1. IIIIII Unit 5 AP Chemistry Periodic Table Trends

2. Periodic Law zWhen elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

3. ALL Periodic Table Trends zInfluenced by three factors: 1. Energy Level yHigher energy levels are further away from the nucleus. 2. Charge on nucleus (# protons) yMore charge pulls electrons in closer. y (+ and – attract each other) 3. Shielding effect y(blocking effect)

4. What do they influence?  Energy levels and Shielding have an effect on the GROUP  Nuclear charge has an effect on a PERIOD

5. zAtomic Radius ysize of atom © 1998 LOGAL Atomic Radius zAtomic Radius yAverage distance in an atom between the nucleus and the outermost electron

6. #1. Atomic Size - Group trends zIncreases going down a group)... zeach atom has another energy level zso the atoms get bigger. yHyH yLi yNa yKyK yRb

7. #1. Atomic Size - Period Trends zGoing from left to right across a period, the size gets smaller. zElectrons are in the same energy level. zBut, there is more nuclear charge. zOutermost electrons are pulled closer. NaMgAl Si P S Ar Cl

8. zAtomic Radius yIncreases to the LEFT and DOWN Atomic Radius

9. #2. Trends in Ionization Energy zIonization energy is the amount of energy required to completely remove an electron (from a gaseous atom). zM + energy  M +1 + e - zRemoving one electron makes a +1 ion. zThe energy required to remove only the first electron is called the first ionization energy.

10. Ionization Energy zThe second ionization energy is the energy required to remove the second electron. yAlways greater than first IE. zThe third IE is the energy required to remove a third electron. yGreater than 1st or 2nd IE.

11. What factors determine IE zThe greater the nuclear charge, the greater IE. zGreater distance from nucleus decreases IE zFilled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. zShielding effect

12. Shielding zThe electron on the outermost energy level has to look through all the other energy levels to see the nucleus. zSecond electron has same shielding, if it is in the same period

13. Ionization Energy - Group trends zAs you go down a group, the first IE decreases because... yThe electron is further away from the attraction of the nucleus, and yThere is more shielding.

14. Ionization Energy - Period trends zAll the atoms in the same period have the same energy level. zSame shielding. zBut, increasing nuclear charge zSo IE generally increases from left to right. zExceptions at full and 1/2 full orbitals.

15. zTrend is opposite of atomic radius. zWhy? yIn small atoms, e - are close to the nucleus where the attraction is stronger Ionization Energy Trends

16. zFirst Ionization Energy yIncreases UP and to the RIGHT E. Ionization Energy

17. zFirst Ionization Energy E. Ionization Energy K Na Li Ar Ne He

18. yFirst Ionization energy yAtomic number He He has a greater IE than H. Both elements have the same shielding since electrons are only in the first level But He has a greater nuclear charge H

19. yFirst Ionization energy yAtomic number H He l Li has lower IE than H l more shielding l further away l These outweigh the greater nuclear charge Li

20. yFirst Ionization energy yAtomic number H He l Be has higher IE than Li l same shielding l greater nuclear charge Li Be

21. yFirst Ionization energy yAtomic number H He l B has lower IE than Be l same shielding l greater nuclear charge l By removing an electron we make filled s -sublevel Li Be B

22. yFirst Ionization energy yAtomic number H He Li Be B C

23. yFirst Ionization energy yAtomic number H He Li Be B C N

24. yFirst Ionization energy yAtomic number H He Li Be B C N O Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital

25. yFirst Ionization energy yAtomic number H He Li Be B C N O F

26. yFirst Ionization energy yAtomic number H He Li Be B C N O F Ne Ne has a lower IE than He Both are full, Ne has more shielding Greater distance

27. yFirst Ionization energy yAtomic number H He Li Be B C N O F Ne l Na has a lower IE than Li l Both are s 1 l Na has more shielding l Greater distance Na

28. zSuccessive Ionization Energies yMg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ yLarge jump in I.E. occurs when a CORE e - is removed. Ionization Energy

29. yAl1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ zSuccessive Ionization Energies yLarge jump in I.E. occurs when a CORE e - is removed. Ionization Energy

30. #3. Trends in Electronegativity zElectronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element. zThey share the electron, but how equally do they share it? zAn element with a big electronegativity means it pulls the electron towards itself strongly!

31. Electronegativity Group Trends zElectronegativity Decreases Down a Group yWhy? yAtomic size increases and valence electrons are farther from the nucleus. yMore energy levels increases shielding. So the pull from the positive nuclear charge is less. yIn General: xNon-Metals have high Electronegativities xMetals have low Electronegativities

32. Electronegativity Period Trend zMetals are at the left of the table. zThey let their electrons go easily zThus, low electronegativity zAt the right end are the nonmetals. zThey want more electrons. zTry to take them away from others zHigh electronegativity.

33. Electronegativity zTrend is also opposite of atomic size zThe smaller the atom, the more electronegative it is because of a greater nuclear charge. zException: Noble gases are not included because they generally do not want to gain electrons. They are already stable.

34. zIonic Radius yCations (+ ions) the ionic radius is smaller than the original atom. yWhy? There is an increased attraction for the fewer electrons that remain. Ionic Radius Na  Na +

35. Ionic Radius yFor Anions (– ions) the ionic radius is larger than the original atom. xWhy? The nuclear attraction is less for an increased number of electrons. xExtra electrons repel each other and spread out – larger!) © 2002 Prentice-Hall, Inc. Cl  Cl -1

36. Ion Group trends zEach step down a group is adding an energy level zIons therefore get bigger as you go down, because of the additional energy level. Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

37. Ion Period Trends zAcross the period from left to right, the nuclear charge increases - so they get smaller. zNotice the energy level changes between anions and cations. Li 1+ Be 2+ B 3+ C 4+ N 3- O 2- F 1-

38. #5. Electron Affinity zThe energy change when an electron is added to an atom forming an anion. zDifferent than Electronegativity because EN is attraction for electrons in a bond! X(g) + e -  X - (g) + Energy zA negative energy indicates energy is released when an electron is added

39. Electron Affinity Trends zTrends are the same as for electronegativity. zExceptions occur at ½ full and full sublevels. zUnits are in kJ/mole LiBeBCNOFNe -60>0-27-1220-141-328>>0

40. Practice zWhich atom is larger H or He? zWhich atom has a greater ionization energy, Ca or Sr? zWhich atom is more electronegative, F or Cl?

41. zWhich atom has the larger radius? yBe orBa yCa orBr Ba Ca Examples

42. zWhich atom has the higher 1st I.E.? yNorBi yBa orNe N Ne Examples

43. zWhich particle has the larger radius? ySorS 2- yAlorAl 3+ S 2- Al Examples