1. Chemical Bonds Ionic Bond Formation of Ions
Electron Configurations of Ions
Ionic Size and Charge density,
Relative Strength of Ionic Bonds
Lattice Energy
Steps in the Formation of an Ionic Compound
The Born-Haber Cycle

2. Chemical Bonds Covalent Bonds Electronegativity
Polarity of Covalent Bonds
Lewis Structures and the Octet Rule
Exceptions to the Octet Rule
Resonance Lewis Structures
Bond Energies
Calculating Enthalpy using Bond Energy
Molecular Shape - The VSEPR Model

3. Review of Atomic Properties
Effective nuclear charge & Atomic Size:
effective nuclear charge increases left to right and decreases down a group;
electronic shell gets smaller left-to-right across period and gets bigger down a group;
Atomic size decreases left to right across period and increases top to bottom down a group:

4. Review of Atomic Properties
Atomic Size and Ionization Energy:
L-to-R: atomic size decreases; ionization energy increases;
Top-to-bottom: atomic size increases; ionization energy decreases;
Ionization energy increases across period (L-to-R), but decreases down a group;

5. Review of Atomic Properties
Electron affinity increases left to right and decreases top to bottom:
smaller atoms have stronger attraction of added electron and larger atoms
Nonmetals have higher tendency to gain electrons than metals and become anion

6. Review of Atomic Properties
Atomic Size and Electron Affinity:
L-to-R: atomic size decreases, electron affinity increases;
Top-to-bottom: atomic size increases, electron affinity decreases;
Electron affinity increases across period (L-to-R), but decreases down a group;

7. Ionic bonds Attractions between cations and anions;
Bonds formed between metals and nonmetals

8. Formation of Cations
Ions formed when metals react with nonmetals - metal atoms lose valence electrons to nonmetals;
Atoms of representative metals lose valence electrons to acquire the noble gas electron configuration;
Cations of representative group have noble gas electron configurations;

9. Formation of Cations From the alkali metals (1A): M  M+ + e-
From the alkaline Earth metals (2A):
M  M e-
From Group 3A metals:
M  M3+ + 3e- ;

10. Formation of Ions
The nonmetal atoms gain electrons to the noble gas electron configuration;
Anions have noble gas electron configuration;

11. Formation of Anions From the halogen family (VIIA): X + e-  X-
From the oxygen family (VIA):
X + 2e-  X2-
From N and P (in Group VA):
X + 3e-  X3-

12. Common Ions of the Representative Elements
Ions isoelectronic to He (1s2): Li+ & H-
Ions isoelectronic to Ne (1s2 2s2 2p6):
Na+, Mg2+, Al3+, F-, O2-, and N3-
Ions isoelectronic to Ar (1s2 2s2 2p6 3s2 3p6):
K+, Ca2+, Sc3+, Cl-, S2-, and P3-

13. Common Ions of the Representative Elements
Ions isoelectronic to Kr (1s22s22p63s23p64s23d104p6):
Rb+, Sr2+, Y3+, Br-, and Se2-;
Ions isoelectronic to Xe (1s22s22p63s23p64s23d104p65s24d105p6)
Cs+, Ba2+, La3+, I-, and Te2-;

14. Ionic Radii Relative size of isoelectronic ions:
Al3+ < Mg2+ < Na+ < Ne < F- < O2- < N3-;
Sc3+ < Ca2+ < K+ < Ar < Cl- < S2- < P3-;
Trend of ionic radii within a group:
Li+ < Na+ < K+ < Rb+ < Cs+;
F- < Cl- < Br- < I-;

15. Cations From Transition Metals
Transition metal atoms lose variable number of electrons;
Cations have variable charges;
Cations do not acquire noble gas electron configurations

16. Electron Configurations of Transition Metal Cations
Examples:
Cr: [Ar] 4s13d5
Cr  Cr e-; Cr2+: [Ar] 3d4
Cr  Cr e-; Cr3+: [Ar] 3d3
Fe: [Ar] 4s23d6
Fe  Fe e-; Fe2+: [Ar] 3d6
Fe  Fe e-; Fe3+: [Ar] 3d5

