1. THE PERIODIC TABLE & ELECTRON CONFIGURATION Chapters 4 & 5

3. The Element Song

4. Dimitri Mendeleev  Invented periodic table  Organized elements by properties  Arranged elements by atomic mass  Predicted existence of several unknown elements  Element 101

5. Mendeleev’s Early Periodic Table GRUPPE I GRUPPE II GRUPPE III GRUPPE IV GRUPPE V GRUPPE VI GRUPPE VII GRUPPE VIII ___ ___ ___ ___ RH 4 RH 3 RH 2 RH R 2 O RO R 2 O 3 RO 2 R 2 O 5 RO 3 R 2 O 7 RO 4 REIHEN 1 2 3 4 5 6 7 8 9 10 11 12 Annalen der Chemie und Pharmacie From Annalen der Chemie und Pharmacie, VIII, Supplementary Volume for 1872, p. 151. H = 1 Li = 7 Be = 9.4 B = 11 C = 12 N = 14 O = 16 F = 19 Na = 23 Mg = 24 Al = 27.3 Si = 28 P = 31 S = 32 Cl = 35.5 K = 39 Ca = 40 ? = 44 Ti = 48 V = 51 Cr = 52 Mn = 55 Fe = 56, Co = 59, Ni = 59, Cu = 63 (Cu = 63) Zn = 65 ? = 68 ? = 72 As = 75 Se = 78 Br = 80 Rb = 85 Sr = 87 ? Yt = 88 Zr = 90 Nb = 94 Mo = 96 __ = 100 Ru = 104, Rh = 104, Pd = 106, Ag = 108 (Ag = 108) Cd = 112 In = 113 Sn = 118 Sb = 122 Te = 125 J = 127 Cs = 133 Ba = 137 ? Di = 138 ?Ce = 140 __ __ __ __ __ __ __ ( __ ) __ __ __ __ __ __ __ __ ? Er = 178 ? La = 180 Ta = 182 W = 184 __ Os = 195, Ir = 197, Pt = 198, Au = 199 (Au = 199) Hg = 200 Tl= 204 Pb = 207 Bi = 208 __ __ __ __ __ Th = 231 __ U = 240 __ __ __ __ __ TABELLE II

6. Examples of Mendeleev’s

7. Modern Periodic Table  Henry G.J. Moseley  Determined the atomic numbers of elements from their X -ray spectra (1914)  Arranged elements by increasing atomic number

8. Modern Periodic Table  Elements are arranged in seven horizontal rows, in order of increasing atomic number from left to right and from top to bottom  Rows are called periods  Elements with similar chemical properties form vertical columns, called groups,  Groups 1, 2, and 13 through 18 are the main group elements  transition elements: groups 3 through 12 are in the middle of the periodic table  Inner transition elements: The two rows of 14 elements at the bottom of the periodic are the lanthanides and actinides

9. Groups to Know  Group 1 = Alkali Metals  Group 2 = Alkaline Earth Metals  Group 17 = Halogens  Group 18 = Noble Gases

10. World of Chemistry

11. IONS  Positive and negative ions form when electrons are transferred between atoms  Cation: an ion with a + charge  Example: Na + Ca 2+  Anion: an ion with a – charge  Example: O 2- F -

12. ELECTRONEGATIVITY  Electronegativity describes how electrons are shared in a compound  The high number means the element has a greater pull on electrons  Fluorine is the most electronegative element

13. SUMMARY OF PERIODIC TRENDS Figure 6.22

14. Light, Energy, and Electrons  e-s are arranged in energy levels (e.l.’s), at different distances from nucleus  Close to nucleus = low energy  Far from nucleus = high energy

15. Rules for “placing” e-s in energy levels  e-s in highest occupied level are “valence e-s”  Only so many e-’s can fit in a particular e.l.  e-s fill lower e.l.’s before being located in higher e.l.’s* Ground state is the lowest energy arrangement of e-s. * There are exceptions we will learn later!)

16. Light, Energy, and Electrons  e-s can jump to higher energy levels if they absorb energy.  They can’t keep the energy so they lose it and “fall back” to lower levels.  When they do this, they release the energy they absorbed in the form of light.

