1. CH 8: Bonding
General Concepts

2. Chapter Outline – Part I
Types of chemical bonds (8.1)
Electronegativity and bond polarity (8.2/3
Ions (8.4)
Energy changes when a binary ionic compound forms (8.5)
Ionic character of covalent bonds (8.6)

3. Introduction to Bonding
Chemical bond – force that holds atoms together so that they function as a unit.
Consider 2 classes of bonds:
Ionic bonding
Covalent bonding

4. Bond Types
Ionic bonds – attractive forces among oppositely charged ions
Forms when a metal loses electron(s) to a nonmetal.
Bond strength can be calculated using Coulomb’s law

5. Ionic Bonds
Strength of the attraction between the ions can be calculated using Coulomb’s law.
E = (2.31 x J nm) (Q1Q2/r)
Q1 and Q2 are the charges on the ions.
r = distance between ion centers in nm

6. Using Coulomb’s Law E = (2.31 x 10-19 J nm) (Q1Q2/r) Sign on E???
The more negative E, the stronger the attractive force between the ions.

7. Using Coulomb’s Law E = (2.31 x 10-19 J nm) (Q1Q2/r) Magnitude of E.
E is more negative when:

8. Covalent Bonds
Covalent bond – bonded atoms share pairs of valence electrons
Covalent bonding results in formation of a molecule.
Covalent bonding occurs between nonmetals.

9. Types of Covalent Bonds
“Pure” covalent bond – electrons are shared by like nonmetals
E.g. diatomic molecules
Results in equal sharing of the electrons
Aka – nonpolar covalent bond

10. Types of Covalent Bonds
Polar covalent bond – unequal sharing of electrons by the bonded atoms
bond between different nonmetals each with its own ability to attract the shared electrons

11. Polar Covalent Bonds Showing bond polarity:
Consider the HF molecule.
See board and/or page 290.
Experimental determination of bond polarity, page 289

12. Bond Polarity
To predict bond polarity…consider the electronegativity (EN) of the bonded atoms.
EN – the ability of an atom in a molecule to attract shared electrons.

13. EN Values
The higher the EN the greater the atom’s ability to attract shared electrons.
EN values and the periodic table
EN ________ down a group.
EN ________ across a period.
See “back” of the periodic table.

14. EN and Bond Polarity
As the difference in EN between bonded atoms increases so does the polarity of the bond.
Can also say that the ionic character of the bond is increasing.
See table 8.1 on page 289.

15. Bond Polarity and Dipoles
Polar molecules have a preferred orientations when placed in an electric field.
Said to have a dipole moment.
Dipole moment – molecule has a center of positive charge and a center of negative charge

16. Bond Polarity and Dipoles
Not all molecule with polar bonds have dipole moments!
Bond polarities cancel each other in molecules with symmetrical dipoles.
Molecules with equal, opposing dipoles.
See page 291and 8.2 on page 292
Dog walking example!

17. Compound Formation
Atoms gain, lose, or share enough electrons to achieve the same stable electron configuration as a noble gas
Nonmetals share electrons
Form molecules with covalent bonds
Representative metals lose electrons to nonmetals in ionic compounds
Ions are isoelectronic to noble gases

18. 8.4 Ion Formation Binary ionic compounds
The metal loses electron(s) to a nonmetal
Focus on represenative metals
The atoms lose/gain enough electrons to obtain a noble gas electron configuration.

19. Cations Group IA metals form ions with a _____ charge. Na atom Na ion
Isoelectronic to: ____________________

20. Anions Group VIA elements form ions with a ______ charge. Sulfur atom
Sulfur ion (called _________________)
Isoelectronic to: ____________________

21. Ionic Compound Consider the compound formed between sodium and sulfur.
Each sodium atom loses 1 electron.
Each sulfur atom needs 2 electrons.
Formula for compound:

22. Ion Size Cations are smaller than their parent atom.
Atoms lose their valence shell when the ion forms.
Na s22s22p63s1 ____ protons
Na s22s22p6 ____ protons

23. Ion Size Anions are larger than their parent atom.
Atoms add electrons to their valence shell when the ion forms – proton # remains the same.
F s22s22p5 ____ protons
F s22s22p6 ____ protons

24. Ion Size
Isoelectronic ions decrease in size as the number of protons increases.
Example: ions with 10 electrons
10 e O F Na Mg2+ Al3+
# p

25. Isoelectronic Ions 10 e O2- F1- Na1+ Mg2+ Al3+ # p 8 9 11 12 13
Radius*
* picometers
The diagram on page 296 should make sense.

26. 8.5 Energy in Binary Ionic Compounds
Lattice energy – change in energy when separated gaseous ions form an ionic solid.
M+(g) X-(g)  MX(s) LE < 0

27. Lattice Energy LE = k (Q1Q2)/r K is the proportionality constant
Q1 and Q2 are the charges on the ions
r is the ionic radius

28. Lattice Energy
LE becomes more exothermic as the ion charges increase and the ion radius decreases.
Small highly charged ions have more exothermic LE
See board for examples.

29. Formation of ionic compounds.
Consider energy changes associated with formation of a binary ionic compound.
5 step process, page 297/298
Most common series of steps is shown on the next slide.

