Periodic Properties - PowerPoint PPT Presentation

periodic properties n.
Skip this Video
Loading SlideShow in 5 Seconds..
Periodic Properties PowerPoint Presentation
Download Presentation
Periodic Properties

play fullscreen
1 / 41
Periodic Properties
Download Presentation
Download Presentation

Periodic Properties

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Periodic Properties

  2. Periodic Trends

  3. Don’t Confuse a trend with the explanation for that trend!!!!! • Example- • Why is Br atom larger than a Kr atom????? • Why is a cation smaller than its parent atom? • Why is an anion larger than its parent atom??

  4. Ionization energy • Endothermic –requires energy to remove outermost electron • Decreases down a group- large atomic radium and sheilding effect means that less energy needed to remove electrons • Increases across a period , the effextive nuclear charge increases, which causes an increased attraction between valence elctron and proton

  5. Ionization Energy- Why drop Be to B and N to O

  6. Ionization Energy Exceptions

  7. What happens to the IE as increasing number of electrons are removed from an atom? Where do the “jumps” occur? • I1 < I2 < I3

  8. Electron Affinity • Energy released when an atom attracts an addition electron • Negative- exothermic process • More negative, the greater the affinity • Elements in Group 2, 15 and 18 have lower than expected EA because of full (s) or ½ full (diamagnetic)(p) sublevels-electron=electron repulsion

  9. Sample MC • Questions 1-3refer to the following elements • O • La • Rb • Mg • N • What is the most electronegative element of the above? • Which element exhibits the greatest number of different oxidation states? • Which of the elements above has the smallest ionic radius for its most commonly found ion?

  10. Sample MC 4. The effective nuclear charge experienced by the outermost electron of Na is different than the effective nuclear charge experienced by the outermost electron of Ne. This difference best accounts for which of the following? • Na has a greater density at standard conditions than Ne. • Na has a lower first ionization energy than Ne. • Na has a higher melting point than Ne. • Na has a higher neutron-to-proton ration than Ne. • Na has fewer naturally occurring isotopes than Ne.

  11. MC Sample 5

  12. Chemical Bonding • Types: ionic- cation and anion • covalent- nonmetal+ nonmetal • Metallic- metal + metal- cations surrounded by a sea of electrons, valence electrons are delocalized

  13. Ionic • Lattice energy- amount of energy needed to separtate 1 mole of an ionic compound into its gaseous ions • Lattice energy increases with: smaller, more highly charged atoms—Coloumb’s law • E= k Q1 Q2/d Q1 and Q2 are charges on ions, d is distance between ions

  14. Covalent • Electronegativity difference determines polarity- 0-0.3; 0.3-1.7, 1.7 and greater • Dipole- molecule having one end with a slight + charge and the otherend with a slight negative charge

  15. Lewis Structures • 1. Count number of valence electrons • 2. Find central atom • 3. place outer atoms around central atom with single bond, place valence lectrons to give them an octet • 4. Determine if there are enough electrons to give central atom octet. If not, multiple bonds! Try Sulfur trioxide!

  16. Formal Charge • Determines best lewis structure. • Equal to # VE-# nonbonding electrons-1/2 number of bonding electrons • Formal charges of 0 means more likely Lewis structure • Example CO2

  17. Exceptions to Octet Rule • Central atom is Be (4 VE) BeI2 B(6 VE) BF3 • Species has an odd # electrons, one atom will have only 7 electrons (NO) • Expanded octet: Larger nonmetal central atom (period 3 and beyond) have d orbitals available for bonding. Also, the larger the atom, the more atoms that can fit around it: • SF6, PCl5, XeF4

  18. Sigma and pi bonds • Single bond—sigma • Double bond—one sigma and one pi bond • Triple bond- one sigma, and 2 pi bonds

  19. MC practice (A) CaO (B) CH2CH2 (C2H4) (C) SeO3 (D) CH2Cl2 (E) NBr3 1. The molecule with only one pi bond 2. The molecule whose molecular geometry is different from its electron geometry 3. The molecule that has trigonal pyramidal geometry

  20. MC practice • 4.Which of the following molecules has the shortest bond length? • A) N2 • B) O2 • C) Cl2 • D) Br2 • E) I2

  21. MC practice • 5. For which of the following molecules is the concept of resonance most likely to describe the bonding satisfactorily? A) BeF2 B) NO2 - C) CO2 D) CN- E) PCl5

  22. MC Practice 6. CCl4, CO2, PCl3, PCl5, SF6 Which of the following does not describe any of the molecules above? A) Square planar B) Tetrahedral C) Trigonal pyramidal D) Linear E) Octahedral

  23. MC Practice 7. All of these molecular shapes can be explained by sp3d hybridization of electrons on the central atom except A) linear B) a square pyramid C) see-saw D) a trigonalbipyramid E) an octahedron

  24. MC Practice 8. The melting point of MgO is higher than that of NaF. Explanations for this observation include which of the following? I. Mg2+ is more positively charged than Na+ II. O2– is more negatively charged than F– III. The O2– ion is smaller than the F– ion A) II only B) I and II only C) I and III only D) II and III only E) I, II, and III

  25. Electrons and wave properties

  26. Electrons and wave properties

  27. Calcuating among wavelength frequency and energy

  28. Balmer series Hydrogen BLS

  29. A line having a wavelength of 656 nm exists in the atomic emission spectrum of hydrogen • For the line, calcuate the: • A. frequency • B. energy of the photon • C. energy of a mole of photons • D. discuss th origin of the line in terms of the Bohr model of the atom

  30. Electron Configuration Rules

  31. Periodic table electron configuration

  32. MC Practice • Questions 1-4 refer to the following electron configurations. • 1s2 2s2 2p5 3s2 3p5 • 1s2 2s2 2p6 3s2 3p6 • 1s2 2s2 2p6 2d10 3s2 3p6 • 1s2 2s2 2p6 3s2 3p6 3d5 • 1s2 2s2 2p6 3s2 3p6 3d3 4s2 • The ground-state configuration for the atoms of a transition element. • The ground-state configuration of a negative ion of a halogen. • The ground-state configuration of a common ion of an alkaline earth element. • An impossible electronic configuration.

  33. MC Practice • 1s ____ 2s  • 1s  2s  • 1s  2s  2p  • 1s  2s  2p  • [Ar] 4s  3d  5. Represents an atom that has four valence electrons. 6. Represents an atom that is chemically unreactive. 7. Represents an atom in an excited state

  34. MC Practice 8. For which of the following electron transitions for a hydrogen atom does the light emitted have the longest wavelength? • n=5 to n=4 • n=5 to n=2 • n=3 to n=2 • n=5 to n=3 • n=4 to n=3

  35. MC Practice 9. Sr, Cr, Co, Hg, P • Gaseous atoms of which of the elements above are paramagnetic? • Sr and P only • Hg and P only • Mg, Cr, and Co only • Cr, Co, and P only • Cr, Co, and Hg only