1. AP Chemistry Ch 8 and 9 Notes
Why Bond?
Atoms bond together to reduce their total potential energy and to increase their stability
A Chemical Bond is a mutual attraction between the nucleus of one atom and the electrons of another atom.
In general atoms bond such that they have an octet of valence electrons. (Hydrogen bonds to get a duet of valence electrons)

2. Ionic vs Covalent
An Ionic Bond is an electrostatic attraction between an atom that has gained electrons (anion) and an atom that has lost electrons (cation). Results when a metal bonds with a nonmetal.
A Covalent Bond results from the sharing of valence electrons between two atoms. Most often occurs between two nonmetals.

3. Ionic Bond Strength
Bond energy is the amount of energy required to break a bond. The greater the bond energy the stronger the bond.
In Ionic Bonds we can use Coulomb’s Law to calculate the bond energy.
E = (2.31 x 10-19J*nm)(Q1*Q2) r
Where E = energy of attraction
Q1 and Q2 are ion charges
r is the distance between ion centers in nm.
note: a negative E indicates an attractive force. (ions together are more stable than separated)

4. Try These: Which has a greater bond energy? NaCl or KBr MgO or Na2O
AlCl3 or AlF3

5. Covalent Bond Length
Bond length is defined as the distance between two atoms when the atoms have the lowest potential energy.
If the atoms are too far apart, they have no interaction and have high PE (low stability)
If the atoms are too close together their nuclei will repel each other and have high PE
If the atoms are at optimal distance, their nuclei will attract the electrons of the other atom and stability will increase.

6. Nonpolar v Polar Covalent Bonds
A Nonpolar covalent bond results when electrons are shared EQUALLY between two atoms (electronegativity is the same)
A Polar Covalent bond results when electrons are shared UNEQUALLY between two atoms. The atom will the higher electronegativity will have a partial negative charge while the atom with the lower electronegativity will have a partial positive charge.
d+ H - Cl d-

7. Electronegativity and Bond Type
The larger the electronegativity difference between two atoms, the greater the ionic character of the bond.
The smaller the electronegativity difference, the greater the covalent character of the bond.

8. Lattice Energy of Ionic Compounds
Def: The energy change that occurs when separate gaseous ions are packed together to form an ionic solid.
We use a variation of Coulomb’s Law to calculate the lattice energy:
lattice energy = k (Q1Q2) r
k is a proportionality constant that varies depending upon the structure of the solid and the electron configurations of the ions.)

9. Uses of Lattice Energy
Because energy is a State Function - a property that is independent of its pathway - we can use lattice energy as one of the steps in an enthalpy of formation calculation for ionic compounds.
Enthalpy of formation is defined as the energy change when one mole of a compound is formed from its elements in the standard state.

10. Example: Find the enthalpy of formation of magnesium fluoride.
Mg (s) + F2 (g) MgF2(s)
Think about what must happen for this reaction to occur…
Solid magnesium must become a gas and then lose two electrons
The fluorine molecule must break apart and then each atom must gain an electron.
The gaseous fluoride ion and magnesium ions must come together to form the product

11. Info you need for this example:
Lattice energy of MgF2 = kJ/mol
First Ionization energy of Mg = 735 kJ/mol
Second Ionization energy of Mg = 1445 kJ/mol
Electron Affinity for F = -328 kJ/mol
Bond Energy for F2 = 154 kJ/mol
Enthalpy of Sublimation of Mg = 150.kJ/mol

12. 1. Sublimation of Mg
Mg (s)  Mg (g)
2. Ionization of Mg
Mg (g)  Mg+ (g) + e-
Mg+ (g)  Mg+2 (g) + e-
3. Break F-F bond
F2 (g)  2F (g)
4. Ionize F atoms
F (g) + e-  F- (g)
5. Bring ions together
Mg+2 (g) + 2 F- (g)  MgF2 (s)
= 150. kJ
=735 kJ
=1445 kJ
=154 kJ
= -328 kJ x 2
= kJ
DH = kJ
Overall Mg (s) + F2 (g)  MgF2

13. Lewis Dot Diagrams
Lewis Dot Diagrams show how atoms in molecules are bonded together.
A dot represents a valence electron in Lewis Dot Diagrams O H H

14. Steps to Lewis Diagrams
Determine the type and number of each atom in the molecule.
Find the total number of valence electrons in the molecule.
Arrange the skeleton structure of the molecule. (If C is present, it goes in the middle. Otherwise put the least electronegative in the center. H is always on an edge.)
Put one pair of electrons between each adjacent atom.

