1. Chapter 10 Molecular Structure: Solids and Liquids Electron Configurations and Dot Formulas Review

2. Electron-Dot Formulas Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings Why look at Electron-Dot Formulas? Ionic Compounds: Helps to determine formulas Covalent Compounds: Help us to understand molecular structures or molecular geometries, and molecular properties.

3. Preview of Molecular Geometries or Shapes

4. Electron-Dot Symbols Electron-dot symbols show the valence electrons of an atom one to four valence electrons as single dots on the sides of an atomic symbol Five to eight valence electrons with one or more pairs of dots on the sides of an atomic symbol

5. Valence Electrons in Some Electron-Dot Symbols

6. Guide to Writing Electron-Dot Formulas STEP 1 Determine the arrangement of atoms. STEP 2 Add the valence electrons from all the atoms. STEP 3 Attach the central atom to each bonded atom using one pair of electrons. STEP 4 Add remaining electrons as lone pairs to complete octets (2 for H atoms). STEP 5 If octets are not complete, form one or more multiple bonds.

7. Multiple Covalent Bonds In a single bond One pair of electrons is shared. In a double bond, Two pairs of electrons are shared. In a triple bond. Three pairs of electrons are shared.

8. Electron-Dot Formulas Electron-dot formulas show The order of bonded atoms in a covalent compound. The bonding pairs of electrons between atoms. The unshared (lone, non- bonding) valence electrons. A central atom with an octet. Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings

9. Number of Covalent Bonds The number of covalent bonds can be determined from the number of electrons needed to complete an octet. Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings

10. Electron-Dot Formula of SF 2 Write the electron-dot formula for SF 2. STEP 1 Determine the atom arrangement. S is the central atom. F S F STEP 2 Total the valence electrons for 1S and 2F. 1S(6e - ) + 2F(7e - ) = 20e -

11. Electron-Dot Formula SF 2 STEP 3 Attach F atoms to S with one electron pair. F : S : F Calculate the remaining electrons. 20e - - 4 e - = 16e - left STEP 4 Complete the octets of all atoms by placing remaining e - as 8 lone pairs to complete octets.             : F : S : F : or : F─S─F :            

12. Electron-Dot Formula ClO 3 - Write the electron-dot formula for ClO 3 −. STEP 1 Determine atom arrangement. Cl is the central atom. O − O Cl O STEP 2 Add all the valence electrons for 1Cl and 3O plus 1e - for negative charge on the ion. 1Cl(7e - ) + 3 O(6e - ) + 1e − = 26e -

13. Electron-Dot Formula ClO 3 - STEP 3 Attach each O atom to Cl with one electron pair. O −   O : Cl : O Calculate the remaining electrons. 26e - - 6 e - = 20e - left

14. Electron-Dot Formula ClO 3 - STEP 4 Complete the octets of all atoms by placing the remaining 20 e - as 10 lone pairs to complete octets.   −   − : O : : O :         │   : O : Cl : O : or : O─Cl─O :             Note: Bonding electrons can be shown by a

15. Remember: Multiple Bonds In a single bond One pair of electrons is shared. In a double bond, Two pairs of electrons are shared. In a triple bond. Three pairs of electrons are shared.

16. Electron-Dot Formula of CS 2 Write the electron-dot formula for CS 2. STEP 1 Determine the atom arrangement. The C atom is the central atom. S C S STEP 2 Total the valence electrons for 1C and 2S. 1C(4e - ) + 2S(6e - ) = 16e -

17. Electron-Dot Formula CS 2 STEP 3 Attach each S atom to C with electron pairs. S : C : S Calculate the remaining electrons. 16e - - 4 e - = 12e - left

18. Electron-Dot Formula CS 2 STEP 4 Attach 12 remaining electrons as 6 lone pairs to complete octets..... : S : C : S :.... STEP 5 To complete octets, move two lone pairs between C and S atoms to give two double bonds......... : S : : C : : S : or : S = C = S :

19. Multiple Bonds in N 2 In nitrogen N 2, Octets are achieved by sharing three pairs of electrons, which is a triple bond. Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings

