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Covalent Bonding

Covalent Bonding. -- atoms share e –. +. +. -- covalent (molecular) compounds tend to be solids with low melting points, or liquids or gases . (–) charge density. -- one shared pair = a single covalent bond of e – (i.e., 2 e – ) .

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Covalent Bonding

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  1. Covalent Bonding -- atoms share e– + + -- covalent (molecular) compounds tend to be solids with low melting points, or liquids or gases (–) charge density -- one shared pair = a single covalent bond of e– (i.e., 2 e–) -- two shared pairs = a double covalent bond of e– (i.e., 4 e–) -- three shared pairs = a triple covalent bond of e– (i.e., 6 e–)

  2. bond polarity: describes the sharing of e– between atoms nonpolar covalent bond: e– shared equally polar covalent bond: e– NOT shared equally electronegativity (EN): the ability of an atom in a molecule to attract e– to itself -- A bonded atom w/a large EN has a great ability to attract e–. max. EN = 4.0 (F) -- A bonded atom w/a small EN does not attract e– very well. min. EN = 0.7 (Cs) (see p. 308) -- EN values have been tabulated.

  3. DEN < 0.5 0.5 < DEN < 2.0 DEN > 2.0 The DEN between bonded atoms approximates the type of bond between them. nonpolar covalent polar covalent ionic bond As DEN increases, bond polarity... increases. DEN for a C–H bond = 0.4, so all C–H bonds (such as those in candle wax) are nonpolar.

  4. H H partial charge O Dipole Moments Polar covalent molecules have a partial (–) and a partial (+) charge and are said to have a dipole moment. d+ d– d+ d+ H–F H–F d– big big big DEN = _____ polarity = _____ dipole moment

  5. Polar molecules tend to align themselves with each other and with ions. NH4+ H–F NO3– H–F H–F H–F NH4+ NO3– H–F H–F ** Nomenclature tip: For binary compounds, the less electronegative element comes first. -- Compounds of metals w/high ox. #’s (e.g., 4+ or higher) tend to be molecular rather than ionic. e.g., TiO2, ZrCl4, Mn2O7

  6. Lewis Structures (or “electron-dot structures”) 1. Sum the valence e– for all atoms. If the species is an ion, add one e– for every (–); subtract one e– for every (+). 2. Write the element symbols and connect the symbols with single bonds. 3. Complete octets for the atoms on the exterior of the structure, but NOT for H. 4. Count up the valence e– on your L.S. and compare that to the # from Step 1. -- If your LS doesn’t have enough e–, place as many e– as needed on central atom. -- If LS has too many e– OR if central atom doesn’t have an octet, use multiple bonds.

  7. Cl O H–C–N O Draw Lewis structures for the following species. .. .. .. .. .. Cl– P –Cl .. .. PCl3 26 e– .. .. .. .. .. HCN 10 e– H–C–N .. .. H–C=N .. .. .. [ ] .. .. 3– .. .. 32 e– PO43– .. .. O– P –O .. .. .. .. ..

  8. H H H–C–C–O–H H H H2 2 e– H–H .. 20 e– CH3CH2OH .. This CO bond is longer (and weaker) than this CO bond. .. .. CO2 16 e– .. .. .. .. O=C=O .. O–C–O .. .. .. As the # of bonds between two atoms increases, the distance between the atoms... decreases.

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