17. Charge Density and Strength Ionic Bond
Charge density = charge/size of ion
Greater charge but small ionic radius  High charge density  stronger ionic bond;
Stronger ionic bond  High lattice energy;
Stronger ionic bond  High melting point;

18. Lattice Energy (UL) M+(g) + X-(g)  MX(s); UL = Lattice energy
Lattice energy - energy released when gaseous ions combine to form solid ionic compound:
M+(g) + X-(g)  MX(s); UL = Lattice energy
Examples:
Na+(g) + Cl-(g)  NaCl(s); UL = -787 kJ/mol
Li+(g) + F- (g)  LiF(s); UL = kJ/mol

19. Lattice energy Lattice energy  k(q1q2/r2)
q1 and q2 = charge magnitude on ions;
r = distance between nuclei, and
k = proportionality constant.
Lattice energy increases with charge magnitude but decreases with ionic size

20. Lattice Energies of Some Ionic Compounds
Lattice Energy, UL(kJ/mol)
The energy required to separate a mole of ionic solids into the gaseous/vapor ions;
MX(s)  M+(g) + X-(g)
Mn+/Xn- F- Cl Br- I- O2- Li Na K Mg Ca
______________________________________________________________________

21. The Born-Haber Cycle for NaCl
Na+(g) + Cl(g) _______________
-349 kJ
+496 kJ _______ Na+(g) + Cl-(g)
Na(g) + Cl(g)___________
+121 kJ
Na(g) + ½Cl2(g)________ ? kJ
+108 kJ
Na(s) + ½Cl2(g)________
-411 kJ
NaCl(s)_________________

22. Chemical Processes in the Formation of NaCl
Na(s)  Na(g); DHs = +108 kJ
½Cl2(g)  Cl(g); ½BE = +121 kJ
Na(g)  Na+(g) + e-; IE = +496 kJ
Cl(g) + e-  Cl-(g); EA = -349 kJ
Na+(g) + Cl-(g)  NaCl(s); UL = ? kJ
Na(s) + ½Cl2(g)  NaCl(s); DHf = -411 kJ
UL = DHf – (DHs + ½BE + IE + EA)
(DHs = Enthalpy of sublimation; IE = Ionization energy;
BE = Bond energy; EA = Electtron affinity; UL = Lattice energy; DHf = Enthalpy of formation)

23. The Born-Haber Cycle for LiF
Li+(g) + F(g) _______________
-328 kJ
+520 kJ _______Li+(g) + F-(g)
Li(g) + F(g)___________
+77 kJ
Li(g) + ½F2(g)________ ? kJ
+161 kJ
Li(s) + ½F2(g)________
-617 kJ
LiF(s)_________________

24. Chemical Processes in the Formation of LiF
Li(s)  Li(g); DHs = +161 kJ
½F2(g)  F(g); ½BE = +77 kJ
Li(g)  Li+(g) + e-; IE = +520 kJ
F(g) + e-  F-(g); EA = -328 kJ
Li+(g) + F-(g)  LiF(s); UL = ?
Li(s) + ½F2(g)  LiF(s); DHf = -617 kJ
UL = DHf – (DHs + ½BE + IE + EA)
(DHs = Enthalpy of sublimation; IE = Ionization energy;
BE = Bond energy; EA = Electtron affinity; UL = Lattice energy; DHf = Enthalpy of formation)

25. The Born-Haber Cycle for MgO
Mg2+(g) + O2-(g) _____________
+737 kJ
Mg2+(g) + O(g)________
+2180 kJ
Mg(g) + O(g)_________
+247 kJ
Mg(g) + ½O2(g)________ ? kJ
+150 kJ
Mg(s) + ½O2(g)________
-602 kJ
MgO(s)_________________

