1. Bonding – General Concepts

6. What is a Bond? A force that holds atoms together.
We will look at it in terms of energy.
Bond energy - the energy required to break a bond.
Why are compounds formed?
Because it gives the system the lowest energy.

7. Ionic Bonding
An atom with a low ionization energy reacts with an atom with high electron affinity.
The electron moves.
Opposite charges hold the atoms together.

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

9. Electronegativity
difference
Bond
Type
Zero
Covalent
Covalent Character
decreases
Ionic Character increases
Polar
Covalent
Intermediate
Ionic
Large

10. Ionic Bonds Electrons are transferred
Electronegativity differences are
generally greater than 1.7
The formation of ionic bonds is
always exothermic!

11. Determination of Ionic Character
Electronegativity difference is not the final determination of ionic character
Compounds are ionic if they conduct electricity in their molten state

12. Coulomb’s Law Q is the charge. r is the distance between the centers.
If charges are opposite, E is negative
exothermic
Same charge, positive E, requires energy to bring them together.
endothermic

13. Size of ions Ion size increases down a group.
Cations are smaller than the atoms they came from.
Anions are larger.
across a row they get smaller, and then suddenly larger.
First half are cations.
Second half are anions.

14. Periodic Trends Li+1 Be+2 B+3 C+4 N-3 O-2 F-1
Across the period nuclear charge increases so they get smaller.
Energy level changes between anions and cations.
Li+1
Be+2
B+3
C+4
N-3
O-2
F-1

15. Ionic Compounds We mean the solid crystal.
Ions align themselves to maximize attractions between opposite charges,
and to minimize repulsion between like ions.
Can stabilize ions that would be unstable as a gas.
React to achieve noble gas configuration

16. Forming Ionic Compounds
Lattice energy - the energy associated with making a solid ionic compound from its gaseous ions.
M+(g) + X-(g) ® MX(s)
This is the energy that “pays” for making ionic compounds.
Energy is a state function so we can get from reactants to products in a round about way.

17. Estimate Hf for Sodium Chloride
Na(s) + ½ Cl2(g)  NaCl(s)
Lattice Energy: Na+(g) + Cl-(g)  NaCl(s)
-786 kJ/mol
Ionization Energy for Na: Na(g)  Na+(g) + e-
495 kJ/mol
Electron Affinity for Cl: Cl(g) + e-  Cl-(g)
-349 kJ/mol
Bond energy of Cl2: Cl2(g)  2Cl(g)
239 kJ/mol
Enthalpy of sublimation for Na: Na(s)  Na(g)
109 kJ/mol

18. Estimate Hf for Sodium Chloride
Na(s) + ½ Cl2(g)  NaCl(s)
Lattice Energy
-786 kJ/mol
Ionization Energy for Na
495 kJ/mol
Electron Affinity for Cl
-349 kJ/mol
Bond energy of Cl2
239 kJ/mol
Enthalpy of sublimation for Na
109 kJ/mol
Na+(g) + Cl-(g)  NaCl(s) kJ
Na(g)  Na+(g) + e kJ
½ Cl2(g)  Cl(g) ½(239 kJ)
Cl(g) + e-  Cl-(g) kJ
Na(s)  Na(g) kJ
Na(s) + ½ Cl2(g)  NaCl(s) kJ/mol

19. 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
Lattice Energies of Salts of the OH- and O2- Ions (kJ/mol)
OH-
O2-
Na+
900
2481
Mg2+
3006
3791
Al3+
5627
15,916

20. Try number 57 from Zumdahl-Zumdahl

21. What about covalent compounds?
The electrons in each atom are attracted to the nucleus of the other.
The electrons repel each other,
The nuclei repel each other.
They reach a distance with the lowest possible energy.
The distance between is the bond length.