17. Light, Energy, and Electrons (See p 129 of text ChemI/IH) Electron energy levels are like rungs of a ladder. Ladder – To climb to a higher level, you can’t put your foot at any level, – you must place it on a rung Electron energy levels – e-s must also move to higher or lower e.l.’s in specific intervals

18. Bohr Model of the Atom (don’t copy this slide)  Interactive Bohr Model Interactive Bohr Model

19. Light, Energy, and Electrons  Quantum-the amount of energy required to move an electron from one E.L. to another.

20. Atomic Emission Spectrum (A.E.S)  Each element emits a color when its excited e- s “fall back.”  Pass this light thru a prism, it separates into specific lines of color.  You can identify an element by its emission spectrum! (no 2 elements have the same AES)

21. Emission Spectra of H, He, Ne (don’t copy this slide)

22. Use of e- waves (don’t copy this slide)  Electron microscope magnifies tiny objects b/c e- wavelength much smaller than visible light snowflake

23. Heisenburg Uncertainty Principle  Def: if you want to locate something, you can shine light on it  When you do this to an electron, the photons send the e- off in an unpredictable direction  (def):Therefore, you can never know BOTH the position and velocity of an e- at the same time

24. Electron Sublevels Each electron has an “address,” where it can be considered to be located in the atom.  Main energy level (principal quantum #) = “hotel”  Sublevel = “floor”  Orbital = “room”  Regions of space outside the nucleus  All orbitals in a sublevel have the same energy  2 electrons max can fit in an orbital

25. Sublevels in Atoms  See Fig 7.5 on p 235 Main energy level Types of sublevels # of orbitals# of electrons 1s1 2spsp 1 3 (4 total) 3spdspd 1 3 5 (9 total) 4-7spdfspdf 1 3 5 7 (16 total)

26. Orbitals  s orbitals are spherical  There is only 1 orbital  p orbitals are dumbbell shaped  There are 3 orbitals, all with = energy  Each is oriented on either x, y, or z axis  They overlap  d orbitals have varying shapes  There are 5 orbitals, all with = energy  f orbitals have varying shapes  There are 7 orbitals, all with = energy

27. Electron Configurations (don’t have to copy. Info in prior slide)  Electrons are always arranged in the most stable (lowest energy) way  This is called“electron configuration” or “ground state”

28. The Periodic Table & Atomic Structure  Shape of p. table is based on the order in which sublevels are filled REGIONS OF THE P. TABLE (see p 244 of book)  s REGION (“block”) - Groups 1 & 2  p REGION (block) - Groups 13-18  d REGION (block)- Groups 3-12 (Transition Elements)  f REGION (block)- (Inner Transition Elements)

29. Regions or “Blocks” of the P. Table (don’t need to copy)

30. Writing e- Configurations for Elements Using the P. Table 1. Always start with Period 1-go from L to R. 2. Go to Period 2-from L to R 3. Go to Period 3- from L to R 4. Continue w/Periods #4-7, L to R, until you arrive at the element you are writing e- configuration for.  Exception: elements in d block are 1 main E.L lower than the period where they are located  Exception: elements in f block are 2 main E.L.s lower than the period where they are located

31. Correct Order of Sublevels (lowest to highest energy)  1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

32. e- configurations 1. Use the P. Table to write the sublevels in increasing order. 2. Add a superscript next to each sublevel that shows how many e-s are in the sublevel 3. Ex: Hydrogen: 1s 1 Helium: 1s 2 Lithium: 1s 2 2s 1 Oxygen: 1s 2 2s 2 2p 4

33. Identifying Valence e-s  Valence e-s are the electrons in the highest occupied main energy level. (don’t copy. In prior slide)  Identify them by finding the “biggest big number” in your e- configuration. Ex: Oxygen: 1s 2 2s 2 2p 4  There are 6 valence e-s in the 2nd main energy level (valence level)

34. Why are d & f block elements’ sublevels out of order?  When you get to the higher main E.L.’s, the sublevels begin to overlap.

35. Exceptions: Some Transition Elements (don’t need to copy)  Titanium - 22 electronsNORMAL  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2  Vanadium - 23 electronsNORMAL  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3  Chromium - 24 electronsEXCEPTION  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 is expected  But this is wrong!!

36. Chromium is actually… (copy this!)  1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1  3d 5 4s 1 Instead of 4s 2 3d 4  There is less repulsion (lower energy) in the 2 nd arrangement 4s3d

37. Noble Gas Notation  Short-cut way of showing e- configuration  A Noble Gas is a Group 18 element. 1.Identify the noble gas in the period above your element of interest. Write this symbol in brackets. 2.Write the e- configuration for any additional e-s that your element of interest has, but the noble gas doesn’t have. Ex: Nitrogen: 1s 2 2s 2 2p 5 becomes [He] 2s 2 2p 5

38. Arrow Orbital Diagram- Used to show e- configuration. SYMBOLS:  A box represents an orbital  Label each box with the sublevel: 1s 2s 2p 2p 2p  An arrow represents an electron  2 arrows (e-s) in the same orbital face opposite directions.  Example: oxygen, see above ↑ ↓ ↑↑

39. Arrow Orbital Diagram- Used to show e- configuration, cont. INSTRUCTIONS:  Fill electrons from lowest to highest sublevel.  Never place 2 e-s in the same orbital of a sublevel until you have placed one in each of the orbitals (Hund’s Rule)