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Covalent Bonds

Covalent Bonds. Both atoms involved (typically nonmetal) “want” to gain e - to become stable Electrons are shared in order to allow this to happen The number of e - shared depends on the element When atoms share two, or more, electron(s), they form a molecule.

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Covalent Bonds

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  1. Covalent Bonds • Both atoms involved (typically nonmetal) “want” to gain e- to become stable • Electrons are shared in order to allow this to happen • The number of e- shared depends on the element • When atoms share two, or more, electron(s), they form a molecule

  2. Compounds formed with covalent bonds are neutral and, generally, follow the OCTET rule! • All atoms are most stable when they have the electron configuration of a noble gases (THIS MEANS HAVING EIGHT* VALENCE ELECTRONS) • For nonmetals the most common exception to the octet rule (when forming bonds) is HYDROGEN.

  3. Highly Stable = Noble gas configuration • Hydrogen - will share 1 e- • Oxygen - will share 2 e- • Nitrogen - will share 3 e- • Chlorine - will share 1 e- • Carbon - will share 4 e-

  4. Covalent Bonding in Hydrogen Electron sharing can occur only when electron orbitals from two different atoms overlap.

  5. FG08_006.JPG Formation of a Covalent Bond

  6. The number of covalent bonds formed by a nonmetallic element is often directly related with the number of electrons it must share (commonly equivalent to the number of lone pairs of e- the atom has) in order to obtain an octet of electrons.

  7. Types of Covalent Bonds • Single - ONE pair of e- shared • Double - TWO pair of e- shared • Triple - THREE pair of e- shared • Coordinate - Both electrons being shared originate from a single atom

  8. (a) A “regular” covalent single bond is the result of overlap of two half-filled orbitals. (b) A coordinate covalent single bond is the result of overlap of a filled and a vacant orbital. - atoms participating in cc bonding generally do not form their normal # of covalent bonds Ex.: HO2Cl, CO

  9. Types of Covalent Bonds

  10. FG08_009.JPG Electron Dot Structuresof HCl

  11. Lewis Structures • Structures which represent in a drawing the arrangement of the atoms and the types of covalent bonds • There are FIVE basic steps to follow. • We’ll use water as an example:).

  12. Step 1: • Arrange the atoms! • Remember, Hydrogens always on the periphery • Use the expected bonding patterns to arrange the atoms • the atom that forms the most bonds is typically in the middle of the structure.

  13. Step 2: • Count the electrons • Obtain a total to work from

  14. Step 3: • Add the bonds & lone pairs • Remember to give each atom (except H) an “octet” of e- (but, don’t exceed the number of e- available)

  15. Step 4: • Use multiple bonds to fill octets when needed • Convert one lone pair to a bonding pair for each pair of e- needed to complete an octet

  16. Step 5: • Exceptions to the octet rule • Less than an octet • Hydrogen & Boron • More than an octet • Phosphorus & Sulfur, Noble gases

  17. CCl4 PBr3 F2 H2S NH4+ SO4-2 O3 Lewis Structures - examples

  18. Double Bonds • O2 • C2H4 • CO2 • CH2CHCHCH2 • NO3- • CO3-2

  19. Triple bonds • N2 • C2H2 • HCN

  20. Exceptions (more than an octet) - Only elements in rows 3 and beyond Why? • PCl3 and NCl3 • PCl5 but not NCl5 • XeF4 • SF6

  21. Resonance • Some molecules have measured values of bond lengths which do not support the Lewis structure drawn for the molecule • Example: Ozone, O3 • To adequately represent such molecules with Lewis structures, you should draw all possible arrangements of ELECTRONS.

  22. Used only with the second element Ex. Dihydrogen monoxide Ex. Tetraphosphorus decoxide Naming Binary Molecular Compounds

  23. Common vs. Chemical Names • Chemical Name

  24. 3-D arrangements of electron pairs Arrangement of valence electron pairs about a central atom that minimize repulsions between the pairs. Since double & triple bonds are multiple electron pairs in the same location, they act like a single pair when determining the geometry of the molecule

  25. 3-D Models of Molecules (a) Acetylene molecule. (b) Hydrogen peroxide molecule. (c) Hydrogen azide molecule.

  26. Bonding

  27. Molecular Geometry - VSEPR

  28. Bond Polarity & Electronegativity • The difference in the electronegativities (ability of an atom to attract electrons) of the two bonded atoms can be used to define the “polarity” of the bond.

  29. Abbreviated periodic table showing Pauling electronegativity values for selected representative elements.

  30. Bond Polarity (a) In the nonpolar covalent bond present, there is a symmetrical distribution of electron density. (b) In the polar covalent bond present, electron density is displaced because of its electronegativity.

  31. Bond Lengths • Bond lengths are measured using nucleus-nucleus distances. • For bonds between the same two atoms: • Single > Double > Triple • Example: C-O

  32. Molecular Polarity (a) Methane is a nonpolar tetrahedral molecule. (b) Methyl chloride is a polar tetrahedral molecule. Hint: Is the molecule SYMMETRICAL?

  33. Geometry Polarity Look @ center atom All bonds non-polar = NON-POLAR molecule Bonds are Polar Only VSEPR bonding groups - go to B. Has VSEPR nonbonding groups = POLAR molecule Look @ attached atoms All attached atoms the same = NON-POLAR molecule One or more different element’s atoms attached = POLAR molecule Determining Molecular Geometry & Polarity- a Shortcut!

  34. Determine the geometry & polarity (look for symmetry) of these “molecules”: HF H2O SF2 NI3 SiBr4 SeO3 CO2 CO3-2 Na2SO4 Examples- use Lewis structures to guide you

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