20. Some Electron-Dot Formulas Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings 2(1) + 6 = 8

21. Electronegativity and Polarity Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings

22. Electronegativity is the relative ability of atoms to attract shared electrons is higher for nonmetals, with fluorine as the highest with a value of 4.0 is lower for metals, with cesium and francium as the lowest with a value of 0.7 increases from left to right going across a period on the periodic table decreases going down a group on the periodic table Electronegativity

23. Some Electronegativity Values for Group A Elements Low values High values Electronegativity increases Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings

24. A nonpolar covalent bond Occurs between nonmetals. Is an equal or almost equal sharing of electrons. Has almost no electronegativity difference (0.0 to 0.4). Examples: Atoms Electronegativity Type of Bond Difference N-N 3.0 - 3.0 = 0.0 Nonpolar covalent Cl-Br 3.0 - 2.8 = 0.2 Nonpolar covalent H-Si2.1 - 1.8 = 0.3 Nonpolar covalent Nonpolar Covalent Bonds

25. A polar covalent bond Occurs between nonmetal atoms. Is an unequal sharing of electrons. Has a moderate electronegativity difference (0.5 to 1.7). Examples: Atoms ElectronegativityType of Bond Difference O-Cl 3.5 - 3.0 = 0.5Polar covalent Cl-C 3.0 - 2.5 = 0.5Polar covalent O-S 3.5 - 2.5 = 1.0Polar covalent Polar Covalent Bonds

26. Comparing Nonpolar and Polar Covalent Bonds Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings

27. Ionic Bonds An ionic bond Occurs between metal and nonmetal ions. Is a result of complete electron transfer. Has a large electronegativity difference (1.8 or more). Examples: Atoms Electronegativity Type of Bond Difference Cl-K 3.0 – 0.8 = 2.2Ionic N-Na 3.0 – 0.9 = 2.1Ionic S-Cs2.5 – 0.7= 1.8Ionic

28. Electronegativity and Bond Types

29. Predicting Bond Types

30. Use the electronegativity difference to identify the type of bond between the following as: nonpolar covalent (NP), polar covalent (P), or ionic (I). A. K-N2.2ionic (I) B. N-O0.5 polar covalent (P) C. Cl-Cl0.0nonpolar covalent (NP) D. H-Cl0.9polar covalent (P) Learning Check

31. Determining Molecular Polarity The polarity of a molecule is determined from its electron-dot formula shape polarity of the bonds dipole cancellation

32. Polar Bonds and Polar Molecules In water, the molecule is not linear and the bond dipoles do not cancel each other. Therefore, water is a polar molecule.

33. Polar Molecules A polar molecule Contains polar bonds. Has a separation of positive and negative charge called a dipole indicated with  + and  - or arrow. Has dipoles that do not cancel.  +  - H–Cl H — N —H dipole │ H dipoles do not cancel

34. Nonpolar Molecules A nonpolar molecule Contains nonpolar bonds. Cl–Cl H–H Or has a symmetrical arrangement of polar bonds. O=C=O Cl │ Cl –C–Cl │ Cl dipoles cancel

35. Polar Bonds and Nonpolar Molecules For example, the bond dipoles in CO 2 cancel each other because CO 2 is linear.

36. Attractive Forces Between Particles Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings

37. Ionic Bonds In ionic compounds, ionic bonds Are strong attractive forces. Hold positive and negative ions together. Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings

38. Dipole-Dipole Attractions In covalent compounds, polar molecules exert attractive forces called dipole-dipole attractions form strong dipole attractions called hydrogen bonds between hydrogen atoms bonded to F, O, or N, and other atoms that are strongly electronegative

39. Dipole-Dipole Attractions

40. Dispersion Forces Dispersion forces are weak attractions between nonpolar molecules caused by temporary dipoles that develop when electrons are not distributed equally

41. Comparison of Bonding and Attractive Forces

42. Melting Points and Attractive Forces Ionic bonds require large amounts of energy to break apart. Ionic compounds have very high melting points. Hydrogen bonds are the strongest type of dipole- dipole attractions. They require more energy to break than other dipole attractions. Compounds with hydrogen bonds have moderate melting points.

43. Melting Points and Attractive Forces (continued) Dipole-dipole attractions are weaker than hydrogen bonds, but stronger than dispersion forces. They have low to moderate melting points. Dispersion forces are weak and little energy is needed to break them. Compounds with dispersion forces have the lowest melting points.