1. CH110 Chapter 4: Compounds & Bonds
Valence Electrons & e– Dot Structures
Octet Rule & Ions
Ionic Compounds
Covalent Compounds
Molecular Shapes & Polarity 1 1

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

4. H Li Na K H Li Lewis Structures Show only Valence Electrons Na 1 7 323 11 Na K

5. He O F N Si S P Ca H Li C Na K Se B Ne Be Al Cl Ar Mg Kr Ga As Br Ge 18 H He 2 3 4 5 6 7 C O Li B N F Ne Be Si Al S Cl Ar Mg P Na Kr Ca Ga As Se Br K Ge

6. Na Na1+or Na+ Ions Na Metals give e-s to make Cations (+) 11 +’s23 11 Na
11 +’s
11 -’s Na
2, 8, 1
Metals give e-s to make Cations (+)
11 +’s
10 -’s
1 +
Na1+or Na+
2, 8 = [Ne]

7. Cl Cl = Cl1– Cl– Ions Cl Nonmetals take e-s to make Anions (–) 17 +’s35 17 Cl
17 +’s
17 -’s Cl
2, 8, 7
Nonmetals take e-s to make Anions (–) Cl 1-
17 +’s
18 -’s
1 -
= Cl1–
Cl–
2, 8, 8 = [Ar]

8. Na + Cl Na+ + Cl _ 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. 15 15

9. Ionic compounds NaCl sodium chloride Not individual molecules
Form crystal arrays
Ions touch many others
Formula represents the average ion ratio
NaCl
sodium chloride Na Cl Cl Na Cl Na 18 18

10. Common ions Representative Elements 1+ 4+ 4- 2+ 3+ 3- 2- 1- H Li Na CsRb K Tl Hg Au Hf Ls Ba Fr Pt Ir Os Re W Ta He Rn At Po Bi Pb Be Mg Sr Ca Cd Ag Zr Y Pd Rh Ru Tc Mo Nb Ac Ra Zn Cu Ti Sc Ni Co Fe Mn Cr V In Xe I Te Sb Sn Ga Kr Br Se As Ge Al Ar Cl S P Si B Ne F O N C Gd Cm Tb Bk Sm Pu Eu Am Nd U Pm Np Ce Th Pr Pa Yb No Lu Lr Er Fm Tm Md Dy Cf Ho Es 2+ 3+ 3- 2- 1-
4 - 6 12 12

11. Na1+ Cl1- Al3+ Cl1- Cl1- Cl1- NaCl AlCl3 Ionic Formulas
Metal Cations + Nonmetal Anions
Na1+
Cl1-
Al3+
Cl1-
Cl1-
Cl1-
NaCl
AlCl3
Sodium Chloride
Aluminum Chloride

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

14. Information in the table
Atomic number 26
Atomic mass (weight)
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. Fe2+ Cl1- Fe3+ Cl1- FeCl2 FeCl3 Transition Metal Ions Cl1- Cl1- Cl1-
Iron(II) Chloride
Iron(III) Chloride
Ferrous Chloride
Ferric Chloride

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

18. Names and Formulas of Common Polyatomic Ions

19. Ionic compounds Na3N NaCl Na2O MgCl2 MgO Mg3N2 AlCl3 Al2O3 AlN
Some simple ions
Anions
Cl-
O2-
N3-
Na+
Na3N
NaCl
Na2O
Cations
Mg2+
MgCl2
MgO
Mg3N2
Al3+
AlCl3
Al2O3
AlN 17 17

20. Ionic compounds NaBr Na2O Na3N MgBr2 MgO Mg3N2 AlBr3 Al2O3 AlN FeBr3
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. Covalent Bonds H H H + Cl Cl + Cl O O O + N N N +

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

24. O=O 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 74 74

25. CO C O C O CO2 O C O O=C=O Covalent Bonds Carbon monoxide
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
Cl2O
N2O5
SiO2
ICl3
P2O5
CCl4
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. 34 34

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

30. Electronegativity Cl H Relative ability of atoms to attract e-. 1.0Br Cl Po Te Se S Bi Sb As P Pb Sn Ge Si F O N Tl Na Cs Rb K Ba Mg Sr Ca In Ga Al H Li Be B C
1.0
0.8
0.9
1.5
1.2
0.7
2.0
1.6
1.7
1.8
2.5
1.9
3.0
2.1
3.5
2.4
4.0
2.8 66 66

31. H Cl – + Bond Polarity, Electronegativity Cl H
This unequal sharing results in polar bonds. + –
H Cl
Slight positive side
Smaller electronegativity
Slight negative
Larger electronegativity 67 67

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

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

34. O=C=O Polarity, Shape O C O d– d+ d– CO2 Electronegativity Difference
< 0.5 Nonpolar
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. O=C=O Polarity, Shape Linear d– d+ d– Nonpolar Compound d–
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. Polarity, Shape C H Cl C H Cl O H H-O-H d+ d– Tetrahedral d+ d– d+
e–’s in 4 directions = 109.5o C H Cl C H Cl d+ d–
Tetrahedral d+
4 directions = 109.5o O H d– d+
H-O-H
Bent

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

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

40. Tetrahedral electron-pair Geometries
Bent
Pyramidal
Tetrahedral H    O
109.5º 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. 55 55

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? 68 68

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” 69 69

44. Boiling and melting points
Chemical Bond Mp Bp
N2 Nonpolar
O Nonpolar
NH Polar
H2O Polar
NaCl Ionic ?
Melting and Boiling points
Very high for ionic compounds
Typically lower for covalent compounds 77 77