26. Chemical Processes in the Formation of MgO
Mg(s)  Mg(g); DHs = +150 kJ
½O2(g)  O(g); ½BE = +247 kJ
Mg(g)  Mg2+(g) + 2e-; IE = kJ
O(g) + 2e-  O2-(g); EA = +737 kJ
Mg2+(g) + O2-(g)  MgO(s); UL = ? kJ
Mg(s) + ½O2(g)  MgO(s); DHf = -602 kJ
UL = DHf – (DHs + ½BE + IE + EA)
(DHs = Enthalpy of sublimation; IE = Ionization energy;
BE = Bond energy; EA = Electron affinity; UL = Lattice energy; DHf = Enthalpy of formation)

27. Covalent Bonds
Bonds between two nonmetals or between a semimetal and a nonmetal atoms
Bonds formed by sharing electron pairs;
One, two or three pairs of electrons shared between two atoms;
A pair of atoms may form single, double, or triple covalent bonds;

28. Potential energy of H-atoms during the formation of H2 molecule

29. Polarity of Covalent Bonds
Covalent bonds - polar or nonpolar;
Nonpolar covalent bonds - bonds between identical atoms or atoms having the same electronegativity.
Polar covalent bonds - bonds between atoms with different electronegativity;

30. Polar Covalent Bonds Bonds have partial ionic character
Bond polarity depends on DEN;
DEN = difference in electronegativity of bonded atoms

31. Electronegativity
Electronegativity = relative ability of bonded atom to pull shared electrons.
Electronegativity Trend:
increases left-to-right across periods; decreases down the group.

32. Electronegativity
Most electronegative element is at top right corner of Periodic Table
Fluorine is most electronegative with EN = 4.0
Least electronegative element is at bottom left corner of Periodic Table
Francium is least electronegative with EN = 0.7

33. General trends:
Electronegativity increases from left to right across a period
For the representative elements (s and p block) the electronegativity decreases as one goes down a group
Electronegativity trend for transition metals is less predictable.

34. Electronegativity and Bond Polarity
Compound F2 HF
LiF
Electronegativity Difference
= 0
= 1.9
= 3.0
Type of Bond
Nonpolar covalent
Polar covalent
Ionic (non-covalent)
In F2 electrons are shared equally and bond is nonpolar
In HF the fluorine is more electronegativity than hydrogen - electrons are drawn closer to fluorine.
H―F bond is very polar

35. Electronegativity and bond polarity
The H-F bond can thus be represented as:
The 'd+' and 'd-' symbols indicate partial positive and negative charges.
The arrow indicates the "pull" of electrons off the hydrogen and towards the more electronegative atom.
In lithium fluoride the much greater relative electronegativity of the fluorine atom completely strips the electron from the lithium and the result is an ionic bond (no sharing of the electron)

36. Predicting Bond Type From Electronegativity
General rule of thumb for bonds:
DEN = 0-0.4, bond is non-polar covalent;
DEN > 0.4, but < 1.9, bond is polar covalent
DEN > 1.8, bond is considered ionic.

37. Potential Energy Diagram for Covalent Bond Formation

38. Bond Length Bond length - distance between the nuclei of bonded atoms.
The larger the atoms that are bonded, the greater the bond length.
Bond length:
single bonds > double bonds > triple bonds . . .

39. Bond Energy
Bond energy - the energy required to break the bonds between two atoms. 
The shorter the bond, the greater the bond energy.
Bond energy:
Triple bonds > double bonds > single bond

40. Bond Length and Bond Energies
Bond length (pm) and bond energy (kJ/mol)
Bond Length Energy Bond Length Energy
_________________________________________________________________________________________________________
H─H H─C
C─C H─N
N─N H─O
O─O H─F
F─F H─Cl
Cl─Cl H─Br
Br─Br H─I
I─I        C─F
C─S C─Cl
C─C C─Br
C─N C─I
C─O     C─C
O─O C=C
O=O C≡C    
C=O N=N ? 418
C=N N≡N

41. Bond Breaking and Bond Formation in the Reaction to form H2O

42. Using Bond Energy to Calculate Enthalpy
Chemical reactions in the gaseous state only involve:
the breaking of covalent bonds in reactants and the formation of covalent bonds in products.
Bond breaking requires energy
Bond formation releases energy
DHreaction = S(Energy of bond breaking) + S(Energy of bond formation)

