1. Bonding: General Concepts
Chapter 13
Bonding:
General Concepts

2. Types of Chemical Bonds
Ionic bonding
Polar covalent bonding
Covalent bonding

4. Lennard-Jones 6-12 potential

5. Ionic bonding
Ionic substances are formed when an atom that loses electrons relatively easily react with an atom that has a high affinity for electrons.
ex. metal-nonmetal compound

6. Covalent Bonding
Electron are shared by nuclei

7. Polar Covalent Bonding
A polar bond is a covalent bond in which there is a separation of charge between one end and the other , in other words in which one end is slightly positive and the other slightly negative.

8. Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself.

9. Calculate Electronegativity
H atom
∆HF=565-1/2( )=272
|xH-xF|=0.102(272)1/2=1.68, xH-4.0=-1.7, xH=2.3
O atom
∆OF=190-1/2( )=40
|xO-xF|=0.102(40)1/2=0.65, xO-4.0=-0.65, xO=3.4
C atom
∆CF=485-1/2( )=234.5
|xC-xF|=0.102(234.5)1/2=1.6, xC-4.0=-1.6, xC=2.4

13. Bond Polarity and Dipole Moments
μ=QR
Q: center of charge of magnitude
R: distance

14. Dipole Moment of HF 1D=3.336×10-30 coulomb meter
μ=(1.6×10-19 C)(9.17×10-11 m)=1.47×10-29
=4.4 D for fully ionic
Measured dipole moment=1.83 D
1.83×3.336×10-30=δ(9.17×10-11)
δ=6.66×10-20
Ionic character=1.83/4.4=41.6%

15. In practice no bond is totally ionic
In practice no bond is totally ionic. There will always be a small amount of electron sharing.  

17. The compounds with more 50% ionic character are normally considered to be ionic solids.

18. Dipole Moment of Polyatomic Molecules
For dipole moment of polyatomic molecules, the dipole is the geometric sum of all bond dipole moment.

22. Achieving Noble Gas Electron Configurations (NGEC)
Two nonmetals react: They share electrons to achieve NGEC.
A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.

23. Isoelectronic Ions Ions containing the the same number of electrons
O2> F > Na+ > Mg2+ > Al3+
largest smallest

24. Formation of Binary Ionic Compounds
Lattice energy:
The change in energy that takes
place when separated gaseous
ions are packed together to form an
ionic solid.
M+(g)+X-(g) →MX(s)

25. Formation of an Ionic Solid
1. Sublimation of the solid metal
M(s) → M(g) [endothermic]
2. Ionization of the metal atoms
M(g) →M+(g) + e- [endothermic]
3. Dissociation of the nonmetal
1/2X2(g) → X(g) [endothermic]

26. Formation of an Ionic Solid (continued)
4. Formation of X ions in the gas phase:
X(g) + e- → X-(g) [exothermic]
5. Formation of the solid MX
M+(g) + X-(g) → MX(s) [quite exothermic]

27. Electron affinity of F
Dissociation of F2
Ionization of Li
Formation
of solid
Sublimation of Li

28. Lithium-Fluoride structure

29. Lattice Energy Calculations
k: a proportionality constant that depends on the structure of the solid and the electron configuration of the ions
Q1 and Q2: charges on the ions
r: the shortest distance between the centers of cations and anions

31. Lattice Energies and the Strength of the Ionic Bond
The strength of the bond between the ions of opposite charge in an ionic compound depends on the charges on the ions and the distance between the centers of the ions when they pack to form a crystal.
An estimate of the strength of the bonds in an ionic compound can be obtained by measuring the lattice energy of the compound.

32. Lattice Energies for Alkali Metals Halides
The bond between ions of opposite charge is strongest when the ions are small.
The lattice energies for the alkali metal halides is therefore largest for LiF and smallest for CsI.

33. Lattice Energies of Alkali Metals Halides (kJ/mol)
Cl-
Br- I-
Li+
1036
853
807
757
Na+
923
787
747
704 K+
821
715
682
649
Rb+
785
689
660
630
Cs+
740
659
631
604

34. Lattice Energies for Salts of the OH- and O2- Ions
The ionic bond should also become stronger as the charge on the ions becomes larger.
The lattice energies for salts of the OH- and O2- ions increase rapidly as the charge on the ion becomes larger.

35. Lattice Energies of Salts of the OH- and O2- Ions (kJ/mol)
Na+
900
2481
Mg2+
3006
3791
Al3+
5627
15,916

36. Lattice Energies and Solubility
The lattice energy of a salt gives a rough indication of
the solubility of the salt in water because it reflects the energy needed to separate the positive and negative ions in a salt.
Sodium and potassium salts are soluble in water because they have relatively small lattice energies. Magnesium and aluminum salts are often much less soluble because it takes more energy to separate the positive and negative ions in these salts.
NaOH is very soluble in water (420 g/L), but Mg(OH)2 dissolves in water only to the extent of g/L, and Al(OH)3 is essentially insoluble in water.

