electrons. Example of Quantum #: 3s 2

23. Electrons and Orbitals (count 2 electrons per orbit)

24. Orbitals, and Electrons per Sublevel Principal Quantum Number (n) Sublevel # of Orbitals # of Electrons per Orbital 1s12 2spsp 1313 2626 3s p d 1 3 5 2 6 10 4s p d f 1 3 5 7 2 6 10 14

25. QOD VOCAB  The principal quantum number corresponds to the:  Which statement is true: The characteristic bright- line spectrum (color) of an element is produced when its electrons…  Move to an excited higher energy state  Return to a lower ground energy state Energy Levels Periods on the periodic table

26. Distribution of Electrons  Atoms are electronically neutral. (for now)  There is an electron for every proton in the nucleus.  The larger the atom, the larger the electron cloud. 1. Pauli Exclusion Principle: only two e- can occupy the same orbital due to the opposite electronic spin.

27. Electron Filling Diagram Sublevels and orbitals are filled as indicated in the diagram. Example: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 … Notice… they don’t go in order ! # electrons in the orbital Sublevel orbital energy level

28. Label your blank periodic table. Read it “like a book”

29. WRITE the Electron Configuration Now try: 1. C 2. Kr 3. Ca 4. Fe 5. Hg

30. QOD VOCAB  What is the total number of electrons that can be held in the third principal energy level?  Quantum Theory of atomic structure states all except :  Electrons orbit the nucleus in perfect paths  Electrons form clouds based on probability of location  Electron clouds form characteristic shapes due to repelling of negative charges  Electrons occupy the lowest energy levels before moving into higher energy levels 2n 2 18

31. Label your blank periodic table. Read it “like a book”

32. DRAW the Electron Configuration  Carbon has 6 e- (same as protons)  Start with lowest energy level and place one electron in each orbital. Spins must be in same direction within orbitals of the same energy level.  If there are remaining e-, pair up singles in same energy level before moving to next highest energy level.

33. Electron Configuration 1s 2s 2p Carbon’s electron config. is: 1s 2 2s 2 2p 2 Superscripts total the number of electrons 2 + 2 + 2 = 6

34. *Notice that you can write the electron configuration based on the orbital diagram. *When asked to draw or diagram, use arrow configuration. *When asked to write, use 1s 2,2s 2 … configuration. Last slide

35. Electron Configuration – Noble Gas Configuration  Electron Configuration demonstrates a periodic trend, so you can write shorthand electron configuration using the electron configuration of the noble gases in Group 18 of the periodic table.  Noble gases have stable configurations.

36. Noble Gas Configuration  When writing shorthand e- config for an element, refer to the noble gas in the energy level (period) just above the element.  Write the symbol of the noble gas in brackets.  Write out the remaining e-config based on the energy filling diagram. Electron Configuration Na = 1s 2 2s 2 2p 6 3s 1 Al = 1s 2 2s 2 2p 6 3s 2 3p 1 Ne = 1s 2 2s 2 2p 6 Shorthand Electron Configuration Na = [Ne] 3s 1 Al = [Ne] 3s 2 3p 1

37. Noble Gas Configuration EX: Na Step 1: Na is in period 3 so refer to the noble gas in period 2 which is Neon. Step 2: Write Ne in brackets. [Ne] Step 3: Now write remaining electrons in standard form. 3s 1. Step 4: Combine. [Ne]3s 1

38. Noble Gas Configuration EX: Br Step 1: Br is in period 4 so refer to noble gas from period 3 which is Argon. Step 2: Write in brackets. [Ar] Step 3: Write remaining electrons. 4s 2 3d 10 4p 5 Step 4: Combine to form: [Ar] 4s 2 3d 10 4p 5 *Check your work: Add the number of electrons from the noble gas (18) to the subscripts of the remaining e-config (17). 18+17=35 which is the electrons for Br.

39. Nobel Gas Configuration Now try: 1. I 2. Kr 3. Na 4. Cu

40. Label your blank periodic table. Read it “like a book”

41. Electron Configuration with Ions  When we write the electron configuration of a positive ion, we remove one electron for each positive charge: Na → Na + 1s 2 2s 2 2p 6 3s 1 → 1s 2 2s 2 2p 6  When we write the electron configuration of a negative ion, we add one electron for each negative charge: O → O 2- 1s 2 2s 2 2p 4 → 1s 2 2s 2 2p 6

42. Electron Configuration with Ions Now try: 1. Ca +2 2. Fe -3

43. Label your blank periodic table. Read it “like a book”

44. QOD VOCAB  What element has completely filled 3p orbitals?  Which of the following is the correct name for Ca +1 ?  Calcium isotope  Calcium  Calcium ion  Calcium with extra electrons Argon (Ar) 1s 2 2s 2 2p 6 3s 2 3p 6

45. Label your blank periodic table. Read it “like a book” * * **

46. 1 1 2 3 4 5 6 7 32 5 4 76 1 98 1 10 1 2 12345 6 1 2 3 4 5 6 123456789 11121314 4 5 3 4 5 6 S - Block D - Block P - Block F - Block 4 5

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