1. Chemical Bonds

2. Bonds
Chemical bonds are the interactions between nuclei and electrons which hold molecules together. All bonds are essentially the same. They are electrons (always in pairs) that are being attracted to positive nuclei from competing atoms.

3. Ionic Bonds
Formed from electrostatic attractions of closely packed, oppositely charged ions.
Formed when an atom that loses electrons reacts with one that has a high electron affinity which means it wants to gain an electron.
The electron is still being held by both nuclei, though one might be holding onto them tighter, depending on their relative electronegativities.

4. Covalent Bonds
Covalent bonds are often described as a sharing of electrons (between two nuclei). In a bond such as H2 or O2 , the electrons are held exactly in the middle between the two nuclei and will form a 100 % covalent bond due to this exactly equal sharing. This is often called a non-polar covalent bond.

5. Energy
The amount of Energy in these interactions can be calculated by a formula called Coulomb’s Law. E = ( 2.31 x J nm) ( Q1 ) ( Q2 ) / r Q is the ionic charge (Eg. Na+1 is + 1) So we would expect MgO to have a higher energy than NaF, due to their +2 charge r is the distance in nanometers Remember: Lower Energy = Greater Stability

6. Bond Length and Energy This diagram is in the text

7. Bond Length and Energy
Notice the far right of the chart. The atoms are too far away to have any interaction so their Energy is Zero.

8. Bond Length and Energy
As we move to the left, the energy starts to decrease (become lower and therefore more stable). This continues until the two atoms find a distance of lowest Energy. This is where they are the most stable. The distance between the two atoms at this point is called the bond length.

9. Bond Length and Energy
You can see in the diagram what the bond length for the H – H bond is >> .04 nm

10. Bond Length and Energy
When we try to move the atoms closer than their lowest energy distance, the energy increases tremendously. The force of interaction between the two positively charged nuclei becomes very high.

11. Bond Length and Energy
The most stable situation for the H2 molecule is when it has lowest energy and that occurs when the two atoms are at a distance apart that allows that to happen. This is called the bond length. The bond length can be measured in meters, nanometers, centimeters, and Angstroms.

12. Electronegativity
Linus Pauling developed the idea of Electronegativity. It describes the amount of strength that each atom has to pull on shared electrons in a bond. Fluorine has the highest value of 4.0 It has a symbol of En

13. Electronegativity
Here is a periodic chart with En values on it. You should be familiar with the trends of En

14. Electronegativity
To determine characteristics of a chemical bond, we sometimes have to calculate the difference in Electronegativity values between the two atoms in the bond. When we do this we always subtract the lower value from the higher value. We always want a positive value. Eg. It would be 4.0 – 1.5 NOT 1.5 – 4.0

15. Electronegativity
So in CaO, the electronegativity difference ( Δ En ) would be 3.5 – 1.0 = 2.5 If the Δ En is < or = to .4 we consider it to be a non-polar covalent bond. If the Δ En is between .4 and 1.67 we consider it to be a polar bond. If the Δ En is > or = to 1.67 we consider it to be an ionic bond.

16. Dipole Moment
When there is a difference in electronegativities between the two atoms in a bond, it creates a dipole moment. It is a description of the charge difference in the bond. We use the symbol δ to show this partial charge. So in NaCl >>> δ+ Na – Cl δ-

17. Polarity
A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.
The Fluorine is negative because it has a higher electronegativity.

18. Ionic Size
When a species becomes an ion, there are changes in its size and volume. Cations are getting smaller than the atoms since they lose electrons and have 1 less shell. Anions are getting larger than the atoms since they are gaining electrons and have a lower proton to electron ratio.

19. Lattice Energy
When we are finding the heat of reaction for the creation of an ionic compound, we need to know about Lattice Energy. Ionic Compounds form when gaseous ions combine. (One negative with one positive). This combining is called the Lattice Energy.

20. Lattice Energy
Lattice Energy is only one of the changes that need to be taken into account. The other changes depend on the problem These problems will resemble the work we did with Hess’s Law.

21. Lattice Energy Using LiF as an example Li (s) + ½ F2 (g)  LiF (s)
We want to create one mole of LiF from its elements.
First we have to get both species into gaseous ions.

