1 / 11

8.4 Bond Polarity and Electronegativity.

8.4 Bond Polarity and Electronegativity.

lotte
Download Presentation

8.4 Bond Polarity and Electronegativity.

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. 8.4 Bond Polarity and Electronegativity. The concept of electronegativity was developed by Linus Pauling. Electronegativity is the ability of an element to attract electrons to itself in a molecule. Electronegativity increases across the periodic table and is at a maximum in the top right hand corner at fluorine, and is at a minimum at the bottom left hand corner at Cesium. Linus Carl Pauling (1901-1994)

  2. ELECTRONEGATIVITY Pauling originally developed the concept from the fact that for ionic compounds the bond energies were much larger for e.g. HF, than expected from the average of the energies for the related homonuclear diatomics, in this case H2 and F2. The more the observed energy of bond formation exceeded the average of the energies of the two related homo-nuclear diatomics, the greater the electronegativity. e-density high e-density spread equally e-density low Covalent – electron Polar covalent – Ionic – one atom Density spread equally one atom has more has attracted most of Over both atoms electron density the electron density

  3. Electronegativities Li to F On the next slide we have a table of electronegativities for elements in the periodic table. One sees that F (EN = 4.0) is the most electronegative element while Cs is the least electronegative (EN = 0.7) The electronegativites increase across the periodic table from Li (EN = 1.0) to Li by 0.5 per element, so that we have: Li Be B C N O F EN: 1.0 1.5 2.0 2.5 3.0 3.5 4.0

  4. Electronegativities of the Elements Cs (EN = 0.7) is least electronegative element F with EN = 4.0 is most electronegative element Au is at the peak of an island of electronegativity, and is most electronegative metal

  5. Electronegativities of some main group elements: H 2.1 Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Ca Ga Ge As Se Br 0.8 1.0 1.6 1.8 2.0 2.4 2.8 Rb Sr In Sn Sb Te I 0.8 1.0 1.7 1.8 1.9 2.1 2.5

  6. Electronegativity and bonding: Electronegativity tells us what kind of bonding we have, i.e. whether it is ionic or covalent. The greater the difference in EN between the two elements forming the bond, the more ionic is the bond. Typical ranges for EN differences are: EN difference bonding type Example EN difference range __________________________________________________________________________________ > 2.0 Ionic LiF 4.0-1.0 = 3.0 0.5-2.0 polar covalent HF 4.0-2.1 = 1.9 <0.5 covalent F-F 4.0-4.0 = 0.0 covalent C-H 2.5-2.1 = 0.4 covalent Li-Li 1.0-1.0 = 0.0 covalent Au-C 2.5-2.4 = 0.1 __________________________________________________________________________________

  7. Relativistic effects. One notes that electronegativity (EN) is at a maximum at F and a minimum at Cs, and increases from left to right, and from bottom to top in the periodic table. An important exception is the ‘island’ of high EN centered on Au. This high EN is due to relativistic effects (RE). The core electrons in heavy atoms such as Au are moving near the speed of light, and this alters the energies of the orbitals in the element. This is because the 1s electrons in an Au atom are circling a nucleus with a charge of +79, and so they must move very rapidly. The effect that this has is that the energies of the s electrons in the Au atom are all much lower than they would be in the absence of RE. This lowering of energies, even of the valence electrons in the 6s orbital of Au, leads to greater EN. The closer an element is to Au in the periodic table, the greater its EN.

  8. Relativistic effects: Relativistic effects arise because the inner core electrons of very heavy elements are traveling at a significant fraction of the speed of light. This increases their mass according to the familiar equation: m = mo/(1 - (v/c))1/2 (m = observed mass of electron, mo = mass of electron at rest, v is the velocity of the electron, and c is the speed of light) See: N. Koltsoyannis, JCS, Dalton Trans, 1997, 1.

  9. Ever wondered at the colors of the group 1B elements, Cu, Ag, and Au? Cu is ‘gold’ colored, then Ag is not, then Au is gold-colored. Why the discontinuity? The answer is that the color gap between the 5d and 6s levels in Au metal is lowered by RE, and so this electronic transition occurs in the visible giving Au metal its gold color.

  10. The chemistry of Au and surrounding elements, and the role of RE. The elements near Au in the periodic table all have high EN, as shown below (gold color = EN > 2.0) : Fe Co Ni Cu Zn Ga Ge EN: 1.8 1.9 1.9 1.9 1.6 1.6 1.8 Ru Rh Pd Ag Cd In Sn EN: 2.2 2.2 2.2 2.11.7 1.7 1.8 Os Ir Pt Au Hg Tl Pb EN: 2.2 2.2 2.2 2.4 2.1 2.01.9 The metals with EN > 2.0 have special chemistry where they can form stable covalent bonds to carbon, for example, and have chemistry that is much more covalent than found for other less electronegative metals.

  11. The remarkable chemistry of the metallic elements with EN > 2.0 Elements such as Pt, Ag, Au, and Hg are extremely covalent in their bonding. Thus, they form stable complexes with bonds to carbon atoms, and other elements with EN values of about 2-2.5. Examples are [Au(CN)2]- and [Hg(CN)2] (CN- = cyanide) or [Au(CH3)2]- and [Hg(CH3)2]. Hg Structure of [Hg(CH3)2]

More Related