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THE PERIODIC TABLE & ELECTRON CONFIGURATION

THE PERIODIC TABLE & ELECTRON CONFIGURATION. Chapters 4 & 5. The Element Song. Dimitri Mendeleev. Invented periodic table Organized elements by properties Arranged elements by atomic mass Predicted existence of several unknown elements Element 101 . Mendeleev’s Early Periodic Table.

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THE PERIODIC TABLE & ELECTRON CONFIGURATION

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

  2. The Element Song

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

  4. Mendeleev’s Early Periodic Table TABELLE II GRUPPE I GRUPPE II GRUPPE III GRUPPE IV GRUPPE V GRUPPE VI GRUPPE VII GRUPPE VIII ___ ___ ___ ___ RH4 RH3 RH2 RH R2O RO R2O3 RO2 R2O5 RO3 R2O7 RO4 REIHEN 1 2 3 4 5 6 7 8 9 10 11 12 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 __ __ __ __ __ From Annalen der Chemie und Pharmacie, VIII, Supplementary Volume for 1872, p. 151.

  5. Examples of Mendeleev’s

  6. 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

  7. 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 calledperiods • Elements with similar chemical properties form vertical columns, calledgroups, • 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

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

  9. World of Chemistry

  10. IONS • Positive and negative ions form when electrons are transferred between atoms • Cation: an ion with a + charge • Example: Na+ Ca2+ • Anion: an ion with a – charge • Example: O2- F-

  11. 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

  12. Figure 6.22 SUMMARY OF PERIODIC TRENDS

  13. 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

  14. 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!)

  15. 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.

  16. 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

  17. Bohr Model of the Atom (don’t copy this slide) • Interactive Bohr Model

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

  19. 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)

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

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

  22. 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

  23. 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

  24. Sublevels in Atoms • See Fig 7.5 on p 235

  25. 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

  26. 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”

  27. 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)

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

  29. 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

  30. 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

  31. 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: 1s1 Helium: 1s2 Lithium: 1s22s1 Oxygen: 1s22s22p4

  32. 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: 1s22s22p4 • There are 6 valence e-s in the 2nd main energy level (valence level)

  33. 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.

  34. Exceptions: Some Transition Elements (don’t need to copy) • Titanium - 22 electrons NORMAL • 1s22s22p63s23p64s23d2 • Vanadium - 23 electrons NORMAL • 1s22s22p63s23p64s23d3 • Chromium - 24 electrons EXCEPTION • 1s22s22p63s23p6 4s2 3d4is expected • But this is wrong!!

  35. Chromium is actually… (copy this!) • 1s22s22p63s23p63d54s1 • 3d54s1Instead of 4s2 3d4 • There is less repulsion (lower energy) in the 2nd arrangement 4s 3d

  36. Noble Gas Notation • Short-cut way of showing e- configuration • A Noble Gas is a Group 18 element. • Identify the noble gas in the period above your element of interest. Write this symbol in brackets. • Write the e- configuration for any additional e-s that your element of interest has, but the noble gas doesn’t have. Ex: Nitrogen: 1s22s22p5 becomes [He] 2s22p5

  37. 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

  38. 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)

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