15. Lewis Diagrams (cont)
Add lone (unshared) pairs of electrons on the outer edges of other atoms until an octet is reached around each atom (H needs only a duet).
Count the total electrons. If it is equal to the number you had available, you are done! If you have too many electrons in your diagram, you must remove lone pairs and make double or triple bonds until the number of electrons in the diagram is equal to the number available.

16. Lewis Diagrams (cont)
If there are too many electrons even if you have all single bonds, this means that the octet rule will not be followed. You will put any extra electrons on the central atom.

17. Resonance
Results when electrons can be arranged in more than one valid way within a Lewis Structure.
The actual molecule is an average of all the valid Lewis Structures.

18. Formal Charge
When more than one nonequivalent Lewis structure can be used to represent a molecule, we use formal charge to determine which structure is most appropriate.
The structure resulting in the formal charges closest to zero for the atoms is most stable.
The sum of the formal charges on a molecule or ion must equal the charge on that species.
Formal charge =
# VE in free atom – [# lone pair electrons + ½ # shared electrons]

19. Formal charge =
# VE in free atom – [# lone pair electrons + ½ # shared electrons]
H N C
H C N
The bottom structure is more stable.

20. VSEPR Valence Shell Electron Pair Repulsion Theory
Predicts the geometries of molecules
Based on the idea that pairs of valence electrons around an atom repel each other and will be positioned as far apart as possible.
To determine the shape of a molecule, you must first draw the Lewis Diagram

21. 2 Regions of High Electron Density: Linear
Only TWO regions of high electron density and two bonded atoms around the central atom (AB2) OR
Only TWO atoms (AB)
A B
B A B
Bond angle is 180º

22. 3 Regions of High Electron Density: Trigonal Planar
THREE regions of high electron density with three bonded atoms around the central atom (AB3)
B A B B
Bond angles are 120º

23. 4 Regions of High Electron Density: Tetrahedral
FOUR regions of high electron density with four atoms bonded to the central atom. (AB4) B
B A B
Bond angles are 109.5º

24. Trigonal Pyramidal B B A B
FOUR regions of high electron density with only THREE bonded atoms. AB3E B
B A B
Bond angles of ~107º

25. Bent
THREE or FOUR regions of high electron density, but only TWO atoms bonded to the central atom. (lone pairs present)
AB2E AB2E2
B A B
B A B
Bond angle of ~120º
Bond angle of ~104.5º

26. 5 Regions of High Electron Density: Trigonal Bipyramidal
FIVE regions of high electron density with five bonded atoms (AB5) B
B A B
B B
Bond angles of 120º and 90º

27. More 5 Regions of High Electron Density
General Formula
Shape Name
AB4E
See-saw
AB3E2
T-Shaped
AB2E3
Linear

28. 6 Regions of High Electron Density: Octahedral
SIX regions of high electron density with six bonded atoms (AB6)
B B
B A B
Bond angles of 90º

29. More 6 Regions of High Electron Density
General Formula
Shape Name
AB5E
Square Pyramidal
AB4E2
Square Planar
AB3E3
T-Shaped
AB2E4
Linear

30. Hybridization Basics
When an atom bonds, the atomic orbitals of the valence shell will change to form a new set of orbitals called hybrid orbitals.
Within a molecule, the hybrid orbitals have lower average energy than the atomic orbitals do.
Hybridization makes sense because in a molecule like CH4, all four bonds are equivalent. If bonding electrons were still in s and p orbitals, we would expect to have two different types of bonds in CH4.

