1. Bonding: General Concepts
Dr. Walker
DE Chemistry

2. Bonding Bond Energy – Energy required to break a chemical bond
Atoms bond in a way that achieves the lowest possible energy for the system
Bond length – distance between the nuclei of two bonded atoms
Shorter bond length = stronger bond = higher bond energy!!

3. Ionic Bonding Electrons are transferred Metals react with nonmetals
Ions paired together have lower energy (greater stability) than separated ions
Ionic compounds conduct electricity in a molten state

4. Coulomb’s Law Shows energy of interaction between a pair of ions
E = Energy, r = distance between ionic nuclei, Q = ionic charge
Lower energy = stronger attraction, more favorable interaction

5. Covalent Bonding Electrons are shared by nuclei
Pure covalent (non-polar covalent)
Electrons are shared evenly
Polar covalent bonds
Electrons are shared unequally
Atoms end up with fractional charges
d+ or d- (partial charges)

6. Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself.
General Trend – Increases to the right, Increases going up
on the Periodic Table

7. Electronegativity and Bonding
The greater the electronegativity difference, the greater the chance of a bond being ionic
This has been quantified, but (as usual), there are exceptions.

8. Bond Polarity and Dipole Moments
Dipolar Molecules
Molecules with a somewhat negative end and a somewhat positive end (a dipole moment) due to a difference in electronegativity
Molecules with preferential orientation in an electric field (see next slide)
All diatomic molecules with a polar covalent bond are dipolar

9. Physical Effect Of Dipole Moments

10. Polar Bonds, Nonpolar Molecules
Covalent molecules can have multiple dipole moments, but can be nonpolar overall due to symmetry.

11. Ions: Electron Configurations and Sizes
Ionic bonds
Electrons are transferred until each species attains a noble gas electron configuration (ns2np6)
This trend is primarily for main group elements and still has exceptions (Pb2+, Sn2+ among others)
Formulas are predicted based on atoms required to form a neutral atom
Covalent bonds
Electrons are shared in order to complete the valence configurations of both atoms
Exceptions can occur below period 2
Note: This SHOULD be review. I’m not sure why the book puts it here.

12. Ionic Size Anions are larger than the parent atom
Cations are smaller than the parent atom
Essentially, they “lose” a shell
Ion size increases within a group (more energy levels)
Isoelectronic ions
Ions with the same number of electrons (for example, Cl-, S2-, and P3- all have 18 electrons)
Size decreases as the nuclear charge (atomic number) increases

13. Ionic Size

14. Energy In Ionic Compounds
Lattice Energy
The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid
Energy change is negative (-DH) as it is released to the surroundings

15. Lattice Energy
Modified form of Coulomb’s Law, shows energy “loss” from lattice formation
k = a proportionality constant dependent on the solid structure and the electron configuration
Q1 and Q2 are charges on the ions
r = shortest distance between centers of the cations and the anions
Lattice energy increases as the ionic charge increases and the distance between anions and cations decreases
More charge, closer ions = more energy “savings”

16. Example Of Lattice Energy
The final step of lattice formation is what makes the process exothermic overall

17. Crystal Lattices
Our previous visualization of ionic salts as “stand-alone” compounds is NOT correct.
Ionic salts form lattice structures, shown on the left. This network formation ultimately makes the energy work out in favor of bond creation.
Ionic formulas are actually empirical formulas, since we don’t count the atoms in an entire lattice, which is impossible

18. Bond Character
Differences in electronegativity cause unequal sharing of electrons in covalent bonds (polarity) or complete transfer of electrons (ionic bonds)
Some cases are ambiguous….it’s hard to predict what they are
Bond character – most bonds are not 100% covalent or 100% ionic

19. Bond Character Bond Character measured by the above equation
Ionic vs. Covalent
Ionic compounds generally have greater than 50% ionic character
Ionic compounds generally have electronegativity differences greater than 1.6
Percent ionic character is difficult to calculate for compounds containing polyatomic ions (each ion has covalent bonds)

20. Electronegativity Values

21. Bonding As A Model
Models are a way to represent what we observe in nature, especially in what we can’t see with our eyes
Models do not equal reality and are often wrong
The way covalent bonds are represented are a model…..
….these have proven to be a good representation in most cases, but some molecules require us to think of them as a unit rather than a collection of “bonds”
Keep this in mind as we progress through the chapter

22. Covalent Bonding – A Review
Electron pairs are shared between two atoms in covalent bonding
Single Bond – one pair of electrons being shared
Double Bond – two pairs of electrons being shared
Triple Bond – three pairs of electrons being shared
More electrons being shared = higher bond energy = shorter bond

23. Bond Energies For Covalent Bonds
These bond energies are endothermic – energy added to break the bonds
Conversely, bond formation is exothermic (as previously mentioned)

24. Localized Electron Bonding Model
Lone electron pairs
Electrons localized on an atom (unshared)
Bonding electron pairs
Electrons found in the space between atoms (shared pairs)
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
Lone Pair
Bonding Pair

25. Localized Electron Bonding Model
Derivations of the Localized Model
Valence electron arrangement using Lewis structures
Prediction of molecular geometry using VSEPR (valence shell electron pair repulsion)
Description of the type of atomic orbitals used to share or hold lone pairs of electrons

