CH110 Chapter 4: Compounds & Bonds

1. CH110Chapter 4: Compounds & Bonds • Valence Electrons & e– Dot Structures • Octet Rule & Ions • Ionic Compounds • Covalent Compounds • Molecular Shapes & Polarity

2. 24 12 Mg Electron arrangement Electrons fill layers around nucleus Low  High 32 18 A new layer is added for each row or period in the table. 8 2 2n2 Shells = Energy levels

3. 1 1 1 4 2 H He 9 4 20 10 7 3 Be Ne Li 40 18 Ar 24 12 Mg 23 11 Na 8 Octet Rule 2 2, 1 2, 2 2, 8 2, 8, 8 2, 8, 1 2, 8, 2

4. 1 1 H H 7 3 Li Li Na 23 11 Na K Lewis Structures Show only Valence Electrons

5. Ge Ga Be H He Ca Mg Al B Li Na Si C P N As S O F Se Cl Br Kr Ar Ne K 1 8 2 3 4 5 6 7

6. Na 23 11 Na Ions 11 +’s 11 -’s 0 2, 8, 1 Metals give e-s to make Cations (+) 11 +’s 10 -’s 1 + Na1+or Na+ 2, 8 = [Ne]

7. 35 17 Cl Cl 1- Cl Cl– Ions 17 +’s 17 -’s 0 2, 8, 7 Nonmetals take e-s to make Anions (–) 17 +’s 18 -’s 1 - = Cl1– 2, 8, 8 = [Ar]

8. Formation of NaCl e–moves from Metal  Nonmetal Stable octets _ Na + Cl Na+ +Cl Metal Cation Nonmetal Anion + and - ions attract to form an ionic bond.

9. Ionic compounds • Not individual molecules • Form crystal arrays • Ions touch many others • Formula represents the average ion ratio NaCl sodiumchloride Na Cl Cl Na Cl Na

10. H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Ls Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Common ions Representative Elements 1+ 4+ 4- 2+ 3+ 3- 2- 1- 4 - 6

11. Ionic Formulas Metal Cations +NonmetalAnions Na1+ Cl1- Al3+ Cl1- Cl1- Cl1- NaCl AlCl3 SodiumChloride Aluminum Chloride

12. Sc Ti V Cr Mn Fe Co Ni Cu Zn Y Zr Nb Mo Tc Ru Rh Pd Ag Cd Ls Hf Ta W Re Os Ir Pt Au Hg Ac Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Common ions Representative Elements 1+ 4+ 4- H He 2+ 3+ 3- 2- 1- Transition Elements Li Be B C N O F Ne Variable Na Mg Al Si P S Cl Ar K Ca Ga Ge As Se Br Kr Rb Sr In Sn Sb Te I Xe Cs Ba Tl Pb Bi Po At Rn Fr Ra

14. Information in the table Atomic number Atomic mass (weight) 26 55.845 Possible Charges (Valence) 2,3 Fe Elemental Symbol Electronic Configuration No longer discussed in text [Ar]3d64s2 Iron Name of the element

16. Transition Metal Ions Fe2+ Cl1- Fe3+ Cl1- Cl1- Cl1- Cl1- FeCl2 FeCl3 Iron(II)Chloride Iron(III) Chloride FerrousChloride Ferric Chloride

17. Polyatomic Ions PO43- Na1+ SO42- NH41+ Na1+ NH41+ NH41+ Na2SO4 (NH4)3PO4 SodiumSulfate Ammonium Phosphate

18. Names and Formulas of Common Polyatomic Ions

19. Ionic compounds Some simple ions Anions Cl- O2- N3- Na+ Na3N NaCl Na2O Cations Mg2+ MgCl2 MgO Mg3N2 Al3+ AlCl3 Al2O3 AlN

20. Ionic compounds Anions Br1- O2- N3- NaBr Na2O Na3N Na1+ Sodium Bromide Sodium Oxide Sodium Nitride Mg2+ MgBr2 MgO Mg3N2 Cations Magnesium Bromide Magnesium Oxide Magnesium Nitride AlBr3 Al2O3 AlN Al3+ Aluminum Bromide Aluminum Oxide Aluminum Nitride FeBr3 Fe2O3 FeN Fe3+ Iron(III) Bromide Ferric Bromide Iron(III) Oxide Ferric Oxide Iron(III) Nitride Ferric Nitride Cu1+ CuBr Cu2O Cu3N Copper(I) Bromide Cuprous Bromide Copper(I) Oxide Cuprous Oxide Copper(I) Nitride Cuprous Nitride

21. Thanks to Christine Neighbors (Fall 2012)

22. H H H H Cl Cl Cl Cl O O O O N N N N Covalent Bonds + + + +

23. H H Cl Cl O O N N N N Covalent Bonds H-H H2 Cl-Cl Cl2 O=O O2 N2

24. Covalent compounds • Covalent compounds • Discrete molecular units • Atoms held together by bonds • Covalent compounds exist in all states • (CO2 - gas, H2O - liquid, SiO2 - solid) • Formula represents atoms in a molecule O=O

25. C O Covalent Bonds CO Carbon monoxide C O CO2 Carbon dioxide O C O O=C=O May modify rules to improve the sound. Example - use monoxide not monooxide.