31. sp3 Hybridization
Occurs when there are 4 regions of high electron density around the atom.
1 s-orbital + 3 p-orbitals = 4 sp3 orbitals
Orbitals are arranged tetrahedrally

32. sp2 Hybridization
Occurs when there are 3 regions of high electron density around the atom.
1 s-orbital + 2 p-orbitals = 3 sp2 orbitals
Orbitals are arranged in trigonal planar geometry

33. sp Hybridization
Occurs when there are 2 regions of high electron density around the atom.
1 s-orbital + 1 p-orbitals = 2 sp orbitals
Orbitals are arranged in linear geometry

34. dsp3 and d2sp3 hybridization
If an atom has more than 4 regions of high electron density, a d-orbital can hybridize.
These atoms must be in at least the 3rd period on the periodic table (the first and second energy level do not have any d-orbitals)
dsp3 hybridization occurs when there are 5 regions of high electron density. (trigonal bipyramidal arrangement)
d2sp3 hybridization occurs when there are 6 regions of high electron density. (octahedral arrangement)

35. Sigma (s) Bonds
Sigma bonds result from the end-to-end overlap of orbitals.
Electron density is concentrated between the two nuclei

36. Pi (p) Bonds Results from a sideways overlap of orbitals.
Electron density is concentrated above and below the line between the nuclei.

37. s and p bonds in ethylene
Draw the dot diagram for C2H4
What is the hybridization of the C atoms?
sp2
The hybrid orbitals of the C atoms overlap with each other and with H orbitals to create s bonds.

38. The unhybridized p-orbitals (in blue) from the C atoms are then available to overlap in a p-bond.

39. s and p of multiple bonds
A single bond consists of only a s bond
A double bond consists of one s and one p bond
A triple bond consists of one s and two p bonds

40. Molecular Orbital Theory
Localized electron models such as Lewis dot structures and hybridization have some limitations because it incorrectly assumes that the electrons reside in orbitals of individual atoms.
In reality, electrons reside in orbitals that are characteristic of molecules as a whole.

41. Bonding v AntiBonding Molecular Orbitals
A bonding molecular orbital has lower energy (greater stability) than the atomic orbitals from which it is formed.
An antibonding molecular orbital has higher energy (less stability) than the atomic orbitals from which it is formed.

42. Placing electrons in a bonding MO results in a stable bond.
In Antibonding MOs electron density is not between the nuclei.
In Bonding MOs electron density is between the nuclei. (Negative “glue” holding the nuclei together.)
Placing electrons in a bonding MO results in a stable bond.
Placing electrons in an antibonding MO results in an unstable bond.

43. Sigma vs Pi MOs
In a sigma (s) MO, the electron density is concentrated symmetrically along a line between the two nuclei.
In a pi (p) MO, the electron density is concentrated above and below the line between the nuclei.
Sigma MOs

44. Pi Bonding and Antibonding Orbitals
AP Chemistry Ch 8 and 9 Notes
Pi Bonding and Antibonding Orbitals
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45. Molecular Orbital Configurations
The number of molecular orbitals formed is equal to the number of atomic orbitals combined
The more stable the bonding MO, the less stable its corresponding antibonding MO
Electrons fill MOs from lowest energy to highest energy
Each MO can accommodate up to 2 electrons with opposite spin
Hund’s Rule applies

46. In H2, the electrons from H’s atomic orbitals are placed in the lowest available molecular orbitals as shown in the diagram. This forms a stable bond.
In He2, the electrons from He’s atomic orbitals are placed in the lowest available molecular orbitals as shown in the diagram. This forms a unstable bond (equal # of electrons in s and s* orbitals).

47. For the second period homonuclear diatomic elements Li2, B2, C2, and N2, the diagram above shows the filling order.

48. For all other homonuclear diatomic elements, this energy diagram shows the correct filling order.
Why the difference? It turns out that only when the bond lengths are relatively short can the two p-orbitals on the bonded atoms efficiently overlap to form a strong p bond. For the larger atoms, the p bond is weaker than the s bond.

49. Bond Order
Bond order indicates the relative strength of a bond. The higher the bond order, the greater the strength.
Bond order =
½ (#electrons in bonding MOs - #electrons in antibonding MOs)
Calculate the bond order in F2, O2, and N2.
Notice that the bond order is equal to the number of electrons being shared between the atoms in the Lewis Diagram.

50. Diamagnetism v Paramagnetism
Diamagnetic substances are (very slightly) repelled by a magnet. All electrons are paired.
Paramagnetic substances are attracted by a magnet. The atom/molecule contains unpaired electrons.