26. Lewis Structures
Lewis structures show how valence electrons are arranged among atoms in a molecule.
Lewis structures reflect the central idea that stability of a compound relates to noble gas electron configuration (octet rule)
Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - )

27. HONC, HONC..
The HONC Rule
Hydrogen (and Halogens) form one covalent bond
Oxygen (and sulfur) form two covalent bonds
One double bond, or two single bonds
Nitrogen (and phosphorus) form three covalent bonds
One triple bond, or three single bonds, or one double bond and a single bond
Carbon (and silicon) form four covalent bonds.
Two double bonds, or four single bonds, or a triple and a single, or a double and two singles

28. I Gotta Be Me…Exceptions To The Rules
2nd row elements C, N, O, F observe the octet rule (HONC rule as well).
2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.
3rd row (like S and P, especially) 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.

29. Completing a Lewis Structure -CH3Cl
Make the atom wanting the most bonds central
Add up available valence electrons:
C = 4, H = (3)(1), Cl = 7 Total = 14
Join peripheral atoms
to the central atom
with electron pairs. H .. ..
Complete octets on
atoms other than
hydrogen with remaining
electrons H .. C .. Cl .. .. .. H

30. Multiple Covalent Bonds: Double bonds
Ethene
Two pairs of shared electrons

31. Multiple Covalent Bonds: Triple bonds
Ethyne
Three pairs of shared electrons

32. Drawing Lewis Structures
This differs from what you learned In previous class, as I GREATLY simplified it.
Step 1: Count the number of valence electrons in the compound
Step 2: Connect the atoms around the central atom by single bonds
Step 3: Place remaining electrons on the outside atoms to fulfill octet rule
Step 4: Place any remaining electrons on the central atom.
Step 5: If the central atom does not have an octet, share pairs of electrons as bonds from outside atoms.

33. Practice

34. Answers

35. Practice - Exceptions

36. Answers - Exceptions

37. The Problem With Lewis
The nitrate ion shows two types of bonds: 1 double bond and 2 single bonds
Experiments show ONLY ONE BOND TYPE with a bond length BETWEEN that of a single and double bond
This is not the only molecule exhibiting this behavior
Presents a limitation of Lewis structures and the Localized Electron Model

38. Resonance
Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule.
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
Benzene, C6H6

39. More Resonance Ozone (O3)
Bond lengths are identical, between that of a single and double bond
Neither structure is
Technically correct!

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

41. Odd-Electron Molecules
Lewis structure models work best with molecules with even numbers of total electrons
The localized electron model is based on pairs, so it doesn’t work so well with these
Example
NO, 11 total valence electrons

42. Formal Charge
Method of handling ions/molecules with unusual arrangements to help determine proper Lewis structure

43. Assigning Formal Charge
This pertains to a Lewis structure that you must draw first
Number of valence electrons on the free atom minus the number of valence electrons assigned to the atom in the molecule
Lone pair (unshared) electrons belong completely to the atom in question
Shared electrons are divided equally between the sharing atoms
The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species
If the charge on an ion is -2, the sum of the formal charges must be -2

44. Assigning Formal Charge
In a molecule with many possibilities, the representation with the lowest formal charges is most appropriate
Bottom choices are best here with no charge on Xenon higher than +1

45. Formal Charge Practice

46. Formal Charge Examples
A) P = 0, O = 0, Cl = -1
B) S = 0, 2 oxygens = 0, 2 oxygens = -1
C) Cl = 0, 3 oxygens = 0, 1 oxygen = -1
D) P = 0, 1 oxygen = 0, 3 oxygens = -1

47. VSEPR Valence Shell Electron Pair Repulsion
The structure around a given atom is determined principally by minimizing electron-pair repulsions.
Translation: all atoms and electrons want their own personal space and get as far away from everyone else as possible

48. Possible Arrangements
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

49. Linear Structure
No unbonded electrons to deal with, chlorines as far apart as possible

50. Linear
AX2
CO2

51. Trigonal Planar
No unbonded electrons on central atom, molecule stays in a single plane

52. Trigonal Planar
AX3
BF3
AX2E
SnCl2

53. Tetrahedral
When four pairs of electrons are present around an atom, they should always be arranged tetrahedrally.
This can be in bonds OR in electron pairs

54. Tetrahedral
AX4
CCl4
AX3E
PCl3
AX2E2
Cl2O

55. Trigonal Bipyramidal
Involves the arrangement of a compound with five electron pairs surrounding a central atom
Notice the term “electron pairs” instead of “atoms”

56. Trigonal Bi-pyramidal
AX5
PCl5
AX4E
SF4
AX3E2
ClF3
AX2E3
I3-

57. Trigonal Bipyramidal Arrangements

58. Octahedral
Six electron pairs arranged around a given atom with 90-degree angles

59. Octahedral
AX6
SF6
AX5E
BrF5
AX4E2
ICl4-

60. Electron Pairs
Lone (unshared) electron pairs require more room than bonding pairs (they have greater repulsive forces) and tend to compress the angles between bonding pairs