2. Bonding Honors Chemistry Unit 6

3. Bond Types  Ionic: transfer of electrons  Covalent: sharing electron pair(s)  Metallic: delocalized electrons

4. Covalent Bonds  Characteristics  Low melting points  Don’t conduct electricity  Most are brittle if solid, but usually gas or liquid  Particle of a covalent compound is called a molecule (most are between nonmetals)

5. Covalent Bonds  Two types:  Polar covalent – one atom attracts shared pair of electrons more strongly (most) sides of bond appear to be partially charged  Nonpolar covalent – electrons are being shared equally, no charge difference (no electronegativity difference) usually between two atoms usually between two atoms of same element of same element

6. Terms to Know  Bond axis: line joining nuclei  Bond angle: angle between 2 axes  Bond length: distance between nuclei  Bond energy: energy to break bond  Bonds are not fixed  More like a stiff spring  Average position is given as bond length or bond angle

7. Molecules  Covalently bonded compounds  Diatomic molecules: always as 2 atoms when in element form (like O 2 )  7 elements, make a 7 in the periodic table (begin with N) and most are in group 17  Elements: Br, I, N, Cl, H, O, F

8. Naming a compound with two non-metals  Use the prefixes:  (1) mono-, (2) di-, (3) tri-, (4) tetra-, (5) penta, (6) hexa-, (7) hepta-, (8) octa-, (9) nona-, (10) deca  If the first element listed has a quantity of just one then you don’t use mono- as a prefix.  Put the appropriate prefix in front of the name of each element change the ending to –ide.  Example: N2O5N2O5N2O5N2O5  Dinitrogen pentaoxide

9. Lewis Dot Structures  Find the total number of valence electrons using group numbers for each element  Arrange atoms to form skeleton structure with lines connecting the atoms. If carbon is present, it is central.  Otherwise, the least electronegative element is electronegative element is central. H is NEVER central. central. H is NEVER central.

10. Lewis Structures (continued)  Each line counts as 2 electrons. Subtract these from total valence electrons.  Compare the electrons left to what each needs to be full. If they are the same, add unshared pairs to give each nonmetal or metalloid a full octet (except H). Add all electrons to see if they equal the valence electrons. (Gr2 and 13 just double their electrons don’t get a full octet, Gr. 2 gets 4, Gr. 13 gets 6)  If there are not enough electrons to give each its own dots, one more line needs to be drawn for each 2 electrons you are short (2 atoms share). Recalculate from the valence electrons and dots can be given.

11. Example 1 H2OH2OH2OH2O  2(1) + 6 = 8 valence electrons  Skeleton:....H-O-H Subtract 2 for each line 8-4 = 4 e - left Put dots to complete octet for oxygen

12. Example 2  CH 3 I  4+3(1)+7 = 14 valence electrons  Skeleton: H | H – C – H H – C – H | : I : : I :....  Subtract 2 for each line 14-8 = 6 e - left  Add dots to I to complete the octet. Other Examples: NH 3, AlI 3, SeO 2, CO 2, SO 3

13. Resonance  Using more than one Lewis structure to explain when bonds are in between drawn structures (from lab measurements)

14. Molecular Shape  Based on VSEPR theory: valence-shell electron-pair repulsion theory  Electrons want to be as far apart as possible (like charges repel)  Pairs around central atom will give angles  2 pairs: linear 180 o angle  3 pairs: trigonal planar 120 o angle  4 pairs: tetrahedral 109.5 o angle  Repulsion is greater for unshared pairs: they push harder on shared pairs, decreasing the expected bond angle 2 unshared>1 shared with one unshared>2 shared (bond angle is smaller with a lone pair

15. VSEPR  Oklahoma State Link  Possible Shapes: p. 186 of book  Linear: 2 bonded atoms 180 o angle  Trigonal planar: 3 bonded atoms 120 o angle  Tetrahedral: 4 bonded atoms 109.5 o angle  Trigonal pyramidal: 3 bonded atoms, 1 unshared (lone pair) <109.5 o angle (107 o angle)  Bent: 2 bonded atoms, 2 unshared (lone pairs) <109.5 o angle (104.5 o angle)

16. Determining Shape  Draw Lewis structure  Count bonded atoms and lone pairs on central atom (ONLY!) to determine shape  Example: H 2 O  Lewis structure:....H-O-H  2 shared (lines), 2 unshared (dot pairs)  Shape: bent, Angles: <109.5 o angle (104.5 o angle)  Other Examples: NH 3, AlI 3, CH 4, HF, SO 3

