1. The Periodic Table

2. Organizing the Elements Demitri Mendeleev (Russian – 1869) Demitri Mendeleev (Russian – 1869)  Published the 1 st periodic table  Based on atomic mass

3. Modern Periodic Table Modern Periodic Table Developed by Henry Mosley Developed by Henry Mosley Organized according to Organized according to atomic number atomic number Periods = rows 1-7 Periods = rows 1-7 Elements in the same period have the same number of principal energy levels Elements in the same period have the same number of principal energy levels

4. Groups = columns 1-18 * also called “families” * also called “families” * elements have similar properties * elements have similar properties * elements have same # of valence electrons * elements have same # of valence electrons * may be numbered according to the “A” system – 1A have 1 valence e’, 2A have 2 system – 1A have 1 valence e’, 2A have 2 valence e’, 3A have 3 valence, etc. valence e’, 3A have 3 valence, etc. * remember, valence electrons are found in the outermost energy level (s and p only) and give an element its physical and chemical properties

5.  Which of these sets of elements have similar physical and chemical properties? a)oxygen, nitrogen, carbon, b)strontium, magnesium, calcium, beryllium c)nitrogen, neon,, fluorine  Name 2 elements that have properties similar to those of the element sodium. Any other Group 1 element: Lithium, Potassium, Rubidium, Cesium, Francium boron nickel

6. The Periodic Law Many of the physical and chemical properties of the elements tend to recur in a systematic manner with increasing atomic number. Many of the physical and chemical properties of the elements tend to recur in a systematic manner with increasing atomic number.

7. Period → Group ↓

8. Metallic Character Approx. 3/4 of the elements are metals Approx. 3/4 of the elements are metals Metallic Properties Metallic Properties Luster Luster Conductivity Conductivity Malleability Malleability Ductility Ductility Sectility Sectility Tend to lose electrons to become positive ions Tend to lose electrons to become positive ions Nonmetallic Properties Nonmetallic Properties Dull luster Dull luster Poor conductors Poor conductors Brittle Brittle Tend to gain electrons to become negative ions Tend to gain electrons to become negative ions Metalloids (semimetals) Metalloids (semimetals) Have some properties characteristic of metals and other properties characteristic of nonmetals Have some properties characteristic of metals and other properties characteristic of nonmetals

9. Metals and Nonmetals Li 3 He 2 C6C6 N7N7 O8O8 F9F9 Ne 10 Na 11 B5B5 Be 4 H1H1 Al 13 Si 14 P 15 S 16 Cl 17 Ar 18 K 19 Ca 20 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Ga 31 Ge 32 As 33 Se 34 Br 35 Kr 36 Rb 37 Sr 38 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48 In 49 Sn 50 Sb 51 Te 52 I 53 Xe 54 Cs 55 Ba 56 He 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Tl 81 Pb 82 Bi 83 Po 84 At 85 Rn 86 Fr 87 Ra 88 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 Mg 12 Ce 58 Pr 59 Nd 60 Pm 61 Sm 62 Eu 63 Gd 64 Tb 65 Dy 66 Ho 67 Er 68 Tm 69 Yb 70 Lu 71 Th 90 Pa 91 U 92 Np 93 Pu 94 Am 95 Cm 96 Bk 97 Cf 98 Es 99 Fm 100 Md 101 No 102 Lr 103 La 57 Ac 89 1 2 3 4 5 6 7   METALS Nonmetals Metalloids

10. Metallic Review Identify each element as a metal, nonmetal or metalloid Identify each element as a metal, nonmetal or metalloid a) Gold - b) Silicon - c) Sulfur - d) Barium - Metal Metalloid Nonmetal Metal

11. Identify each property below as more characteristic of a metal or a nonmetal Identify each property below as more characteristic of a metal or a nonmetal a) Brittle - b) Malleable - c) Poor conductor of electricity - d) Shiny - e) Tend to gain electrons -  In which pair of elements are the chemical properties of the elements most similar? a) sodium and chlorine b) nitrogen and phosphorus c) boron and oxygen Nonmetal Metal Nonmetal Metal Nonmetal

12. Classifying the Elements Alkali Metals – Alkali Metals – Alkaline Earth Metals - Group 2 (2A) Alkaline Earth Metals - Group 2 (2A) Boron Group - Group 13 (3A) Boron Group - Group 13 (3A) Carbon Group - Group 14 (4A) Carbon Group - Group 14 (4A) Nitrogen Group - Group 15 (5A) Nitrogen Group - Group 15 (5A) Oxygen Group – Group 16 (6A) Oxygen Group – Group 16 (6A) Halogens – Group 17 (7A) Halogens – Group 17 (7A) Noble Gases – Group 18 (8A) Noble Gases – Group 18 (8A) Representative Elements – Representative Elements – Transition Metals – All of the d-block (Group B) Transition Metals – All of the d-block (Group B) Inner Transition Metals - All of the f-block Inner Transition Metals - All of the f-block Group 1 (1A) Groups 1A – 8A

13. The Oxidation Number is the charge an element usually takes when the atoms become ions. The Oxidation Number is the charge an element usually takes when the atoms become ions. The oxidation number is positive if the element likes to give up electrons (metals do this). The oxidation number is positive if the element likes to give up electrons (metals do this). The oxidation number is negative if the element likes to receive electrons (non-metals do this). The oxidation number is negative if the element likes to receive electrons (non-metals do this). Losing or gaining electrons gives the atom a full outer shell Losing or gaining electrons gives the atom a full outer shell Elements in 1A like a +1 charge Elements in 1A like a +1 charge 2A like a +2 charge 2A like a +2 charge 3A like a +3 charge 3A like a +3 charge 4A like a +4 or a -4 charge 4A like a +4 or a -4 charge 5A like a -3 charge 5A like a -3 charge 6A like a -2 charge 6A like a -2 charge 7A like a -1 charge 7A like a -1 charge 8A have no charge – outer shell already filled 8A have no charge – outer shell already filled

