2. Warm Up 3/3/09 How do elements bond? What is does the word covalent mean?

3. Two types of bonds Ionic and Covalent. Ionic bonds are electrical attraction (like magnets) between a cation (+) and anion (-). ex. Na + and Cl - form an Ionic bond. NaCl Usually this is a metal and a nonmetal. Electrons are not shared. The metals give away their electron. The nonmetal accepts the electrons.

4. Two types of bonds Ionic and Covalent. Covalent bonds occur when valence electrons are shared. Ex. Nitrogen wants to gain 3e -. Iodine wants to gain 1e -. NI 3 Covalent bonds occur between nonmetals. There are polar covalent and nonpolar covalent bonds.

5. A compound is made of multiple atoms. Both ionic compounds and molecular. A molecule is a compound made up of covalent bonds. An ionic compound is made of ionic bonds.

6. Chapter 6

7. Classify the following compounds as molecular or ionic. MgBr 2 PCl 3 NH 3 FeO IonicMolecular Molecular Ionic Ionic bonds Covalent Bonds Ionic Bond

8. Remember elements want to be like the Noble Gases. They strive to have a full outer valence shell of 8 electrons. Lets look at F 2 and HCl.

9. The Octet Rule Noble gas atoms are unreactive because their electron configurations are especially stable. This stability results from the fact that the noble-gas atoms’ outer s and p orbitals are completely filled by a total of eight electrons. Other atoms can fill their outermost s and p orbitals by sharing electrons through covalent bonding. Such bond formation follows the octet rule: Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

10. The Octet Rule, continued Exceptions to the Octet Rule Exceptions to the octet rule include those for atoms that cannot fit eight electrons, and for those that can fit more than eight electrons, into their outermost orbital. Hydrogen forms bonds in which it is surrounded by only two electrons. Boron has just three valence electrons, so it tends to form bonds in which it is surrounded by six electrons. Main-group elements in Periods 3 and up can form bonds with expanded valence, involving more than eight electrons. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

11. BC SKr

12. MgAl HeP

13. When we draw structures we need to draw an estimated structure of the atom. 1 st Determine the elements involved and total valence electrons. ex. For Carbon Tetrachloride CCl 4 4 x Cl = 4 x 7 valence e - = 28e - 1 x C = 1 x 4 valence e - = 4e - 32e -

14. CCl 4 Next we draw a skeleton of the atom. The first element in the chemical (molecular) formula is in the middle. And give all the elements their octet. (Except H and B) Then count to make sure all electrons have been accounted for. C Cl

15. Try H 2 S

16. Double bonds exist when atoms share many electrons. We can drawn these structures as well. Try O 2

17. Write the electron dot structure for the following: 1) Arsenic 3) Aluminum Draw the Lewis structure for the following PF 3 O 3 (PO 4 ) -3

18. Resonance - when there is more than one proper way to draw a Lewis structure. :O = O – O:  :O – O = O:.. The Double Bond Switches Places. These are resonance structures.

19. These are ions made up of many atoms. Poly means many. Many atoms (NH 4 ) + (SO 4 ) -2 (CrO 4 ) -2

20. Ionic Compounds Most of the rocks and minerals that make up Earth’s crust consist of positive and negative ions held together by ionic bonding. example: table salt, NaCl, consists of sodium and chloride ions combined in a one-to-one ratio— Na + Cl – —so that each positive charge is balanced by a negative charge. An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. Section 3 Ionic Bonding and Ionic Compounds Chapter 6

21. Ionic Compounds, continued The chemical formula of an ionic compound represents not molecules, but the simplest ratio of the compound’s ions. A formula unit is the simplest collection of atoms from which an ionic compound’s formula can be established. Section 3 Ionic Bonding and Ionic Compounds Chapter 6