17. Hybridization  Hybridized orbitals merge s and p orbitals by borrowing empty p orbitals to put one electron in each. This allows them to share that orbital with an electron from another atom in a covalent bond.  The new hybrids have an energy that is in between that of s and p  Examples: Be, Al & B, C & Si (& others)

18. Hybrid Orbitals Hybrid Orbitals Count bonds to see how many orbitals are sp hybrids orbitals are sp hybrids needed. Start with s, then needed. Start with s, then add p orbitals to make enough. add p orbitals to make enough. This names the hybrids. This names the hybrids. sp 2 hybrids sp 3 hybrids

19. Predicting Bonds  Based on electronegativity difference  Examples of calculation: (use table on p. 151)  H-F 4.0 – 2.1 = 1.9  H-Br 2.8 – 2.1 =.7  H-I 2.5 – 2.1 =.4  The greater the difference, the stronger the bond

20. Bond Character  Large difference: ionic bond  Small difference: covalent bond  Dividing line is 1.7 > 1.7 is ionic, 1.7 is ionic, < 1.7 is covalent, = 1.7 is 50% ionic and 50% covalent  Unless bonded to the same type atom, the bond has both ionic and covalent character (use chart on back of per. table)

21.  Find the electronegativity difference in black in the chart  The percent ionic character is underneath in red  Subtract from 100 to find the covalent character  Example: H-F difference was 1.9  Ionic character is listed as 59%, so covalent character is 100-59 = 41% covalent  H-Br, H-I

22. Bonding Demo  Record color and intensity (brightness) as each bond forms in your journal, then calculate the % character for each bond  S-O  3.5 – 2.5 = 1.0 difference  22% ionic, 78% covalent  Mg-O  3.5 – 1.2 = 2.3 difference  74% ionic,26% covalent

23. Polarity  If charge of polar bonds is distributed equally in all directions, the molecule is nonpolar nonpolar  If charge of polar bonds is not equal in all directions, the molecule is polar  Look for something that makes the charge asymmetrical (either of these makes it polar)  Bonded atoms are not all the same element attached to the central atom  Unshared pairs of electrons on the central atom  A polar molecule is called a dipole (has + and – poles)  Polarity is measured as dipole moment

24. van der Waals Forces  Intermolecular: Weak forces between molecules (van der Waals forces)  Intramolecular: strong forces inside a molecule holding atoms together (bonds)  Types of van der Waals forces  Dipole-dipole: between polar molecules  Dipole-induced dipole: between dipole and nonpolar (peer pressure model)  London Dispersion Forces: temporary dipoles that happen because of electron movement  Induced by concentrations of electrons in nonpolar molecules  Only attractive force operating in nonpolar substances  85% of force in most polar molecules (exceptions: NH 3, H 2 O)

25. Induced Dipole  Peer pressure model Electrons of nonpolar molecule are disturbed by presence of charged particle (ion or dipole)

26. Dipole-Induced Dipole

27. Temporary Dipole  Movement of electrons may cause electron distribution to become asymmetrical for an instant

28. Effects of IM Forces  Properties are affected by IM forces  Boiling and melting points give an indication of how strong the IM forces are  Nonpolar substances have the weakest IM forces: gases or lowboiling liquids (lower melting and boiling points)  Polar substances have dipole forces that are stronger: liquid or solid at room temp (higher melting and boiling points)

29. Soaps and Detergents  There are polar and nonpolar sides to a soap molecule  The nonpolar side embeds or dissolves in greasy dirt  The polar side is attracted to water molecules (polar)  Agitation breaks globule up into small pieces which are then pulled away into the water and washed away.  Detergents have an additive to keep soap scum from forming.

30. Chromatography  Fractionation (separation) based on polarity  Two phases:  Mobile phase: mixture to be separated dissolved in liquid or gas  Stationary phase: solid or liquid adhering to a solid  Types: column, paper, gas

31. Column Chromatography  Stationary phase is in a column.  Used for delicate separations such as vitamins, hormones, and proteins.  HPLC and ion are special kinds of column chromatography

32. Paper Chromatography  Separation on paper into spots or lines on the strip  Has limitations

33. Gas Chromatography  Used to analyze volatile liquids and gas or vapor mixtures.  Mixed with inert gas (like He) in mobile phase  Interpreted by computer

34. Gas Chromatogram

35. Chromatography Applications  Drug testing uses column and gas chromatography  Car emissions are done with gas chromatography