1. The History of the Modern Periodic Table See separate slide show for Periodic Table History

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

3. Chemical Reactivity Families  Similar valence e - within a group result in similar chemical properties Alkali Metals Alkaline Earth Metals Transition Metals Halogens Noble Gases

4. Periodic Table Reveals Periodic Trends Effective Nuclear charge atomic size or radius ionization energy electron affinity electronegativity metallic character Reactivity bonding characteristics crystal configurations acidic properties densities Melting/Boiling points

5. Electron screening or shielding Electrons are attracted to the nucleus Electrons are repulsed by other electrons Electrons would be bound more tightly if other electrons weren’t present. The net nuclear charge felt by an electron is called the effective nuclear charge ( Z eff ).

6. Quantum Mechanical Model Z eff is lower than actual nuclear charge. Z eff increases toward nucleus ns > np > nd > nf This explains certain periodic changes observed.

7. Effective Nuclear Charge ( Z eff ) The effective nuclear charge acting on an electron equals the number of protons in the nucleus, Z, minus the average number of electrons, S that are between the nucleus and the electron in question. Z eff = # protons  # shielding electrons Z eff = attractive forces  repulsive forces Z eff = Z  S

8. For Example, Lithium vs. Carbon Li Z eff = 3  2 = 1 C Z eff = 6  2 = 4 So, carbon has a much smaller atomic radius compared to lithium: R carbon =77 pm R lithium = 152 pm When moving across a row: The greater the Z eff value, the smaller the atom’s radius.

9. Trend #1 Atomic Radii Increases to Left and Down Why larger going down? Why smaller to the right? Higher energy levels have larger orbitals Shielding - core e - block the attraction between the nucleus and the valence e - Increased nuclear charge without additional shielding pulls e - in tighter

10. Practice… Referring to a periodic table, arrange the following atoms in order of increasing size: – Phosphorus – Sulfur – Arsenic – Selenium S < P < Se < As

11. Atomic radii

12. The Periodic Table & Radii

13. Periodic Trend is Due to Effective Nuclear Charge Atomic Radii vs. Zeff:

14. Trends in Ionic Radii Using your knowledge of Z eff, how would the size of a cation compare to neutral atom? Anion?

15. Trends in Ionic Radii The cation of an atom decreases in size. The more positive an ion is, the smaller it is because Z eff increases The anion of an atom increases in size. The more negative an ion, the larger it is because Z eff decreases.

16. Cations  lose electrons, become smaller

17. Anions  gain electrons, become bigger

18. Ion Radii +3 +4 -3 -2 -1 Increases down Increases moving across, but depends if cation OR anion

19. Ions and Ionic Radii

21. Practice… Arrange the following atoms and ions in order of decreasing size: – Mg 2+ – Ca 2+ – Ca Which of the following ions is the largest: – S 2- – S – O 2-

22. Practice… Arrange the following ions in order of decreasing size: – S2 - – Cl - – K + – Ca 2+ Which of the following ions is the largest? – Rb + – Sr 2+ – Y 3+

23. Trend in Ionization Energy Ionization NRG is the NRG required to remove an electron from an atom

24. Successive Ionization NRG Ionization energy increases for successive electrons from the same atom.

25. *Notice the large jump in ionization energy when a core e  is removed. Why do you think there is such a big jump for Mg 3+ ?

26. The smaller the atom, the higher the ionization energy due to Z eff Bigger atoms have lower ionization NRG due to the fact that the electrons are further away from the nucleus and therefore easier to remove.

27. Increases Decreases

28. Practice… Which of the following elements would have the highest second ionization energy? Justify your answer. – Sodium, Sulfur, or Calcium Which will have the greater third ionization energy, Ca or S? Justify your answer.

29. Practice… Referring to a periodic table, arrange the following atoms in order of increasing first ionization energy (Ne, Na, P, Ar, K) Justify your answer. Based on the trends discussed in this section, predict which of the following atoms (B, Al, C or Si) has the lowest first ionization energy and which has the highest first ionization energy.

30. Electron Affinity The energy change associated with the addition of an electron Tends to increase across a period Tends to decrease as you go down a group Abbreviation is E ea, it has units of kJ/mol. Values are generally negative because energy is released. Value of E ea results from interplay of nucleus electron attraction, and electron–electron repulsion.

31. Ionization NRG vs. Electron Affinity Ionization energy measures the ease with which an atom loses an electron Electron affinity measures the ease with which an atom gains an electron

32. Electron Affinity

33. Trends in Electronegativity tendency for an atom to attract electrons when it is chemically combined with another atom. decreases as you move down a group increases as you go across a period from left to right.

34. Trend #5 Metallic Character The metallic character of atoms can be related to the desire to lose electrons. The lower an atom’s ionizatoin energy, the greater its metallic character will be. On the periodic table, the metallic character of the atoms increase down a family and decreases from left to right across a period.

35. MetalsNonmetals Shiny Luster Various colors (most silvery) Solids are malleable and ductile Good conductors of heat and electricity Most metal oxides are ionic solids that are basic Tend to form cations in aqueous solution No luster Various colors Brittle solids Poor conductors of heat and electricity Most nonmetal oxides are molecular substances that form acidic solutions Tend to form anions or oxyanions in aqueous solution

36. Metallic Character Increases moving down and across to the left Fr CsBa Ra Lower left corner -- elements most likely to lose their valence electrons Rb

37. Metals and Nonmetals Low ionization energies of metals means they tend to form cations (positive ions) relatively easily Due to their electron affinities, nonmetals tend to gain electrons when they react with metals.

38. # 6 Melting/Boiling Points Highest in the middle of a period (generally).

39. Some Important Properties of Alkali Metals Soft metallic solids Easily lose valence electrons (Reducing Agents) – React with halogens to form salts – React violently with water Large Hydration NRG – Positive ionic charge makes ions attractive to polar water molecules

40. Alkaline Earth Metals… Harder and more dense than Alkali Metals Less reactive than alkali metals (lower first ionization energies) Reactivity increases as you move down the periodic table.

41. The Halogens… “Salt Formers” Melting and Boiling Points increase with atomic number. Highly negative electron affinities Tendency to gain electrons and form halide ions

42. Noble Gases … Monoatomic ions Gases at room temperature Large 1 st ionization energies “Exceptionally” unreactive

43. Practice… Look at Sample Integrative Exercise 7 on page 264