22. Lattice Energy
We will start with the Lithium species. Equation Process Li (s)  Li (g) Hsublimation Li (g)  Li+1 (g) 1st ionization Energy

23. Lattice Energy
Now we will do the Fluorine species Equation Process ½ F2 (g)  F (g) Bond Energy F (g)  F-1 (g) Electron Affinity Li+1 (g) + F-1 (g)  LiF (s) Lattice Energy

24. Lattice Energy

25. Lattice Energy
From the previous slide, you can see that the Lattice Energy is – 1,047 kJ and is one of the components in finding the overall Heat of Reaction (which is – 617 kJ). The negative sign tells us that the formation of the compound is more stable than the ions.

26. Summary of Steps
Using LiF as an example Li (s) + ½ F2 (g)  LiF (s) Li (s)  Li (g) Hsublimation Li (g)  Li+1 (g) 1st ionization Energy ½ F2 (g)  F (g) Bond Energy F (g)  F-1 (g) Electron Affinity Li+1 (g) + F-1 (g)  LiF Lattice Energy

27. Remember Hess’s Law ?
Li (s)  Li (g) 161 Li (g)  Li+1 (g) + e ½ F2 (g)  F (g) 77 e-1 + F (g)  F-1 (g) Li+1 (g) + F-1 (g)  LiF (s) Overall Reaction: Li (s) + ½ F2  LiF (s)

28. Lattice Energy
Another look at the diagram

29. Remember Hess’s Law ?
We add up the individual heat changes for each equation just like we did for Hess’s Law. Doing this gives us: Δ Hrxn = kJ per mole of LiF Li (s) + ½ F2 (g)  LiF (s)

30. Percent Ionic Character
We can determine the percent Ionic Character from a bond from its dipole moment. % ionic character = (measured dp / calculated dp ) x 100 % We need to know this more on a qualitative level, however there are a few values that are important.

31. Percent Ionic Character
If we have a species like O2, the electrons are being pulled equally by two identical nuclei. Therefore we can say that they are being equally shared and 100 % covalent The next slide has a reminder from an earlier slide.

32. Percent Ionic Character
If the Δ En is < or = to .4 we consider it to be a non-polar covalent bond. The percent ionic character is about 4 % when the Δ En is .4 If the Δ En is between .4 and 1.67 we consider it to be a polar bond. The percent ionic character at Δ En of 1.67 is 50 % If the Δ En is > or = to 1.67 we consider it to be an ionic bond. There can’t be a 100 % ionic bond because that would mean a complete transfer of the electron and therefore there wouldn’t be a bond.

33. Bond Energy It is the energy required to break a bond.
Therefore according to the Law of Conservation of Energy, it is also the amount of Energy released when a bond is formed.
Bond energies are measured in kJ / mole

34. Bond Energy
Multiple bonds will have a greater bond energy than single bonds. Usually, the shorter the bond, the stronger the bond. We can calculate Δ Hrxn using bond energies.

35. Bond Energy
In a Chemical Reaction, the reactants must break the bonds and then the material must reform to make new materials (products) which needs new bonds to form.
The procedure is:
Δ H = (Σ n * (bonds broken)) - (Σ n * (bonds formed))
In other words: Bonds broken – Bonds formed
reactant side product side

36. Bond Energy - Example C2H6 + O2  H2O + CO2
Calculate the Δ Hrxn for the following reaction
C2H6 + O2  H2O + CO2
Remember that it is crucial to understand the procedure needed to solve a problem.
The first step is to balance the equation.
2 C2H6 + 7 O2  6 H2O + 4 CO2

37. Bond Energy - Example
Here is the balanced equation written structurally. 2 C2H6 + 7 O2  6 H2O + 4 CO2 O = O H – O – H O = C = O O = O H – O – H O = C = O O = O H – O – H O = C = O O = O H – O – H O = C = O O = O H – O – H O = O H – O – H O = O

38. Bond Energy - Example
The third step is to reference the bond energies from your table. With the table, calculate the total bond energies of each species. Reactants: C2H6  5,650 kJ O2  3,465 kJ Products: H2O  5,604 kJ CO2  6,392 kJ

39. Bond Energy - Example
Make sure that you counted the bonds accurately and multiplied by the moles in the equation. The next step is to calculate the total amount of bonds broken (reactants) and the total amount of bonds formed (products) Bonds broken: 5, ,465 = 9,115 kJ Bonds formed: 5, ,392 = 11,996 kJ

40. Bond Energy - Example
Since it is Bonds Broken – Bonds Formed the final work will look like this: 9,115 – 11,996 = - 2,881 kJ Just like before, the negative sign tells us that it is an exothermic reaction.

41. O-2 > F-1 > Na+1 > