1. 4 Periodic Trends: 1) Atomic Radius 2) Ionic Radius 3) Ionization Energy 4) ElectroNegativity ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

2. Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties. local.ans.org/mi/Teacher_CD/.../Hist_PeriodicTable.ppt

3. Trends in Atomic Size First problem: Where do you start measuring from? The electron cloud doesn’t have a definite edge. They get around this by measuring more than 1 atom at a time. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

4. Atomic Size Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

5. Trends in Atomic Size Influenced by three factors: 1. Energy Level  Higher energy level is further away. 2. Charge on nucleus  More charge pulls electrons in closer. 3. Shielding effect e e repulsion ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

6. Group trends As we go down a group... each atom has another energy level, so the atoms get bigger. H Li Na K Rb ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

7. Periodic Trends As you go across a period, the radius gets smaller. Electrons are in same energy level. More nuclear charge. Outermost electrons are closer. NaMgAlSiPSClAr ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

8. Atomic Radii Figure 8.9 xbeams.chem.yale.edu/~batista/113/chapter8/ch8.ppt

9. Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by the outermost electrons. Z* is the nuclear charge experienced by the outermost electrons. Z* increases across a period owing to incomplete shielding by inner electrons. Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* by --> [ Z - (no. inner electrons) ] Estimate Z* by --> [ Z - (no. inner electrons) ] Z=Atomic Number Z=Atomic Number Charge felt by e- in Li:Z* = Z-#Inner e-=3 - 2 = 1 Charge felt by e- in Li:Z* = Z-#Inner e-=3 - 2 = 1 Be Z* = 4 - 2 = 2 Be Z* = 4 - 2 = 2 B Z* = 5 - 2 = 3and so on! B Z* = 5 - 2 = 3and so on! xbeams.chem.yale.edu/~batista/113/chapter8/ch8.ppt

10. Shielding The electron in the outermost energy level experiences more inter-electron repulsion (shielding). Second electron has same shielding, if it is in the same period ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

11. Ion Sizes CATIONS are SMALLER than the atoms from which they come. CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES. The electron/proton attraction has gone UP and so size DECREASES. Li,152 pm 3e and 3p Li +, 78 pm 2e and 3 p + Forming a cation. xbeams.chem.yale.edu/~batista/113/chapter8/ch8.ppt

12. Ion Sizes ANIONS are LARGER than the atoms from which they come. ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES. The electron/proton attraction has gone DOWN and so size INCREASES. Trends in ion sizes are the same as atom sizes. Trends in ion sizes are the same as atom sizes. Forming an anion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p - xbeams.chem.yale.edu/~batista/113/chapter8/ch8.ppt

13. Trends in Ionic Size Cations form by losing electrons. Cations are smaller that the atom they come from. Metals form cations. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

14. Ionic size Anions form by gaining electrons. Anions are bigger that the atom they come from. Nonmetals form anions. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

15. Group trends Adding energy level Ions get bigger as you go down. Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+ ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

16. Periodic Trends Across the period, nuclear charge increases so they get smaller. Energy level changes between anions and cations. Li 1+ Be 2+ B 3+ C 4+ N 3- O 2- F 1- ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

17. Trends in Ion Sizes Figure 8.13 xbeams.chem.yale.edu/~batista/113/chapter8/ch8.ppt

18. Trends in Ionization Energy The amount of energy required to completely remove an electron from a gaseous atom. Removing an electron makes a +1 ion. The energy required to remove the first electron is called the first ionization energy. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

19. Ionization Energy The second ionization energy is the energy required to remove the second electron(s). Always greater than first IE. The third IE is the energy required to remove a third electron, which is greater than 1st or 2nd IE. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

20. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276 ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

21. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276 ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

22. Ionization Energy See Screen 8.12 IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg + (g) + e- xbeams.chem.yale.edu/~batista/113/chapter8/ch8.ppt

23. Mg (g) + 735 kJ ---> Mg + (g) + e- Mg + (g) + 1451 kJ ---> Mg 2+ (g) + e- Mg 2+ (g) + 7733 kJ ---> Mg 3+ (g) + e- Energy cost is very high to dip into a lower shell. Ionization Energy See Screen 8.12 xbeams.chem.yale.edu/~batista/113/chapter8/ch8.ppt

24. What determines IE The greater the nuclear charge, the greater IE. Greater distance from nucleus decreases IE Shielding effect ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

25. Group trends As you go down a group, first IE decreases because... The electron is further away. More shielding. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

26. Periodic trends All the atoms in the same period have the same energy level. Same shielding. But, increasing nuclear charge So IE generally increases from left to right. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

27. Electronegativity The tendency for an atom to attract electrons to itself when it is chemically combined with another element. High electronegativity means it pulls the electron toward it. Atoms with large negative electron affinity have larger electronegativity. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

28. Group Trend The further down a group, the farther the electron is away, and the more electrons an atom has. More willing to share. Low electronegativity. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

29. Periodic Trend Metals are at the left of the table. They let their electrons go easily Low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away from others High electronegativity. ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

30. Ionization energy and ElectroNegativity INCREASE ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt

31. Atomic size INCREASES, shielding constant Ionic size increases ibchem.com/IB/ibfiles/periodicity/per_ppt/pt_trends.ppt