37. The Covalent Chemical Bond

38. Experimental result : 1652 kJ/mol C(g)+4H(g) →CH4(g) + 1652 kJ/mol
Bond Energy of CH4
Experimental result : 1652 kJ/mol
C(g)+4H(g) →CH4(g) kJ/mol
An average C-H bond energy per mole
of C-H bond: 1652/4=413 (kJ/mol)

39. Stepwise Decomposition of CH4
CH4(g) →CH3(g)+H(g) 435 kJ/mol
CH3(g) →CH2(g)+H(g) 453 kJ/mol
CH2(g) →CH(g)+H(g) kJ/mol
CH(g) →C (g)+H(g) kJ/mol

40. H = D(bonds broken)  D(bonds formed)
Bond Energies
Bond breaking requires energy (endothermic).
Bond formation releases energy (exothermic).
H = D(bonds broken)  D(bonds formed)
energy required
energy released

41. (C-Cl)+3(C-H)=1578 (C-Cl)+3(413)=1578 C-Cl=339 (kJ/mol)
Bond Energy of CH3Cl
C(g)+Cl(g)+3H(g) →CH3Cl(g)+1578kJ/mol
(C-Cl)+3(C-H)=1578
(C-Cl)+3(413)=1578
C-Cl=339 (kJ/mol)

42. Covalent Bond Energies and Chemical Reactions
H2+F2→2HF
ΔH=ΣD (bonds broken)-ΣD (bonds formed)
ΔH=DH-H+DF-F-2DH-F=1×432+1×154-2×565
=-544 kJ

43. CH4+2Cl2+2F2→CF2Cl2+2HF+2HCl
Reactants bonds broken:
CH4: 4×413=1652,
2Cl2: 2×239=478
2F2: 2×154= Total energy required: 2438kJ
Products bonds formed:
CF2Cl2: 2×485=970 (C-F) and 2×339=678 (C-Cl)
HF: 2×565=1130
HCl: 2×427=854 Total energy released: 3632 kJ
ΔH= =-1194 kJ (-1126 kJ)

44. Localized Electron Model
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
Description of valence electron arrangement (Lewis structure).
Prediction of geometry (VSEPR model).
Description of atomic orbital types used to share electrons or hold long pairs.

45. Lewis Structure
Shows how valence electrons are arranged among atoms in a molecule.
Reflects central idea that stability of a compound relates to noble gas electron configuration.

46. Resonance
Occurs when more than one valid Lewis structure can be written for a particular molecule.
These are resonance structures. The actual structure is an average of the resonance structures.

47. Comments About the Octet Rule
2nd row elements C, N, O, F obey the octet rule.
2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.
3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.
When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

48. Formal Charge
One method involves estimating the charge on each atom in the various possible Lewis structures and using the charges to select the most appropriate structure.
It allows chemists to determine the location of charge in a molecule as well as compare how good a Lewis structure might be.

49. Calculation of Formal Charge
Formal Charge=(number of valence electrons on a free atom)-(number of valence electrons assigned to the atom in the molecule)

50. Assumptions for Formal Charge
Lone pair electrons belong entirely to the atom in question.
Shared electrons are divided equally between the two sharing atoms.
(Valence electrons) assigned=(number of lone
pair electrons)+1/2(number of shared
electrons)

51. Consider the molecule H2CO2-1 +1

52. Consider the molecule H2CO2

53. The two possible Lewis structures are
shown above. They are connected by a
double headed arrow and placed in brackets.
The non-zero formal charge on any atoms
in the molecule have been written near the
atom.

54. Valence Shell Electron Pair Repulsion (VSEPR Model)
It is used to predict the geometries of molecules formed from nonmetals.
Postulate: the structure around a given atom is determined principally by minimizing electron pair repulsion.
The bonding and nonbonding pairs should be positioned as far apart as possible.

55. Predicting a VSEPR Structure
Draw Lewis structure.
Put pairs as far apart as possible.
Determine positions of atoms from the way electron pairs are shared.
Determine the name of molecular structure from positions of the atoms.

56. For non-metals compounds, four pairs of
electrons around a given atom prefer prior
to form a tetrahedral geometry to minimize
the electron repulsions.

57. Draw the Lewis structure
Count the pairs of electrons and arrange them to minimize repulsions
Determine the positions of the atoms
Name the molecular structure

59. Lone pairs require more space than bonding pair.
The bonding pairs are increasingly squeezed together as the number of lone pairs increases.

60. The bonding pair is shared between two nuclei; and the electrons can be close to either nucleus.
A lone pair is localized on only one nucleus, so both electrons are close to that nucleus only.

61. Lone pairs require more room than bonding pairs
square planar