1. AP Chemistry Periodic Table, Electron Configurations and Ions

2. Modern Periodic Table Elements are arranged by increasing atomic number. Elements are arranged by increasing atomic number. –Recall that atomic number gives the number of protons in the nucleus of an atom. –It also represents the number of electrons in the neutral atom

3. Electron Configurations and the Periodic Table

4. Orbitals fill in a specific order Lowest energy to higher energy. Lowest energy to higher energy. Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. However, half filled orbitals have a lower energy, and are next best However, half filled orbitals have a lower energy, and are next best Makes them more stable.Makes them more stable. Changes the filling orderChanges the filling order

5. Write the electron configurations for these elements: Titanium - 22 electrons 1111s22s22p63s23p64s23d2 Vanadium - 23 electrons 1111s22s22p63s23p64s23d3 Chromium - 24 electrons 1111s22s22p63s23p64s23d4 (expected) BBBBut this is not what happens!!

6. Chromium is actually: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Why? Why? This gives us two half filled orbitals (the others are all still full) This gives us two half filled orbitals (the others are all still full) Half full is slightly lower in energy. Half full is slightly lower in energy. The same principal applies to copper. The same principal applies to copper.

7. Copper’s electron configuration Copper has 29 electrons so we expect: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 Copper has 29 electrons so we expect: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 But the actual configuration is: But the actual configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 This change gives one more filled orbital and one that is half filled. This change gives one more filled orbital and one that is half filled. Remember these exceptions: d 4, d 9 Remember these exceptions: d 4, d 9

8. Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel

9. Octet Rule Most elements want to have 8 valence electrons (exception He) Most elements want to have 8 valence electrons (exception He) Elements will gain or lose electrons in their valence shell in order to achieve an octet. Elements will gain or lose electrons in their valence shell in order to achieve an octet. –This process forms Ions. Ions have the same electron configuration as their nearest noble gas. Ions have the same electron configuration as their nearest noble gas.

10. Electron Configurations - Ions If you know how many electrons an element or ion has, you can write its electron configuration or draw its orbital filling diagram. If you know how many electrons an element or ion has, you can write its electron configuration or draw its orbital filling diagram. Even for an ion! Even for an ion! Example - Na has 11 electrons, Na +1 (the ion) has 10 electrons. Example - Na has 11 electrons, Na +1 (the ion) has 10 electrons. Ions will form so that it ends with a completely filled main energy level. Ions will form so that it ends with a completely filled main energy level.

11. Writing Electron Configuration of Sodium Ion 1s 2s 2p 3s 3p 4s Energy 1s 2 2s 2 2p 6 3s 1 Na + Means it lost 1 electron Na+ electron configuration 1s 2 2s 2 2p 6

12. Positive Ions Elements with less than 4 valence electrons lose their outermost “s” and “p” sublevel electrons Elements with less than 4 valence electrons lose their outermost “s” and “p” sublevel electrons Example Example Be: 1s 2 2s 2 has 2 valence electrons Be 2+ : 1s 2 notice the ion has the same configuration as He –They always form positive ions.

13. Negative Ions Elements with more than 4 valence electrons gain electrons in their outermost “p” sublevel Elements with more than 4 valence electrons gain electrons in their outermost “p” sublevel Example Example O: 1s 2 2s 2 2p 4 has 6 valence electrons O 2- : 1s 2 2s 2 2p 6 now has 8 valence electrons just like Ne –They always form negative ions.

14. Ion Formation Summary Atoms gain or lose electrons to become more stable. Atoms gain or lose electrons to become more stable. –They achieve the same electron configuration as the Noble Gases.

15. Periodic Table Organization Periods Periods Horizontal rows on the periodic table are called periods. Horizontal rows on the periodic table are called periods. Elements in the same period have the same number of valence shells. Elements in the same period have the same number of valence shells. Elements in the same period share no similar chemical properties. Elements in the same period share no similar chemical properties.

16. Groups / Families Vertical columns on the periodic Table are called groups. Vertical columns on the periodic Table are called groups. Groups are numbered 1-18 Groups are numbered 1-18 Elements in the same group have the same chemical properties because they have the same number of valence electrons. Elements in the same group have the same chemical properties because they have the same number of valence electrons. Some groups have special names. Some groups have special names.

17. Group 1 Metals Known as Alkali Metals Known as Alkali Metals Most active metals on periodic table Most active metals on periodic table React violently with water, and become more active down the group React violently with water, and become more active down the group Only have 1 valence electron (because they all end in s 1 ) Only have 1 valence electron (because they all end in s 1 ) They all form ions with a +1 charge. They all form ions with a +1 charge.

18. Group 2 Metals Known as Alkaline Earth Metals Known as Alkaline Earth Metals Also extremely reactive, but not as much as Alkali Metals Also extremely reactive, but not as much as Alkali Metals Contain 2 valence electrons (they end in s 2 ) Contain 2 valence electrons (they end in s 2 ) They all form ions with a +2 charge. They all form ions with a +2 charge.

19. Groups 3-12 These groups are known as Transition Metals These groups are known as Transition Metals Not very reactive metals Not very reactive metals Generally have 2 valence electrons Generally have 2 valence electrons Some of their “d” orbital electrons jump around into “p” orbitals so many of these elements have more than one possible positive charge. Some of their “d” orbital electrons jump around into “p” orbitals so many of these elements have more than one possible positive charge. Form colorful ion solutions (bright vibrant colors) Form colorful ion solutions (bright vibrant colors)

20. Group 17 Known as Halogens Known as Halogens Extremely reactive nonmetals; reactivity increases up a group (Fluorine the most active) Extremely reactive nonmetals; reactivity increases up a group (Fluorine the most active) Have 7 valence electrons (e - configuration ends in s 2 p 5 ) Have 7 valence electrons (e - configuration ends in s 2 p 5 ) Most of these elements are toxic in their natural state Most of these elements are toxic in their natural state Most are also diatomic in their natural state (F 2, Br 2, Cl 2, I 2 ) Most are also diatomic in their natural state (F 2, Br 2, Cl 2, I 2 ) They all form ions with a (-1) charge. They all form ions with a (-1) charge.

21. Group 18 Known as Noble gases Known as Noble gases Unreactive nonmetals Unreactive nonmetals All have 8 valence electrons All have 8 valence electrons (e- configuration ends in s 2 p 6 ), except Helium (which has 2 electrons) (e- configuration ends in s 2 p 6 ), except Helium (which has 2 electrons) They do not form ions. They do not form ions.