22. Covalent Bonds Polar-Covalent bonds Nonpolar-Covalent bonds
Electrons are unequally shared
Electronegativity difference between .3 and 1.7
Nonpolar-Covalent bonds
Electrons are equally shared
Electronegativity difference of 0 to 0.3

23. Covalent Bonding Forces
Electron – electron
repulsive forces
Proton – proton
repulsive forces
Electron – proton
attractive forces

24. Bond Length Diagram

25. Bond Order and Bond lengths
Bond Order (B.O.): describes if a bond is single (B.O. = 1), double (B.O. = 2), triple (B.O. = 3), or some intermediate value (structures with resonance).
This is concept that we will return to, later in this chapter.
Bond Length: the distance between the centers of two atoms joined by a covalent bond.
Higher bond order corresponds to shorter bond lengths, and, in general, to stronger bonds.
Bond Order and Bond lengths 25

26. Bond Order and Bond lengths26

27. Bond-dissociation energy (D): the quantity of energy required to break one mole of covalent bonds in a gaseous species, usually expressed in kJ/mol.
Bond Energies
Bond Breakage:
H2(g)  2 H(g) DH = D(H-H) = kJ/mol
Bond Formation:
2 H(g)  H2(g) DH = -D(H-H) = kJ/mol 27

28. Bond Energies
Average bond energy: the average of bond dissociation energies for a number of different species containing a particular covalent bond.
In H2O, more energy is required to break the first bond than the second. The second bond broken is the O-H radical. The O-H bond energy in H2O is the average of the two values: kJ/mol.
Bond Breakage:
H-OH (g)  H(g) + OH(g) DH = kJ/mol
Bond Breakage:
O-H (g)  H(g) + O(g) DH = kJ/mol 28

29. Bond Energies 29

30. Calculation of enthalpy of reaction from bond energies
Hypothetically, thermochemical data can be obtained from: gaseous reactants  gaseous atoms  gaseous products
Calculation of enthalpy of reaction from bond energies
Some of the BE are likely to be average bond energies rather than bond dissociation energies. 30

31. Calculation of enthalpy of reaction from bond energies
Note: In calculating DH for this reaction, some of the terms cancel out because the same bond types appear as both reactants and products (black bonds above).
Only the net number and types of bonds broken (red bonds above) and formed (blue bonds above) are included in the calculation. 31

32. Calculation of enthalpy of reaction from bond energies32

33. Try number 65 from Zumdahl-Zumdahl

38. Fundamental Properties of Models
A model does not equal reality.
Models are oversimplifications, and are therefore often wrong.
We must understand the underlying assumptions in a model so that we don’t misuse it.

39. The Octet Rule
Combinations of elements tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
Diatomic Fluorine

40. Comments About the Octet Rule
2nd row elements C, N, O, F observe 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.

41. Lewis Structures
 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.

42. Resonance
 Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule.
Benzene, C6H6
  The actual structure is an average of the resonance
structures.
  The bond lengths in the ring are identical, and
between those of single and double bonds.

43. Resonance Bond Length and Bond Energy
 Resonance bonds are shorter and stronger than single bonds.
  Resonance bonds are longer and weaker than double
bonds.

44. Resonance in Polyatomic Ions
Resonance in a carbonate ion:
Resonance in an acetate ion:

45. Localized Electron Model
Lewis structures are an application of the “Localized Electron Model”
L.E.M. says: Electron pairs can be thought of as “belonging” to pairs of atoms when bonding
Resonance points out a weakness in the Localized Electron Model.

46. VSEPR – Valence Shell Electron Pair Repulsion
X + E
Overall Structure
Forms 2
Linear
AX2 3
Trigonal Planar
AX3, AX2E 4
Tetrahedral
AX4, AX3E, AX2E2 5
Trigonal bipyramidal
AX5, AX4E, AX3E2, AX2E3 6
Octahedral
AX6, AX5E, AX4E2
A = central atom
X = atoms bonded to A
E = nonbonding electron pairs on A

47. VSEPR: Linear
AX2
CO2

48. VSEPR: Trigonal Planar
AX3
BF3
AX2E
SnCl2

49. VSEPR: Tetrahedral
AX4
CCl4
AX3E
PCl3
AX2E2
Cl2O

50. VSEPR: Trigonal Bi-pyramidal
AX5
PCl5
AX4E
SF4
AX3E2
ClF3
AX2E3
I3-

51. VSEPR: Octahedral
AX6
SF6
AX5E
BrF5
AX4E2
ICl4-