1. NATURE OF COVALENT BONDS 8.2 Cont’d

2. The difference…

3. Review  Atoms will share enough electrons so that they each have a 8 electrons in their outer shell  Noble gas configuration  Single Covalent Bond When they share a pair of electrons (each give 1 each to share)

4. Terminology Lone Pair Unshared Pair Nonbonding Pair All the same! -pair of electrons that is not shared between atoms/ is not involved in bonding

5. Examples  Double bonds  Carbon dioxide (CO 2 )

6. Use your book to define…  coordinate covalent bond  One atom is going to give 2 of its electrons to share  The other atom does not have a free pair to share and they both do not follow the octet rule (usually only one does)  So the one atom will share its lone pair, then it will still follow the octet, and the other atom will too!

7. Example  Coordinate covalent bond  Carbon Monoxide (CO) CO........ C O.. Carbon isn’t happy!

8. Carbon Monoxide  Shared pair came from one of the atoms, as opposed to one electron from each atom  Now both atoms follow the octet rule C O..

9. Polyatomic Ions  Use a combo of regular covalent bonds and coordinate covalent bonds  Example of polyatomic ion we have used before: NH 4 + (ammonium) H+H+

10. Practice  With negative polyatomic ions, extra electrons will come into play  SO 3 2- means that 2 extra electrons will be needed to make this molecule stable O O O S............ 2-2-

11. Bond Dissociation Energies  Energy needed to break a covalent bond between two atoms  Large bond dissociation energy means it is a strong bond  Single vs. Double vs. Triple bonds BondLength (pm)Energy (kJ/mol) C-C154348 C=C134614 C=C120839

12. Resonance Structures  When you can draw more than one valid Lewis structure, then you have resonance structures  All the atoms have the same number of electron pairs, bonds may just be moved around  The atoms will not move, only the electrons

13. Exceptions to the Octet Rule  There are some molecules that do not follow the octet rule but still occur in nature  If the total number of valence electrons is ODD, then not all electrons can be paired  Sometimes an atom can have less than 8 valence electrons and still create a stable compound (BF 3 )  Sometimes an atom can have more than 8 valence electrons and still create a stable compound (SF 5 )

14. POLARITY 8.4

15. Electronegativity and Polarity  Some atoms are more electronegative than others  Have a stronger pull on electrons  Fluorine is the most electronegative AtomElectronegativity Fluorine4.0 Oxygen3.5 Nitrogen3.0 Chlorine3.0 Bromine2.8 Carbon2.5 Sulfur2.5 Hydrogen2.1 Pg 177

17. Polar Covalent Bonds  The more electronegative atom will attract the electrons more strongly  The more electronegative atom will then have a slightly negative charge  The less electronegative atom will then have a slightly positive charge  When electrons are shared unequally between two atoms it is called a polar covalent bond

18. HCl....

19. Polar Covalent Bonds  The electron cloud image shows a darker/more dense cloud around the more electronegative atom  The charge on the chlorine is less than 1, so it is only a partial charge  We use the lowercase Greek letter for delta to indicate the word “partial” δ H – Clor H – Cl

20. Nonpolar Covalent Bonds  If electrons are shared equally and there is no partial charge, then it is a nonpolar bond  This happens when identical atoms are bonded together  It may also happen if the difference in electronegativities is very low

21. Water  Is water polar?  Look at the individual bonds and shape 2.1 3.5 δ+δ+ δ+δ+ δ-δ-

22. Polar Molecules  If a single polar bond is present, then the entire molecule is polar  A molecule that has 2 poles (positive and negative) is called a dipole  HCl is a dipole  These molecules would be attracted to a magnetic field

23. Nonpolar Molecules  A molecule can have 2 polar bonds but if they are exact opposites of each other they can cancel out Symmetrical, nonpolar

24. Types of Solids (network)