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Covalent Bonding -- atoms share e – -- covalent (molecular) compounds tend to be solids with low melting points, or liquids or gases -- three shared pairs.

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Presentation on theme: "Covalent Bonding -- atoms share e – -- covalent (molecular) compounds tend to be solids with low melting points, or liquids or gases -- three shared pairs."— Presentation transcript:

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

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

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

4 “Oh, man… I forgot what the most electronegative elements are.” “Shee-oot… Ow teh ye… for a nickel.” FO’ NCl.” F = 4.0 O = 3.5 N = Cl = 3.0 C = 2.5 H = 2.0

5 Dipole Moments Polar covalent molecules have a partial (–) and a partial (+) charge and are said to have a dipole moment. H–F  +  – partial charge O HH –– ++ ++ big  EN = _____ polarity = _____ dipole moment bigbig

6 Polar molecules tend to align themselves with each other and with ions. H–F NO 3 – NH 4 + ** 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., TiO 2, ZrCl 4, Mn 2 O 7

7 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. -- 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. 4. Count up the valence e – on your L.S. and compare that to the # from Step 1.

8 26 e – Draw Lewis structures for the following species. PCl 3 P–ClCl– Cl.. HCN 10 e – H–C–N H–C=N H–C–N.. [ ] PO 4 3– P–OO–.. O O 32 e – 3–

9 H2H2 CH 3 CH 2 OH CO 2 As the # of bonds between two atoms increases, the distance between the atoms... 2 e – H–H 20 e – H H–C–C–O–H H HH 16 e – O=C=O O–C–O.. decreases. This CO bond is longer (and weaker) than this CO bond.


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