22. NaCl and CsCl Crystal Lattices Section 3 Ionic Bonding and Ionic Compounds Chapter 6

23. Ionic Vs. Covalent Bonding Section 3 Ionic Bonding and Ionic Compounds Chapter 6

24. Formation of Ionic Compounds The sodium atom has one valence electrons and the chlorine atom has seven valence electrons. Atoms of sodium and other alkali metals easily lose one electron to form cations. Atoms of chlorine and other halogens easily gain one electron to form anions. Section 3 Ionic Bonding and Ionic Compounds Chapter 6

25. Some examples of Lewis structures of polyatomic ions are shown below. Section 3 Ionic Bonding and Ionic Compounds Chapter 6 Polyatomic Ions, continued

26. VSEPR Theory The abbreviation VSEPR (say it “VES-pur”) stands for “valence-shell electron-pair repulsion.” VSEPR theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible. example: CO 2 The central carbon atom is surrounded by only the two electron pairs it shares with the oxygen atoms. According to VSEPR, the shared pairs will be as far away from each other as possible, so the bonds to fluorine will be 180° apart from each other. The molecule will therefore be linear: Section 5 Molecular Geometry Chapter 6

27. VSEPR Theory, continued Sample Problem E Use VSEPR theory to predict the molecular geometry of boron trichloride, BCl 3. Section 5 Molecular Geometry Chapter 6

28. VSEPR Theory, continued Sample Problem E Solution, continued Section 5 Molecular Geometry Chapter 6 Notice the electron pairs spacing. As far apart as possible. Its geometry should therefore be trigonal-planar.

29. VSEPR theory can also account for the geometries of molecules with unshared electron pairs. examples: ammonia, NH 3, and water, H 2 O. The Lewis structure of ammonia shows that the central nitrogen atom has an unshared electron pair: Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued VSEPR theory postulates that the lone pair occupies space around the nitrogen atom just as the bonding pairs do.

30. Taking into account its unshared electron pair, NH 3 takes a tetrahedral shape. The shape of a molecule refers to the positions of atoms only. The geometry of an ammonia molecule is that of a pyramid with a triangular base. H 2 O has two unshared pairs, its geometry takes the shape of a “bent,” or angular, molecule. Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued

31. Unshared electron pairs repel other electron pairs more strongly than bonding pairs do. This is why the bond angles in ammonia and water are somewhat less than the 109.5° bond angles of a perfectly tetrahedral molecule. Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued

32. Expanded Octet Draw the following. SF 6 BrF 5

33. 1-16Shape Name# of bonding e - pairs # of Lone e - pairs Bond angles Orbital Hybridization Example of the Molecule 1 2 3 4 5 6 7 8 9 10 Create a table (set up horizontally)

34. Draw these structures. The classify them as on of the “molecule types” in your chart. Show work for partial credit, show resonance when necessary. 1. CO 2 2. PCl 5 3. SF 4 4. CH 4 5. NH 3 6. O 2 7. BF 3 8. XeF 4 9. (OH) - 10. SF 6 11. XeF 2 12. BrF 5 13. ClF 3 14. SO 2 15. N 2 16. H 2 O

35. Warm Up 1) Draw the electron dot structure for carbon. 2) Draw the Lewis Structure for CH 4 3) Draw the orbital diagram for the valence electrons of Hydrogen and Carbon. Challenge Question. Draw the structure for XeF 2 what is the geometry for this structure?

36. Hybridization VSEPR theory is useful for predicting and explaining the shapes of molecules. A step further must be taken to explain how the orbitals of an atom are rearranged when the atom forms covalent bonds. For this purpose, we use the model of hybridization, which is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies. Section 5 Molecular Geometry Chapter 6

37. Polar Covalent Bond – Electrons are not shared evenly in the bond. One atom is more electronegative than the other. H ---- F +----  Nonpolar Covalent Bond – Electrons are shared evenly. O ---- O Both are equally hungry for electrons. Nonpolar Electrons are pulled closer to the Fluorine. 

