1. Review – Packet #7  Bonds can be classified as being either polar or non-polar.  Polarity: tendency of a molecule, or compound, to be attracted or repelled by electrical charges because of an asymmetrical arrangement of atoms around the nucleus.  Think of it like a game of tug of war, if one end of the compound is pulling on the electrons more than the other, there is an unequal pull, and therefore, the substance is polar. If there is an equal pull, then the substance is non-polar.  This concept of polarity is determined by electronegativity.

2. Ionic Bond  Attraction between oppositely charged ions  Occurs when electrons are transferred from one ion (charged particle) to another  Electronegativity difference 1.7+  Metals react with Nonmetals to form ionic compounds  Always Polar !!!

3. Lewis Dot Structure of Ionic Compounds  KCl  CaBr 2  KNO 3  (NH 4 ) 3 PO 4

4. Possible Combinations for Ionic Compounds Formula (+)( - )Examples MNMNaCl, KI, and CaF 2 MPLiNO 3 and Sr(CN) 2 PNMNH 4 Br, Hg 2 S, and (H 3 O) 2 Se PPNH 4 NO 3 and H 3 OMnO 4

5. Properties of Ionic Compounds  Hard  Good conductors of electricity in liquid or aqueous form only, because ions can move in solution and in liquid form, but not in solid form.  High melting and boiling points  Solid at room temperature  Dissolve in polar substances: like water. (Polar – opposite charges).

6. Covalent Bonds  Formed when 2 atoms (both nonmetals) share electrons. [Example Cl 2 or H 2 O]  Neither atom pulls strongly enough to remove an electron from the other  The EN difference is < 1.7  Unpaired electrons pair up in such a way that the atoms complete their outer shells  Covalent compounds also referred to as molecular compounds

7. Properties of Covalent Bonds  Gases, liquids or solids  Soft  Nonmetals  Poor conductors of heat and electricity because they are not charged particles. (No ions or mobile electrons)  Low melting and boiling points because of weak attraction between molecules.

8. Polar vs. Non-Polar Covalent Bonds  Unlike an ionic compound, a covalent compound can be classified as either a polar covalent bond, or a non-polar covalent bond.  If the EN of the atoms are different then it is a polar covalent bond.  If the EN of the atoms are the same or very similar then it is a non-polar covalent bond. 0.0 - 0.4 = non-polar covalent 0.5 - 1.6 = polar covalent

9. Number of Covalent Bonds  Single covalent bond: one pair of shared electrons; 2 electrons total  -Double covalent bond: two pairs of shared electrons; 4 electrons total  -Triple covalent bonds: three pairs of shared electrons; 6 electrons total

10. Must know how to draw these!! HFH 2 OCO 2 CH 4 NH 3 HClH 2 SCS 2 CF 4 PH 3 HBrSiO 2 CCl 4 HISiS 2 CBr 4 CI 4

11. Partially Positive & Negative  In a polar covalent bond, both of the elements are non-metals, and therefore there is no “true” + or – charges; instead there are partially (+) and partially (-) charges.  The element with the higher EN is partially (-) and the one with the lower EN is partially (+)

12. This is a SNAP! S ymmetric are N onpolar A symmetric are P olar

13. Other Types of Covalent Bonds Coordinate Covalent Bond:  When one atom donates both of the electrons that are shared  Example: NH 4 + and H 3 O +  Nitrogen donates a pair of electrons to share with H + forming a coordinate covalent bond between nitrogen and hydrogen

14. Other Types of Covalent Bonds Network Solids:  Solids that have covalent bonds between atoms linked in one big network or one big macromolecule with no discrete particles. This gives them some different properties from most covalent compounds.  They are hard, poor conductors of heat and electricity, and have high melting points  Examples include: Diamond (C), silicon carbide (SiC), and silicon dioxide (SiO 2 )

15. Metallic Bond  Occurs only in metals (Example Copper)  Metals have low ionization energy meaning they hold onto their valence electrons very loosely  As a result the electrons in metallic substances move about very easily and are not associated with any particular atom  Therefore, the particles of a metal are usually positive ions surrounded by a mobile sea of electrons  The attraction between the positive cations and the moving electrons is what holds the metal together  Properties of Metallic Bonds are that of metals: hard, good conductors of heat & electricity, malleable, ductile, etc...

16. Dipole-Dipole Attractions Positive end of a polar molecule is attracted to the negative end of an adjacent polar molecule.

17. Hydrogen Bonding An intermolecular attraction between a hydrogen atom in one molecule to a nitrogen, oxygen, or fluorine atom in another molecule The strongest intermolecular force Substances with hydrogen bonds tend to have much higher melting and boiling points than those without hydrogen bonds Example: The boiling point of H 2 O is much higher than H 2 S

18. London Dispersion Forces AKA: van der Waals Forces  Weak intermolecular forces between non-polar molecules (like diatomic molecules)  Dispersion forces make it possible for small, non- polar molecules to exist in both liquid or solid phases under conditions of high or low temperatures.  Increases with molecular size, Ex. As you go down group 17, dispersion forces increase and boiling point increases.

19. Molecule-Ion Attraction  Attraction between the ions of an ionic compound such as NaCl, and a molecule such as water (or any other polar covalent compound).  When you put NaCl into water, the Na + from the salt is attracted to the O from the water which is partially (-), and the Cl - from the salt is attracted to the H + of the water.

20. THE VSEPR Model 180° 120°

21. THE VSEPR Model 104.5° 107° 109.5 °

22. THE VSEPR Model 120° 90°

23. THE VSEPR Model 90°

24. © 2009, Prentice-Hall, Inc. sp Hybrid Orbitals Consider BeF 2 –In beryllium’s ground electronic state, it would not be able to form bonds with fluorine because it has no singly-occupied orbitals –Linear structure

25. © 2009, Prentice-Hall, Inc. sp 2 Hybrid Orbitals Using a similar model lets consider BF 3 Boron (1s 2 2s 2 2p 1 ) does not have three single electrons to share with the three F atoms An electron is promoted to the 2p orbital s,p,p (sp 2 ) Trigonal Planar

26. © 2009, Prentice-Hall, Inc. sp 3 Hybrid Orbitals Using a similar model lets consider CH 4 Carbon (1s 2 2s 2 2p 2 ) does not have four single electrons to share with the four hydrogen atoms An electron is promoted to the 2p orbital s,p,p,p (sp 3 ) Tetrahedral

27. Copyright © Cengage Learning. All rights reserved 27 sp 3 d Hybrid Orbitals Combination of one s orbital, three p orbitals, and one d orbital Gives a trigonal bipyramidal arrangement of hybrid orbitals. Examples include: PF 5, SF 4, BrF 3

28. sp 3 d 2 Hybrid Orbitals  Combination of one s orbital, three p orbitals, and two d orbital  Gives a octahedral arrangement of hybrid orbitals.  Examples include: SF 6, ClF 5, XeF 4, PF 6 -

29. Hybrid Orbitals The type of hybridization can be determined based on the electron-domain geometry of a particular molecule.

30. Hybrid Orbitals