2. Why do the atoms of elements get smaller when moving from left to right within a row (period) across the periodic table?

3. Why do Atoms get smaller?

4. To answer this question let’s look at the 2nd row in the periodic table

5. Charged particles inside the atom determine its size. A neutral lithium (Li) atom has 3 protons in its nucleus and 3 electrons in 2 energy levels in the space surrounding the nucleus. protons electrons

6. The diameter of an atom is determined by the strength with which the protons in the nucleus pull on the outermost shell of electrons. This electron “sees” 2 electrons and 3 protons when “looking” inward. An effective nuclear charge of 1+. The larger the effective nuclear charge, the stronger the pull, the shorter the distance across the atom.

7. Now let’s compare a Li atom to a Be atom. Li Be What does an electron in the outer energy level of a Be atom “see” when it “looks” inward? 4 protons and 2 electrons. Electrons don’t see electrons in the same energy level. An effective nuclear charge of 2+.

8. Be This larger effective nuclear charge pulls the outermost level (valence shell) of electrons closer to the nucleus.

9. Be This makes the distance from one side of the atom to the other smaller.

10. Now let’s compare a Be atom to a B atom. Be What does an electron in the outer energy level of a B atom “see” when it “looks” inward? 5 protons and 2 electrons. An effective nuclear charge of 3+. B An effective nuclear charge of 2+.

11. This larger effective nuclear charge pulls the outermost level (valence shell) of electrons closer to the nucleus. B

12. B This makes the distance from one side of the atom to the other smaller.

13. Now let’s recap what happens to the atomic diameter as one moves across the periodic table from left to right. A Li atom has an effective nuclear charge of 1+ A Be atom has an effective nuclear charge of 2+ A B atom has an effective nuclear charge of 3+ The larger the effective nuclear charge the _____________(bigger, smaller) the atom. Smaller. Let’s look at diagrams to confirm

14. Li 3 p 1+, 2e 1- ENC - 1+ Be 4 p 1+, 2e 1- ENC - 2+ B 5 p 1+, 2e 1- ENC - 3+

15. Smaller, Why? Larger effective nuclear charge Now let’s model the atoms of row 3 elements

16. 11 Na row 3 3 energy levels group 1 1 e 1- in outer (valence) shell 11 e 1- total, 11 p 1+ 2e 1- 8e 1- 1e 1- Outermost electrons “see” 11p 1+, 10e 1- Effective Nuclear Charge is 1+

17. 12 Mg row 3 3 energy levels group 2 2 e 1- in outer (valence) shell 12 e 1- total, 12 p 1+ 2e 1- 8e 1- 2e 1- Outermost electrons “see” 12p 1+, 10e 1- Effective Nuclear Charge is 2+

18. 12p 1+ 2e 1- 8e 1- 2e 1- Larger effective nuclear charge makes a Mg atom smaller than a Na atom

19. 13 Al row 3 3 energy levels group 3 3 e 1- in outer (valence) shell 13 e 1- total, 13 p 1+ 2e 1- 8e 1- 3e 1- Outermost electrons “see” 13p 1+, 10e 1- Effective Nuclear Charge is 3+

20. 13p 1+ 2e 1- 8e 1- 3e 1- Larger effective nuclear charge makes an Al atom smaller than a Mg atom

21. 14 Si row 3 3 energy levels group 4 4 e 1- in outer (valence) shell 14 e 1- total, 14 p 1+ 2e 1- 8e 1- 4e 1- Outermost electrons “see” 14p 1+, 10e 1- Effective Nuclear Charge is 4+

22. 13p 1+ 2e 1- 8e 1- 4e 1- Larger effective nuclear charge makes a Si atom smaller than an Al atom

23. Now let’s look at a graph of Atomic Radius in 10 -10 m vs Atomic Number.

24. Atomic Radius vs Atomic Number Radius m x 10 -10 0.5 1.0 1.5 2.0 2.5 5101520 Atomic Number Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca

25. Here is a graph for the 1st 40 Elements

27. In general Atomic Radius tends to get smaller across the periodic table and larger down the periodic table. This can be summarized by

28. Smaller atomic radius Smaller atomic radius

29. Most of the trends in the Periodic Table can be explained by comparing atomic radii. For instance the trend in Electronegativity. Electronegativity is the relative attractive force of an atom for the electrons in a bond. Generally the smaller the atom the greater the electronegativity. This is because the positively charged nucleus is closer to the shared electrons in a smaller atom.