38. Covalent Bonds Classify the following as polar or nonpolar. Indicate the partial negative and positive charge if necessary. 1) H – Cl 2) C – F 3) Br – Br 4) Cl – I

39. To see why compounds love or fear water we need to look at the structure. HH - - O H H + + Water is a polar molecule. It has positively charged pole and a negatively charged pole. The charges are ASYMMETRIC! Remember… Hydrogen has positive charge because Oxygen is more Electronegative. Electronegative. The more electronegative the atom, the more hungry the atom is for electrons and will fight to keep electrons closer to itself. The negative charge on Oxygen are due to nonbonding electron pairs.

40. Non-polar molecules have an over all structure that is SYMMETRIC! Thus, Charges cancel each other out.

41. Classify the following molecules as polar or non-polar.. AmmoniaHydrochloric AcidMethane + + + + + + + + - - -

42. Intermolecular Forces The forces of attraction between molecules are known as intermolecular forces. The boiling point of a liquid is a good measure of the intermolecular forces between its molecules: the higher the boiling point, the stronger the forces between the molecules. Intermolecular forces vary in strength but are generally weaker than bonds between atoms within molecules, ions in ionic compounds, or metal atoms in solid metals. Boiling points for ionic compounds and metals tend to be much higher than those for molecular substances: forces between molecules are weaker than those between metal atoms or ions. Section 5 Molecular Geometry Chapter 6

43. Comparing Ionic and Molecular Substances Section 5 Molecular Geometry Chapter 6

44. Intermolecular Forces, continued A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole. The dipole created by a hydrogen chloride molecule is indicated as follows: Section 5 Molecular Geometry Chapter 6

45. Intermolecular Forces, continued The negative region in one polar molecule attracts the positive region in adjacent molecules. So the molecules all attract each other from opposite sides. Such forces of attraction between polar molecules are known as dipole-dipole forces. Dipole-dipole forces act at short range, only between nearby molecules. Dipole-dipole forces explain, for example the difference between the boiling points of iodine chloride, I–Cl (97°C), and bromine, Br–Br (59°C). Section 5 Molecular Geometry Chapter 6

46. Comparing Dipole-Dipole Forces Section 5 Molecular Geometry Chapter 6

47. Intermolecular Forces, continued A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons. The result is a short-range intermolecular force that is somewhat weaker than the dipole-dipole force. Section 5 Molecular Geometry Chapter 6

48. Intermolecular Forces, continued Some hydrogen-containing compounds have unusually high boiling points. This is explained by a particularly strong type of dipole-dipole force. In compounds containing H–F, H–O, or H–N bonds, the large electronegativity differences between hydrogen atoms and the atoms they are bonded to make their bonds highly polar. This gives the hydrogen atom a positive charge that is almost half as large as that of a bare proton. Section 5 Molecular Geometry Chapter 6

49. Intermolecular Forces, continued The small size of the hydrogen atom allows the atom to come very close to an unshared pair of electrons in an adjacent molecule. The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule is known as hydrogen bonding. Section 5 Molecular Geometry Chapter 6

50. Intermolecular Forces Hydrogen bonds are usually represented by dotted lines connecting the hydrogen-bonded hydrogen to the unshared electron pair of the electronegative atom to which it is attracted. An excellent example of hydrogen bonding is that which occurs between water molecules. The strong hydrogen bonding between water molecules accounts for many of water’s characteristic properties. Section 5 Molecular Geometry Chapter 6

51. Visual Concepts Hydrogen Bonding Section 5 Molecular Geometry Chapter 6

52. Intermolecular Forces, continued London Dispersion Forces Even noble gas atoms and nonpolar molecules can experience weak intermolecular attraction. In any atom or molecule—polar or nonpolar—the electrons are in continuous motion. As a result, at any instant the electron distribution may be uneven. A momentary uneven charge can create a positive pole at one end of an atom of molecule and a negative pole at the other. Section 5 Molecular Geometry Chapter 6

53. Intermolecular Forces, continued London Dispersion Forces, continued This temporary dipole can then induce a dipole in an adjacent atom or molecule. The two are held together for an instant by the weak attraction between temporary dipoles. The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles are called London dispersion forces. Fritz London first proposed their existence in 1930. Section 5 Molecular Geometry Chapter 6