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Ch 9-section-1


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Ch 9-section-1

  1. 1. Chapter 9<br />Covalent Bonding<br />
  2. 2. Review….<br />What is a chemical bond?<br />Force that holds two atoms together<br />What is an ionic bond?<br />An electrostatic force that holds oppositely charged particles together in an ionic compound<br />Compounds formed from metal & nonmetal<br />Forms when….?<br />What are atoms always trying to achieve?<br />Stability<br />Complete set of valence electrons… OCTECT<br />
  3. 3. What is a covalent bond?<br />Chemical bond that results from sharingof valence electrons<br />Occurs b/w nonmetal & a nonmetal<br />Balance b/w attractive and repulsive forces<br />2 Hydrogen Atoms<br />Sharing their 1 Ve-<br />
  4. 4. Molecules<br />Compound made when 2 or more atoms are bonded covalently<br />Diatomic molecules<br />In nature, sometimes two atoms of the same element are more stable when they are covalently bonded than the individual atom alone…<br />BrINClHOF (pronounced “Brinkle Hoff”)<br />Br2 I2 N2 Cl2 H2 O2 F2<br />
  5. 5. H<br />Cl<br />Unshared or<br />Lone pair (LP)<br />shared or Bond pair<br />Single Covalent Bonds<br />A single covalent bond –Atom shares 1 pair (2) electrons.<br />Shared pairs – both elements count the electron pair to achieve octet<br />Lonepairs– pair of electrons that are not shared b/w the atoms<br />Lewis structures- Use electron dot diagram to show how atoms are arranged in a molecule.<br />. .<br />. .<br />. .<br />
  6. 6. In the Fluorine Molecule…..<br />How many bonding pairs are there in each?<br />1<br />How many lone pairs are there each?<br />3<br />
  7. 7. Multiple Covalent Bonds<br />Double covalent bonds share two pairs of electrons. <br />CO2 O=C=O<br />Triple covalent bonds share three pairs of electrons.<br />N2 :N=N:<br />
  8. 8. Covalent Bond Formation in Hydrogen<br /><ul><li>Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron-electron repulsion.
  9. 9. However, if atoms get too close, the internuclear repulsion greatly raises the energy.</li></li></ul><li>The attractive and repulsive forces in covalent bonding must be balanced.<br />
  10. 10. Bond Length - In general, the closer the electrons are held by the atoms, the shorter the bond length and the higher the bond energy.<br />Multiple bonds result in stronger, shorter bonds.<br />
  11. 11. Bond Energy - The amount of energy required to break a bond. The greater the energy, the stronger the bond.<br />Bond breaking is an endothermic process, so bond breaking enthalpies are positive.<br />
  12. 12. Comparing Bond Length and Bond Strength<br />Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength:<br />(a) S F, S Br, S Cl<br />(b) C = O, C O, C O <br />
  13. 13. Sigma () Bonds<br />Sigma bonds are characterized by<br />Head-to-head overlap.<br />Cylindrical symmetry of electron density about the internuclear axis.<br />
  14. 14. Pi () Bonds<br />Pi bonds are characterized by<br />Side-to-side overlap.<br />Electron density above and below the internuclear axis.<br />
  15. 15. Single Bonds vs. Multiple bonds<br />Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering.<br />In a multiple bond one of the bonds is a  bond and the rest are  bonds.<br />
  16. 16. Orbital overlap<br />
  17. 17. Lewis Dot Structures<br />Determine the number of Valence e- for all atoms in the molecule<br />Divide the Ve- by 2 to get pairs (2 dots or 1 line)<br />Decide on central atom (least electronegative or furthest to the left).<br />Hydrogen & halogens are terminal atoms<br />Connect all atoms to the central atom by a bonding pair (single line)<br />Place remaining pairs around all atoms before moving on to central atom.<br />Check for octet (not H)<br />If atom does not have an octet, move lone pairs from a terminal atom to create a double or a triple bond (except grp 7).<br />