30. ++ Electronegativity is a measure of the attractive force + + Greater the distance the weaker the force

31. Smaller atomic radius Smaller atomic radius Increasing Electronegativity

32. Ionization energy is a measure of the amount of energy required to remove an electron from an atom. In general, the smaller the atom, the harder it is to remove its electrons. As an atom gets smaller the energy required to remove its outermost electron increases. This can be illustrated by the following:

33. + 2e 1- 8e 1- 1e 1- + 2e 1- 1e 1- Stronger attractive force Weaker attractive force

34. Smaller atomic radius Smaller atomic radius Increasing Electronegativity Increasing 1st ionization energy

35. Multiple Ionization Energies 12p 1+ 2e 1- 8e 1- 2e 1- energyenergy 1 st e 1- 2 nd e 1- 3 rd e 1- If multiple electrons are removed from a single atom the energy required changes in predictable ways.

36. Multiple Ionization Energies energyenergy 1 st e 1- 2 nd e 1- 3 rd e 1- After the 1st electron is removed the valence shell gets a little closer so it takes more energy to remove the 2nd electron. 12p 1+ 2e 1- 8e 1- 1e 1-

37. Multiple Ionization Energies energyenergy 1 st e 1- 2 nd e 1- 3 rd e 1- Since the 2nd shell is much closer than the valence shell it takes a lot more energy to remove an electron from this shell. 12p 1+ 2e 1- 8e 1-

38. Multiple Ionization Energies energyenergy 1 st e 1- 2 nd e 1- 3 rd e 1- 12p 1+ 2e 1- 7e 1-

39. Multiple Ionization Energies energyenergy 1 st e 1- 2 nd e 1- 3 rd e 1- 12p 1+ 2e 1- 7e 1- 4 th e 1-

40. Multiple Ionization Energies energyenergy 1 st e 1- 2 nd e 1- 3 rd e 1- 12p 1+ 2e 1- 6e 1- 4 th e 1-

41. When looking at a graph of multiple ionization energies the number of valence shell electrons can be deduced. If the element’s period is known then the element can sometimes be deduced. energyenergy 1 st e 1- 2 nd e 1- 3 rd e 1- 4 th e 1- If this atom comes from an element in period 4 what element is it? Ga Group 3

42. energyenergy 1 st e 1- 2 nd e 1- 3 rd e 1- 4 th e 1- If this atom comes from an element in period 5 what element is it? Rb Since the 4th ionization energy is not even 2x’s greater than the 3rd, and the 2nd is more than 2x’s bigger than the 1st this atom must only have 1 valence shell electron.

43. Electron Affinity is the amount of energy released when an atom gains an electron. In general the smaller the atom the greater the quantity of energy released when it gains an electron. This can be summarized as a trend using the Periodic table.

44. Smaller atomic radius Smaller atomic radius Increasing Electronegativity Increasing 1st ionization energy Increasing electron affinity

45. Metallic properties are due to the ease with which metal atoms lose their electrons. In general the more easily metal atoms lose electrons the more metallic its properties. The smaller a metal’s atoms are, the more tightly held are its electons, the less metallic its properties.

46. Smaller atomic radius Smaller atomic radius Increasing Electronegativity Increasing 1st ionization energy Increasing electron affinity Decreasing metallic properties

47. When metals react they usually do so by losing electrons. The more easily they lose electrons the more reactive they are. The smaller a metal’s atoms the less reactive they are. This results in the following trend for metallic reactivity:

48. Smaller atomic radius Smaller atomic radius Increasing Electronegativity Increasing 1st ionization energy Increasing electron affinity Decreasing metallic properties Decreasing metallic reactivity

49. Non-metals which don’t have completely filled outer shells can react by gaining electrons. In general the smaller these non-metallic element’s atoms are the stronger their attraction for electrons, the more reactive they are. In general reactive non-metals get more reactive from bottom to top, and more reactive from left to right in the peridic table.

50. Increasing non-metallic reactivity