Periodic Table

1. Periodic Table The most useful tool in the Lab

2. Early Organization • J.W. Dobereiner (1829) organized elements in triads • Triad – three elements with similar properties (ex: Cl, Br, I) • J.R. Newlands (1864) organized elements in octaves • Octave – repeating group of 8 elements

3. Development of the PeriodiceTable • Dmitri Mendeleev taught chemistry in terms of properties. • Mid 1800’s - molar masses of elements were known. • Wrote down the elements in order of increasing mass. • Found a pattern of repeating properties.

4. Mendeleev’s Table • Grouped elements in columns by similar properties in order of increasing atomic mass. • Found some inconsistencies - felt that the properties were more important than the mass, so switched order. • Also found gaps. • Must be undiscovered elements. • Predicted their properties before they were found.

5. The Modern Periodic Table • Henry Moseley – British physicist • Arranged elements according to increasing atomic number • The arrangement today • Symbol, atomic number & mass

6. The New Way • Elements are still grouped by properties. • Similar properties are in the same column. • Order is by increasing atomic number. • Added a column of elements (noble gases) • Weren’t found because they are unreactive.

7. Organization • Horizontal rows = periods • There are 7 periods • Each period represents an energy level • Every element in the same period has • the same # of energy levels and • the same core electron configuration

8. Organization • Vertical column = group or family • Similar physical & chemical prop. • Same # of valence electrons • Same common oxidation state • Identified by number & letter

9. Horizontal rows are called periods • There are 7 periods

10. Group 1A are the alkali metals • Group 2A are the alkaline earth metals

11. Group 7A is called the Halogens • Group 8A are the noble gases

12. These are called the inner transition elements, and they belong here The group B are called the transition elements

13. 1A 8A 2A 3A 4A 5A 6A 7A • The elements in the A groups are called the representative elements outer s or p filling

14. Lanthanides – the 4f orbital fills for these elements

15. Actinide series – the 5f orbitals are being filled for these elements.

16. Types of elements • Metals • Non-metals • Metalloids or semi-metals

17. Metals • Good conductor of heat and electricity • Malleable • Ductile • High tensile strength • High luster • Solid at room temperature • React by losing electrons

18. Nonmetals • Poor conductors of heat and electricity • React by gaining electrons • Some gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br)

19. Semi-Metals • Heavy, stair-step line • Metalloids border the line • Properties intermediate between metals and nonmetals • Learn the general behavior and trends of the elements, instead of memorizing each element property • B, Si, Ge, As, Sb, Te

20. Families • Group IA – alkali metals • most reactive metals • Silvery in appearance • Soft • Combine easily with non-metals • Melting point is higher than the boiling point of water • Have 1 valence electron

21. Families • Group 2 – Alkaline Earth Metal Family • Harder, stronger, denser, higher melting point, and less reactive than alkali • Usually not found as free elements, but as compounds • Have 2 valence electrons

22. Families • Group 7 – Halogens • Most reactive family • Non-metals • Have seven valence electrons • Group 8 – Noble Gas • Inert, unreactive • Have full set of valence electrons

23. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

24. S- block Alkali metals all end in s1 Alkaline earth metals all end in s2 really should include He, but it fits better later. He has the properties of the noble gases. s1 s2

25. He 2 Ne 10 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 Ar 18 Kr 36 Xe 54 Rn 86

26. p1 p2 p6 p3 p4 p5 The P-block

27. 1 2 3 4 5 6 7 • Each row (or period) is the energy level for s and p orbitals.

28. Areas of the periodic table • Group A elements = s & p blocks • representative elements • Wide range of phys & chem prop.

29. Transition Metals -d block s1 d5 s1 d10 d1 d2 d3 d5 d6 d7 d8 d10

30. d orbitals fill up after previous energy level, so first d is 3d even though it’s in row 4. 1 2 3 4 5 6 7 3d

31. f6 f13 f1 f2 f3 f4 f5 f7 f8 f10 f12 f14 f11 f9 F - block • inner transition elements

32. 1 2 3 4 5 6 7 • f orbitals start filling at 4f 4f 5f

33. Atomic Size } Radius • First problem: Where do you start measuring from? • The electron cloud doesn’t have a definite edge. • Atomic Radius = half the distance between two nuclei of a diatomic molecule.

34. 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 (blocking effect)

35. WHAT HAPPENS TO ATOMIC RADII? • Does a negative ion (anion) get larger or smaller? • Does a positive ion (cation) get larger or smaller?

36. Trends in Ionic Size • Cations form by losing electrons. • Cations are smaller than the atom they come from. • Metals form cations. • Cations of representative elements have noble gas configuration.

37. Ionic size • Anions form by gaining electrons. • Anions are bigger than the atom they come from. • Nonmetals form anions. • Anions of representative elements have noble gas configuration.

38. WHAT IS IONIZATION ENERGY? • The energy required to remove an electron • Which element has the highest ionization energy? Why?

39. What determines Ionization Energy? • The greater the nuclear charge, the greater IE. • Greater distance from nucleus decreases IE • All the atoms in the same period have the same energy level. • But, increasing nuclear charge • So IE generally increases from left to right.

40. Ionization Energy • The energy required to remove the first electron is called the first ionization energy • Thesecond ionization energy is the energy required to remove the second electron. • Always greater than first IE. • The third IE is the energy required to remove a third electron. • Greater than 1st or 2nd IE.

41. Driving Force • Full Energy Levels require lots of energy to remove their electrons. • Noble Gases have full orbitals. • Atoms behave in ways to achieve noble gas configuration.

42. WHAT IS ELECTRONEGATIVITY? • The ability of an atom to pull off an electron. • Which element has the highest electronegativity? Why?

43. 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.

44. Trends in Electron Affinity • The energy change associated with adding an electron to a gaseous atom. • Easiest to add to group 7A. • Gets them to full energy level. • Increase from left to right: atoms become smaller, with greater nuclear charge. • Decrease as we go down a group.