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8.2

8.2. The Nature of covalent bonding. Polyatomic ions. POLY = many ATOMIC  atoms So POLYATOMIC IONS are IONS made of 2 or more atoms. MORE TYPES OF COVALENT BONDS. We already saw SINGLE , DOUBLE , and TRIPLE covalent bonds

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8.2

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  1. 8.2 The Nature of covalent bonding

  2. Polyatomic ions • POLY = many • ATOMIC  atoms • So POLYATOMIC IONS are IONS made of 2 or more atoms

  3. MORE TYPES OF COVALENT BONDS • We already saw SINGLE, DOUBLE, and TRIPLE covalent bonds • SINGLE, DOUBLE, and TRIPLE refer to the number of shared electrons pairs • COORDINATE COVALENT BONDS • Sometimes, both electrons in the pair will come from the same atom. • Example: CO (Carbon Monoxide) • These get a special notation as well • We still draw a line, but put an arrow that goes in the direction of giving; i.e. facing the atom that did not contribute the pair.

  4. BOND DISSOCIATION ENERGY • Dissociate: the opposite of associating together • So, bond dissociation energy is the energy needed to break apart a covalent bond. • Just keep this in mind for now. We will visit that concept in more detail later on

  5. RESONANCE • Sometimes multiple dot structures are valid that represent the possible covalent bonds between atoms. • A lot of times, these are simply mirror images of another. • Often, these structures will have slightly different chemical properties • Our example is OZONE

  6. OCTET RULE EXCEPTIONS • The Octet rule is broken in molecules whose total valence electrons is an odd number. • Molecules with unpaired electrons are called free radicals and are reactive • Some molecules can also have atoms with fewer or more than an octet of valence electrons • Our example is NO2 (Nitrogen Dioxide) • Others are NO (Nitric Oxide) and ClO2 (Chloride Dioxide)

  7. VseprtHEORY • σ: Sigma: Symmetric around the axis containing the two molecules. • π: Pi: Not symmetric around the axis. • For now, just know about these… • All the valence electron pairs want to stay as far from each other as possible. • Remember: Like charges repel • Linear: makes a straight line • Planar: Lies completely flat

  8. Molecular shapes w/ & w/out electron pairs • The diagrams of common molecular shapes to ONLY show the atoms in the molecule. • They do NOT illustrate any electron pairs. • So technically all these molecules have a tetrahedral shape when we look at all the things surrounding their central atoms. • But, when we JUST look at the involved atoms, they have shapes other than tetrahedral: • E.g. Methane (CH4), Ammonia (NH3), and Water (H2O)

  9. The shape not including unshared electron pairs (Column 4) is the shape we regularly refer to for these common molecules. If you notice, the number of atoms plus unshared electron pairs are all 5 for each of the molecules listed. They all have the same number of things around the central atom, but a different number of total atoms.

  10. The tetrahedron • The closest we can get four spheres is a tetrahedron • Try it yourself with 4 marbles • This explains some of the angled structures we get

  11. Polar covalent bonds • Form a pole (think north and south pole – 2 opposing ends) • The more electronegative atom attracts more strongly and gains a slightly negative charge and vice-versa. • This forms what is called a dipole (2-poles) • We represent this with a sword looking symbol pointing to the negative charge and a lower-case Greek d: δ (delta). • Water is our best example. δ- δ+

  12. NONPOLAR COVALENT BONDS • Do not form a pole… • The atoms in the bond all pull equally • The bonding electrons are shared equally • Example: Carbon Dioxide

  13. NETWORK SOLIDS • These are solids of covalent bonds. • Melting a network solid requires breaking down covalent bonds throughout the entire solid.

  14. Attractions between molecules • Before, we saw attractions between just atoms. • Now, we’ll look at forces acting in between molecules. • These 3 are very similar, but have subtle differences.

  15. Van Der Waals forcES • Dipole interactions • Occurs with POLAR molecules • Opposites attract • Different dipole charges • Dispersion forces • Occur among POLAR molecules and even NONPOLAR molecules • Like charges repel • Because electrons are always moving, sometimes there might just happen to be more electrons on one side of a molecule (and they sort of disperse from one side) at a particular moment. • That, in turn, influences the electrons on a neighboring molecule. • And again, opposites attract, so we’ll have a temporary dipole effect • Let’s look at dispersion forces among diatomic molecules from the Halogens: • This happens often enough with Bromine that it is a liquid at room temperature and pressure • This also happens with Iodine, but because it has more electrons, this effect is stronger and Iodine is a solid at room temperature and pressure.

  16. Graphite (pencil lead)

  17. Hydrogen bonds • ONLY OCCURS WITH HYDROGEN • Hydrogen MUST be locked into a covalent bond with an atom. Normally, they are very electronegative. • It is WEAKLY bonded to an UNSHARED electron pair of another atom. That atom is normally very electronegative as well. • Example: WATER • Because water molecule want to stay together, water remains a liquid and not a gas at room temperature. • These bonds are responsible for surface tension in water. • Oxygen is fairly electronegative

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