1. Periodic Table

2. Atom Neutral atom has equal number of protons and electrons + charges = - charges Overall charge of zero (neutral)

3. Ions Atom can gain or lose electrons to form ion Gain electrons → negative ion (anion) + charges < - charges → - ion Lose electrons → positive ion (cation) + charges > - charges → + ion

5. Principal Energy Level Element# e - 1 2 34 H11 He22 Li32 1 Be42 2 C62 4 F92 7 Ne102 8 Na112 8 1 Mg122 8 2 Cl172 8 7 Ar182 8 8 K192 8 81

6. .

7. Periodic Table and Periodic Properties Periodic Law proposed by Mendeleev (1869): “Elements arranged according to their atomic weights (masses) show a distinct periodicity (regular variation) of their properties” Modern Periodic Law: Physical and chemical properties of elements are periodic functions of their atomic numbers

9. Mendeleev Arranged elements according to: 1.Properties 2.Atomic mass

10. When properties did not match up, Mendeleev left “holes” and assumed that the “missing” elements had not yet been discovered ScGaGe Atomic mass of I is126.90 amu Atomic mass of Te is 127.60 amu Assume that mass was wrong and better data would show new atomic masses Arranged them to match properties

11. .

12. Periods are rows (1-7) Properties change as you go across table LiBeBC N OFNe Groups are columns (A and B elements) Elements within a group (family) have similar properties Li Na K Rb Cs

13. .

14. Why do elements in group have similar properties? Need to look at a more detailed arrangement of electrons in atom Valence electrons are electrons in outermost shell (principal energy level)

15. Principal Energy Level Element# e - 1 2 34 H11 He22 Li32 1 Be42 2 C62 4 F92 7 Ne102 8 Na112 8 1 Mg122 8 2 Cl172 8 7 Ar182 8 8 K192 8 81

16. Octet A full octet (8 electrons) in outermost energy level is very stable Atoms try to achieve full octets Valence electrons are involved in reactions and in forming ions

17. Representative (Main Group) Elements - Group A elements Transition elements - Group B elements Alkali metals - Group IA Alkaline earth metals - group IIA Halogens - group VIIA Noble gases - group VIIIA

18. Group IA—Alkali Metals Soft metal—easy to cut Silvery color Very reactive Not found as free elements Have 1 outer shell electron—easily removed Form +1 ions Li NaKRbCsFr

19. Group IIA—Alkaline Earth Metals Harder and denser than Group IA Less reactive than Group IA Has two outer shell electrons Form 2+ ions Be MgCaSrBaRa

20. Group VIIA—Halogens Occur as diatomic molecules F2 and Cl 2 gases at room temperature Br 2 liquid at room temperature I 2 solid at room temperature Reactivity of elements F > Cl > Br > I Have 7 outer shell electrons—easily take on one more to get “full” octet Form -1 ions

21. Group VIII—Noble Gases Colorless gases at room temperature Exist as single atoms Have 8 outer electrons (full octet) Do NOT react with other elements He Ne Ar KrXeRn

22. Trends in the Periodic Table We will look at the following trends: Atomic radius (size of atom) Ion Size Ionization Energy Electron Affinity

23. Trends in Atomic Size The size of the atom increases as you go down a group in the periodic table More “shells” Higher energy levels are farther from nucleus The size of the atom decreases as you go across a period (from left to right) More protons in nucleus exert a greater “pull” on electrons

24. Size of atoms Numbers show radii in pm

25. Cation size Cations (positive ions) are always smaller than their parent atoms Formed by losing electrons from their outermost energy level Cations have more protons than electrons. Extra protons exert a greater “pull” on electrons that remain

26. Anion Size Anions are always larger than their parent atoms Anions have more electrons than protons, so the effective pull on the electrons is lower and electrons are held less strongly

27. Ion Size (Radii in pm)

28. Ionization Energy Ionization energy (I.E) is the energy required to remove an electron from an atom (in the gaseous state) I.E. + K →K + + e - The more strongly an electron is held by the nucleus, the harder it is to remove an electron and the higher the ionization energy needed to remove the electron

29. Trends in Ionization Energy Ionization energy decreases as you go down a group because the electrons are farther from nucleus and not held as tightly so they are easier to remove Ionization energy increases as you go across a period. More protons in the nucleus exert a greater pull on the electrons, so they are harder to remove

30. Ionization Energy FG07_006.JPG

31. Electron Affinity Electron Affinity is the energy released when a single electron is added to an isolated atom Br + e - → Br - + energy Can think of it as ability of an atom to attract an electron to itself Atoms with high electron affinity form anions that are more stable than atom alone

32. Trends in Electron Affinity Electron affinities generally increase as you go across a period (from left to right) (exception: Noble gases) Electron affinities generally decrease as you go down a group

34. Atoms which have the greatest tendency to lose or gain electrons to form full octet will be most reactive

35. .

36. Metals Tend to lose electrons and form positive ions Exist as solids at room temperature (except Hg) Most of known elements (87) Shiny, lustrous, malleable, ductile, Conduct heat and electricity

37. Nonmetals Tend to gain electrons and form negative ions Several are gases at room temperature. Others are solids. Br is liquid Often dull or brittle Poor conductors of heat or electricity

38. Metalloids (semi-metals) Properties intermediate between metals and nonmetals Semiconductors(Si, As, Ge in computers) Follow zigzag line between metals and nonmetals B SiGeAsSbTe(Po At)

39. Information from Periodic Table Atomic number→ 80 (protons) Atomic symbol → Hg Atomic mass → 200.6(amu) (weighted average)

40. Electronegativity Electronegativity is a measure of the ability of an atom to attract electrons to itself in a chemical bond Developed by Linus Pauling Arbitrary scale ranging from 0 to 4.0 (most electronegative)

41. General Values Most electronegative atom is F (E n = 4.0) –F has high ionization energy –F has high electron affinity –F attracts electrons to itself

42. Alkali metals such as Na have low electronegativity (E n = 0.9) –Tend to have low ionization energy –Tend to have low electron affinity –Easily lose electrons –Do not attract electrons very strongly

43. Noble gases have E n = 0 –Do not bond with other atoms –Have VERY high ionization energy

44. Representative Values F = 4.0O = 3.5 C = 2.5Cl = 3.0 H = 2.1N = 3.0 Na = 0.9

45. Electronegativity Values

46. Trends in Electronegativity Highest Electronegativity in upper right corner of periodic table Lowest Electronegativity in lower left corner of periodic table Noble gases have 0 electronegativity

47. Summary Electronegativity increases as you go from left to right on periodic table EXCEPT for Noble Gases E n = 0 Electronegativity decreases as you go down the periodic table