1. REVIEW DAY

2. Orbitals are described with 2 different numbers called Quantum Numbers 1)Principal Quantum Number (n) Integer values that describe the energy of the orbital Larger n means larger average distance of an electron from nucleus (thus larger orbital) 2)Angular Momentum Quantum Number (l) Tells us the “shape” of the orbitals

3. Orbital Diagrams

4. Electron Configuration 1s11s1 row # shell # possibilities are 1-7 7 rows subshell possibilities are s, p, d, or f 4 subshells group # # valence e- possibilities are: s: 1 or 2 p: 1-6 d: 1-10 f: 1-14 Total e- should equal Atomic # What element has an electron configuration of 1s 1 ?

5. The Order of Orbitals ■1s, ■2s, 2p, ■3s, 3p, ■4s, 3d,4p, ■5s, 4d, 5p, ■6s, 4f, 5d, 6p, ■7s, 5f, 6d, 7p, ■(8s, 5g, 6f, 7d, 8p, and 9s)

6. Orbital filling table

7. RULES!! Memorize These!! ■Three rules (To MEMORIZE) –Aufbau: Electrons fill orbitals starting in the lowest energy levels and moving out. –Pauli: No two electrons can have the same spin and occupy the same orbital. –Hunds: Electrons fill each orbital singly before any orbital get a second electron.

8. Shorthand Notation ■Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. ■Step 2: Find where to resume by finding the next energy level. ■Step 3: Resume the configuration until it’s finished.

9. Shorthand Notation ■Chlorine –Longhand is 1s 2 2s 2 2p 6 3s 2 3p 5 You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s 2 2s 2 2p 6 The next energy level after Neon is 3 So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17 [Ne] 3s 2 3p 5

10. Speed = wavelength x frequency c = (length/second) = (length/cycle) x (cycle/second) Hence, = c / and = c /

11. Hydrogen spectrum ■Emission spectrum because these are the colors it gives off or emits ■Called a line spectrum. ■There are just a few discrete lines showing 410 nm 434 nm 486 nm 656 nm Spectrum

12. Energy Example Problems ■Calculate the frequency of radiation with a wavelength of 442 nm. ■Calculate the frequency of radiation with a wavelength of 4.92 cm. 1) Convert nm to m: 442 nm x (1 m / 10 9 nm) = 4.42 x 10¯ 7 m 2) Substitute into λν = c: (4.42 x 10¯ 7 m) (x) = 3.00 x 10 8 m s¯ 1 x = 6.79 x 10 14 s¯ 1 Since the wavelengths are already in cm, we can use c = 3.00 x 10 10 cm s¯ 1 and not have to do any conversions at all. (4.92 cm) (x) = 3.00 x 10 10 cm s¯ 1 x = 6.10 x 10 9 s¯ 1

13. Example What is the frequency of radiation with a wavelength of 5.00 x 10¯ 8 m? In what region of the electromagnetic spectrum is this radiation? Solution: 1) Use λν = c to determine the frequency: (5.00 x 10¯ 8 m) (x) = 3.00 x 10 8 m/s x = 6.00 x 10 15 s¯ 1 2) Determine the electromagnetic spectrum region: Consult a convenient reference source. This frequency is right in the middle of the ultraviolet region of the spectrum.

14. Spectrum of Electromagnetic Radiation RegionWavelength (Angstroms) Wavelength (centimeters) Frequency (Hz) Energy (eV) Radio> 10 9 > 10< 3 x 10 9 < 10 -5 Microwave10 9 - 10 6 10 - 0.013 x 10 9 - 3 x 10 12 10 -5 - 0.01 Infrared10 6 - 70000.01 - 7 x 10 -5 3 x 10 12 - 4.3 x 10 14 0.01 - 2 Visible7000 - 40007 x 10 -5 - 4 x 10 -5 4.3 x 10 14 - 7.5 x 10 14 2 - 3 Ultraviolet4000 - 104 x 10 -5 - 10 -7 7.5 x 10 14 - 3 x 10 17 3 - 10 3 X-Rays10 - 0.110 -7 - 10 -9 3 x 10 17 - 3 x 10 19 10 3 - 10 5 Gamma Rays< 0.1< 10 -9 > 3 x 10 19 > 10 5

15. The Bohr Ring Atom n = 3 n = 4 n = 2 n = 1

16. Ionization Energy ■Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. –First ionization energy is that energy required to remove first electron. –Second ionization energy is that energy required to remove second electron, etc.

17. Ionization Energy ■It requires more energy to remove each successive electron. ■When all valence electrons have been removed, the ionization energy shows a HUGE increase.

18. Electronegatvity ■Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. ■The Pauling scale is the most commonly used. –Fluorine (the most electronegative element) is assigned a value of 4.0 –Values range down to cesium and francium which are the least electronegative at 0.7.

19. Electron Affinity ■Electron affinity is defined as the change in energy of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion. ■The neutral atom's likelihood of gaining an electron