26. Naming Covalent Compounds In the names of covalent compounds, prefixes are used to indicate the number of atoms (subscript) of each element. (mono is omitted for the first element, not the second) Prefixes Used in Naming Covalent Compounds

27. Rules for Naming Binary Compounds Containing Two Nonmetals • Write the name of the first nonmetal as it appears on the periodic table. Use a prefix if there is more than one atom. • Use a prefix to indicate the number of atoms for the second nonmetal. • Write the stem of the second nonmetal. • Add the suffix –ide.

28. Naming covalent compounds carbon monoxide carbon dioxide dichlorine monoxide dinitrogen pentoxide silicon dioxide iodine trichloride diphosphorous pentoxide carbon tetrachloride May modify rules to improve the sound. Example - use monoxide not monooxide.

29. Cl H Bond Polarity, Electronegativity H Cl Electrons in covalent bonds rarely get shared equally.

30. 2.1 Cl H H 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Li Be B C N O F Na Mg Al Si P S Cl 0.9 1.5 1.8 2.1 2.5 3.0 1.2 K Ca Ga Ge As Se Br 0.8 1.0 1.6 1.8 2.0 2.4 2.8 Rb Sr In Sn Sb Te I 0.8 1.0 1.7 1.8 1.9 2.1 2.5 Cs Ba Tl Pb Bi Po At 1.8 1.9 1.9 2.0 2.1 0.7 0.9 Electronegativity • Relative ability of atoms to attract e-.

31. Cl H H Cl Bond Polarity, Electronegativity • This unequal sharing results in polar bonds. + – • Slight positive side • Smaller electronegativity • Slight negative • Larger electronegativity

32. Cl H H Cl Bond Polarity, Electronegativity This unequal sharing results in polar bonds. – + • Slight positive side • Smaller electronegativity • Slight negative side • Larger electronegativity

33. Cl H Bond Polarity, Electronegativity Electronegativity Difference < 0.5 Nonpolar 0.5-1.7 Polar >1.8 Ionic 2.1 3.0 H Cl d+ d– Polar Covalent

34. Polarity, Shape CO2 Electronegativity Difference < 0.5 Nonpolar 0.5-1.7 Polar >1.8 Ionic O C O 3.5 2.5 3.5 O=C=O d– d+ d– Polar Covalent Bonds Linear Shape (180o) Nonpolar Compound

35. Polar and Non-Polar Thanks to Paula H. & Judy M. (Summer 1976)

36. Polarity, Shape e–’s in 2 directions = 180o O=C=O Linear d– d+ d– Nonpolar Compound e–’s in 3 directions = 120o d– Trigonal planar Polar Compound d+

37. H H C H H C H H Cl Cl O H H H-O-H Polarity, Shape e–’s in 4 directions = 109.5o d+ d– Tetrahedral d+ 4 directions = 109.5o d– d+ Bent

38. N N H H H H H H Polarity, Shape e–’s in 4 directions = 109.5o d- d+ d+ d+ Pyramidal

39. Some common geometries e- directions around Shape central atom Example___ Linear 2 O=C=O Trigonal Planar 3 Tetrahedral 4

40. Tetrahedral electron-pair Geometries Pyramidal Bent Tetrahedral H    109.5º O N C H H H H H H 105º H H 107º Water, H2O 2 bond pairs Ammonia, NH3 3 bond pairs Methane, CH4 4 bond pairs

41. Molecular geometry • Molecules have specific shapes. • Determined by the number of electron pairs around the central atom • Bonded and unbonded pairs • Geometry affects factors like polarity and solubility.

42. Geometry and polar molecules • For a molecule to be polar • - must have polar bonds • - must have the proper geometry • CH4 non-polar • CH3Cl polar • CH2Cl2 polar • CHCl3 polar • CCl4 non-polar • WHY?

43. Polarity and solubility • Solubility The maximum amount of a solute • that dissolves in a given solvent • Depends on the forces of attraction between molecules - intermolecular • Types of intermolecular attractions most often encountered • Dipole-Dipole • Hydrogen bonding • Van der Wall forces • General rule “Like dissolves like”

44. Boiling and melting points • Chemical Bond Mp Bp • N2Nonpolar -210 -196 • O2 Nonpolar -219 -183 • NH3Polar -78 -33 • H2O Polar 0 100 • NaCl Ionic 804 ? • Melting and Boiling points • Very high for ionic compounds • Typically lower for covalent compounds