14. Classifying Review Which of the following are symbols for representative elements Which of the following are symbols for representative elements Na, Mg,, Cl?  Which of these metals is not a transition metal? a) Aluminum b) Silver c) Iron d) Zirconium Fe, Ni

15. Atomic Radius Atoms are roughly spherical Atoms are roughly spherical Sphere size can be determined using the radius Sphere size can be determined using the radius Problem: Problem: Solution: half the distance between the nuclei of identical atoms that are bonded together Atomic Radii – half the distance between the nuclei of identical atoms that are bonded together Edges of orbitals are fuzzy and difficult to measure…

16.  Trends: 1. Across a period, the atomic radius → Why? Nuclei have larger positive charges Nuclei have larger positive charges Electrons are pulled in closer Electrons are pulled in closer 2. Down a group, the atomic radius → 2. Down a group, the atomic radius →Why? Greater number of energy shells Valence electrons are further away from the nucleus decreases increases

18. Atomic Radius Review Which element in each pair has a larger atomic radius? Which element in each pair has a larger atomic radius? a) & lithium b) & bromine c) carbon & d) & neon  Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum, and sodium. Sodium → Aluminum → Sulfur → Chlorine sodium calcium germanium fluorine

19. Ionization Energy  What is an ion? An electrically charged atom that has an unequal number of protons and electrons An electrically charged atom that has an unequal number of protons and electrons Atoms may either gain or lose electrons Atoms may either gain or lose electrons  Loss of electrons = (+) charge →  Gain of electrons = (-) charge →  Ionization energy = cation anion the amount of energy required to remove the most loosely bound electron A + energy  A + + e -

20. 1. Across a period, the ionization energy → Why? Electrons are held closer to the nucleus Electrons are more difficult to remove Electrons are more difficult to remove 2. Down a group, the ionization energy → Why? Electrons are further away from the nucleus Electrons are easier to remove Electrons are easier to remove increases decreases

22. Second Ionization Energy The energy required to remove a second electron from a cation The energy required to remove a second electron from a cation Energy increases for each successive electron removed Energy increases for each successive electron removed Larger than expected energy spikes occur when trying to remove an electron from a Noble gas configuration Larger than expected energy spikes occur when trying to remove an electron from a Noble gas configuration A + + energy  A 2+ + e -

23. Ionization Review Which element in each pair has a greater ionization energy? Which element in each pair has a greater ionization energy? a) lithium, b) magnesium, c) cesium,  Arrange the following groups in order of increasing ionization energy a) Be, Mg, Sr b) Bi, Cs, Ba c) Na, Al, S Which atom would show the largest increase between the 2 nd and 3 rd ionization energies? Which atom would show the largest increase between the 2 nd and 3 rd ionization energies? Na Mg Al Na Mg Al boron strontium aluminum Sr → Mg → Be Cs → Ba → Bi Na → Al → S

24. Ionic Size Cations form when Cations form when atoms lose e -. atoms lose e -. The cation is smaller than the atom from which it came The cation is smaller than the atom from which it came Anions form when Anions form when atoms gain e -. atoms gain e -. The anion is larger than the atom from which it came. The anion is larger than the atom from which it came.

25. Why is a cation smaller than the uncharged atom? When e - are lost, the attraction between the protons and remaining electrons increases. This means e - are pulled closer to the nucleus. Why is an anion larger than the uncharged atom? When e - are added, the protons can’t pull the electrons as tightly because there are more of them. This means the radius increases.

26. Ion Size Review Which particle has the larger radius in each atom/ion pair? Which particle has the larger radius in each atom/ion pair? a) Na, b) S, c) I, d) Al,  The ions Na + and Mg 2+ each have 10e -. Which ion would you expect to have a smaller ionic radius? Na + S 2- I-I-I-I- Al 3+

27. Electronegativity A measure of the ability of an atom to attract a pair of electrons when bonded to another atom A measure of the ability of an atom to attract a pair of electrons when bonded to another atom Which atom “wins the tug-o-war” in the HF molecule? Which atom “wins the tug-o-war” in the HF molecule?

28. 1. Across a period, the electronegativity → Why? Atoms are very close to a stable octet (full outer shell) They ‘pull’ electrons from other atoms 2. Down a group, the electronegativity → 2. Down a group, the electronegativity →Why? Top atoms have few electrons, and hold them tightly Bottom atoms have numerous electrons and little desire to acquire more increases decreases

30. Electronegativity Review Which element in each pair has a higher electronegativity value? Which element in each pair has a higher electronegativity value? a) Cl, b) C, c) Mg, d) As,  Which element in each pair has a greater attraction for electrons? a) Ca or b) O or c) S F N Ne Ca O F or K

31.  Metallic Character  1. Across a period, metallic character → Why? Why? More electrons in outer shells Less tendency to lose electrons 2. Down a group, metallic character → Why? Why? Fewer electrons in outer shells Greater tendency to lose electrons decreases increases

32. Decreases Increases Electronegativity DecreasesIncreases Ionization Energy IncreasesDecreases Atomic Radius IncreasesDecreases Metallic Character Variation down a Group Variation across a Period Periodic Property Summary

33. This periodic table has the f-block inserted in the middle.

34. Chinese Periodic Table

38. Stowe Periodic Table

39. A Spiral Periodic Table

40. Triangular Periodic Table

41. “Mayan” Periodic Table

42. Giguere Periodic Table

43. Orbital filling table