30. Formation of ionic compounds.
Sublime the metal.
Ionize the gaseous metal atoms.
ionization energy(ies)
Dissociate the nonmetal (if diatomic).
Bond energy
Ionize the gaseous nonmetal atoms.
Electron affinity
Form the solid from the gaseous ions LE

31. Born Haber Practice… NaF MgF2 IE – page 272; EA – page 275
Bond energies – page 306
Sublimation: Na 109 kJ; Mg 147 kJ
Lattice energy: NaF -923kJ; MgF2 2913

32. 8.6 Partial Ionic Character
When atoms with different EN bond the result is either a polar covalent or an ionic bond.
There’s evidence that some level of electron sharing occurs in all bonds.
Even in what we consider as ionic bonds.

33. 8.6 Partial Ionic Character
Classify a bond as ionic if it conducts electricity when melted.
Essentially all compounds with metals meet this criteria.
These compounds generally have more than 50% ionic character.

34. 8.8 Covalent Bond Energies
Strength of a given bond depends upon the compound.
Not all C-H bonds are of the same energy!
See page 305.
Bond energies given in tables are averages based on experimental data.

35. Bond Energies Consider the bond energies on page 306.
Compare the bond energies and bond length associated with single, double, and triple bonds between a given pair of atoms.

36. 8.8 Using Bond Energies
The DH for a reaction can be estimated from bond energies.
DH = energy needed to break bonds of reactants –
energy released when product bonds form

37. 8.8 Using Bond Energies
Estimate the DH for.

38. CH 8 Part II: Bonding Models
Introduction to models (8.7)
Localized Electron (LE) Bonding Model (8.9)
Lewis Structure (8.10)
Resonance (8.12)
Exceptions to the Octet Rule(8.11)
VSEPR Theory (8.13)
Key pages: 326/27

39. 8.7 Models
Read “Fundamental Properties of Models” on page 350.

40. 8.9 LE Bonding Model Localized electron bonding model
Assumes a molecule is made of atoms bound together by sharing pairs of electrons using the orbitals of the bonding atoms.

41. 8.9 LE Bonding Model Localized electron bonding model
Shared electrons are pictures to be localized in the space between the atoms
Called bonding pairs
Non-bonding valence electrons are pictured to be localized on the parent atom.
Called lone pairs
Consider HCl

42. 8.10 Lewis Structures
Lewis structures show the arrangement of the valence electrons in molecules (and ions).
Representative atoms will have the same number of valence electrons as one of the noble gases
2 electrons to be like H
8 electrons to be like all other noble gases

43. Lewis Structures Lewis structures illustrate LE bonding model.
Show the bonding electrons and the lone pairs.
Lewis structures can be used to predict the 3D geometry of a molecule.
Requires application of VSEPR Theory
More to come on this…..

44. 1st Goal: To Write Lewis Structures
Sum the valence electrons.
Use a pair of electrons to form a bond between each of the bonded atoms.
Put the atom that needs the most electrons in the center when the molecule contains more than 2 atoms.
Arrange the remaining electrons to satisfy the duet rule for H and the octet rule for elements in the 2nd row of elements.

45. Writing Lewis Structures
Practice!
H2O O2
HCN
NO31-
PH3

46. 8.12 Resonance
More than one valid Lewis structure can often be drawn for molecules with multiple bonds (double, triple..)
Consider NO21-
2 valid Lewis structures can be drawn.

47. Resonance Structures
Lewis structure just drawn indicate 2 types of bonds in NO single bond and a double bond
However….the data shows that both bonds in NO21- are of the same energy and bond length
Both bonds are stronger and shorter than a single bond, but not as strong or short as a double bond!

48. Exceptions to the Octet Rule
Less than an octet.
Be and B
More than an octet
3rd period elements and “up”
Odd number of electrons

49. Exceptions to the Octet Rule
Less than an octet.
Be - satisfied/stable with 4 electrons
B - satisfied/stable with 6 electrons

50. Exceptions to the Octet Rule
2. More than an octet
Atoms in the 3rd period and “up*” can use their unfilled d orbitals to accommodate more than 8 electrons
Commonly see 10 electrons and 12 electrons around the central atom.
“Up” refers to periods 4, 5,6,…

51. More than an octet
ICl3
PF5

52. Exceptions to the Octet Rule
Odd number of electrons
A small number of molecules have an odd number of electrons
Called “free radicals”
Molecules are highly reactive/unstable
“steal” an electron from other molecules
Example: NO

54. VSEPR Theory Valence Shell Electron Pair Repulsion Theory
Structure around a given atom is determined by minimizing e-pair repulsions
Atoms arrange themselves in 3D space in a manner that minimizes electron pair repulsive forces

55. VSEPR Theory
Predictions based on VSEPR theory agree closely with experimental data.
CO2
BF3
SO2

56. Goals – much practice in lab
Applying LE and VSEPR Theory
From Lewis structure to:
Electron pair geometry (electronic geometry, EG)
molecular geometry (MG)
bond angles
Molecular polarity
Hybridization (CH 9)
Bond type – sigma, pi

61. Sigma and Pi Bonds
Sigma bonds – bonding electrons overlap (bond) on the internuclear axis
Pi bonds – unhybridized p orbitals overlap (bond) above and below the inter nuclear axis

63. Single bond = 1 sigma bond
Double bond = 1 sigma bond, 1 pi bond
Triple bond = 1 sigma bond, 2 pi bonds

64. O=C=O

65. Count # electron densities around the central atom
Write Lewis structure.
Count # electron densities around the central atom
Determines electronic geometry (EG)
Also determines hybridization of central atom
Count # lone pairs around central atom.
This combined with #2 determines molecular geometry (MG)
Determine polarity of each bond.
From this determine polarity of molecule.
Other….Label sigma and pi bonds, identify orbitals involved.

66. Lab – when polarities differ