43. Calculating Enthalpy Reaction Using Bond Energy
Example: use bond energy to calculate DH for the following reaction in gaseous state:
CH3OH + 2 O2  CO2 + 2H2O;
DHreaction = S{BE(in reactants)} - S{BE(in products)}

44. Using bond energy to calculate enthalpy
S{BE(in reactants) =
3 x BE(C─H) + BE(C─O) + BE(O─H) + 2 x BE(O═O)
= (3 x 413) (2 x 495) = 3054 kJ
S{BE(in products)
= 2 x BE(C═O)* + 4 x BE(O─H)
= (2 x 799) + (4 x 495) = 3578 kJ
DHreaction = S{BE(in reactants)} - S{BE(in products)}
DHreaction = 3054  3578 = -524 kJ

45. Lewis Structures for Molecules or Polyatomic ions
Step-1:
Calculate number of valence electrons;
For polyatomic ions, add one additional electron for each negative charge, or subtract one for each positive charge on the ion.

46. Lewis Structures for Molecules and Polyatomic ions
Step-2:
Choose a central atom (the least electronegative atom)
(Hydrogen and Fluorine cannot become central atoms)
Connect other atoms to the central atom with single bonds (a pair of electrons).

47. Lewis Structures for Molecules and Polyatomic ions
Step-3:
Complete the octet state of all terminal atoms, except hydrogen.
Place remaining pairs of electrons (if present) on central atom as lone pairs.

48. Octet State of Central Atom
Step-4:
If central atom has not acquired octet state but no more electrons available, move lone-pair electrons from terminal atoms, one pair at a time, to form double or triple bonds to complete octet of the central atom.

49. Lewis Symbols and Formation of Covalent Molecules

50. Lewis Structures of CH4, NH3 and H2O

51. Lewis Structures of CO2, HCN, and C2H2

52. Resonance Lewis Dot Structures for CO32-

53. Exception to Octet Rule
If central atom is from group 2A or 3A, octet state is not acquired - the central atom has incomplete octet.
Central atoms from periods 3, 4, 5, …may have more than 8 valence electrons (expanded octet)
Molecules with odd number of electrons will contain unpaired electrons.

54. Covalent Molecules with Central Atoms have Expanded Octet State

55. Evaluate Formal Charge
Evaluate formal charges (fc) on each atom in the molecule to determine best correct or best Lewis structures.
Formal charge is apparent charge on an atom in a Lewis formula; it is determined as follows:
Formal charge =
(Atom’s group #) – (# of lone-pair electrons on the atom) – (# of covalent bonds the atom forms)

56. Assigning Formal Charges

57. Choosing the correct or best Lewis structures based on formal charges
If two or more Lewis dot structures that satisfy the octet rule can be drawn, the most stable one will be the structure in which:
The formal charges are as small as possible.
Any negative charges are located on the more electronegative atoms.

58. Which Lewis structures of CO2 & N2O are correct?

59. The Shape of Water Molecules

60. Molecular Shapes of BeI2, HCl, IF2-, ClF3, and NO3-

61. Lewis Structures, Molecular Shapes & Polarity

62. The Shapes of Methane and Ammonia Molecules

63. The VSEPR Shapes

64. Linear and Trigonal Planar Electron-Pair Geometry

65. The Tetrahedral Electron-Pair Geometry

66. Trigonal Bipyramidal Electron-Pair Geometry

67. The Octahedral Electron-Pairs Geometry

68. Lewis Structures of HF, H2O, NH3, & CH4

69. Lewis Symbols for O, F, and Na

70. Lewis Model for the Formation of Covalent Bonds and Covalent Molecules

71. Covalent Bonds and Lewis Structures Some Molecules

72. Resonance Lewis Structures of PO43-

73. Assigning Appropriate Formal Charges

74. Structures and Shapes of Formaldehyde and Ethylene