Periodic Trends

1. Periodic Trends Why do atoms with all the same parts (protons, neutrons, and electrons) behave so differently?

2. Atomic Size

3. Atomic Size • The problem: Where do you start measuring? • The electron cloud doesn’t have a definite edge. • We get around this by measuring more than 1 atom at a time.

4. Atomic Size } • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Radius

5. Trends in Atomic Size • Influenced by two factors. • Energy Level: Higher energy levels have more orbitals, therefore larger • The charge on the nucleus (the atomic number): More charge pulls electrons in closer.

6. Group trends H • As we go down a group • Each atom has another energy level, • So the atoms get bigger. Li Na K Rb

7. Periodic Trends • As you go across a period the radius gets smaller. • Same energy level. • More nuclear charge. • Electrons are pulled closer. Na Mg Al Si P S Cl Ar

8. Rb Overall K Na Li Atomic Radius (nm) Kr Ar Ne H 10 Atomic Number

9. Ionic Size

10. Ionic Size - Cations • Cations form by losing electrons (are positive ions). • Cations are smaller than the atom they come from. • Metals form cations. • Cations of representative elements have noble gas configuration.

11. Ionic size - Anions • Anions form by gaining electrons (are negative ions). • Anions are bigger than the atom they come from. • Nonmetals form anions. • Anions of representative elements have noble gas configuration.

12. Group trends for Ions • Adding energy level • Ions get bigger as you go down. Li+1 Na+1 K+1 Rb+1 Cs+1

13. Periodic Trends for Ions • Across the period nuclear charge increases so they get smaller. • Reminder: Cations are smaller than the atom they came from and anions are larger. N-3 O-2 F-1 B+3 Li+1 C+4 Be+2

14. Ionization Energy

15. Ionization Energy • The amount of energy required to completely remove an electron from a gaseous atom. • Removing one electron makes a +1 ion. • The energy required is called the first ionization energy.

16. Ionization Energy • The second ionization energy is the energy required to remove the second electron. • Always greater than first IE. • The third IE is the energy required to remove a third electron. • Greater than 1st of 2nd IE.

17. Symbol First Second Third 11815 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne

18. What determines IE • The greater the nuclear charge (# of protons) the greater IE. • Less distance from nucleus increases IE • Shielding decreases IE.

19. Shielding • The electron on the outside energy level has to “look through” all the other energy levels to “see” the nucleus • There is less “pull” on it. • Lower IE

20. Group trends • As you go down a group first IE decreases because • The electron is further away. • More shielding, therefore, less pull on the electrons

21. Periodic trends • All the atoms in the same period have the same energy level. • All have the same shielding. • Increasing nuclear charge from left to right • So IE generally increases from left to right.

22. He • He has a greater IE than H. • same shielding • greater nuclear charge H First Ionization energy Atomic number

23. He • Li has lower IE than H • more shielding • further away • outweighs greater nuclear charge H First Ionization energy Li Atomic number

24. He • Be has higher IE than Li • same shielding • greater nuclear charge H First Ionization energy Be Li Atomic number

25. He • B has lower IE than Be • same shielding • greater nuclear charge • By removing an electron we make s orbital half filled H First Ionization energy Be B Li Atomic number

26. He C H First Ionization energy Be B Li Atomic number

27. He N C H First Ionization energy Be B Li Atomic number

28. He N • Breaks the pattern because removing an electron gets to 1/2 filled p orbital O C H First Ionization energy Be B Li Atomic number

29. He F N O C H First Ionization energy Be B Li Atomic number

30. Ne He F N • Ne has a lower IE than He • Both are full, • Ne has more shielding • Greater distance O C H First Ionization energy Be B Li Atomic number

31. Ne He • Na has a lower IE than Li • Both are s1 • Na has more shielding • Greater distance F N O C H First Ionization energy Be B Li Na Atomic number

32. First Ionization energy Atomic number

33. Driving Force • Full Energy Levels are very low energy. • Noble Gases have full orbitals. • Atoms behave in ways to achieve noble gas configuration.

34. Electronegativity

35. Electronegativity • The tendency for an atom to attract electrons to itself when it is chemically combined with another element. • How fair it shares. • Big electronegativity means it pulls the electron toward it.

36. Group Trend • The further down a group the further the valence electrons are away from the nucleus and the more electrons an atom has. • Less tightly held. • Low electronegativity.

37. Periodic Trend • Metals are at the left end. • They let their electrons go easily • Low electronegativity • At the right end are the nonmetals. • They want more electrons. • Try to take them away. • High electronegativity.

38. Ionization energy and electronegativity INCREASE

39. Atomic size increases, shielding constant Ionic size increases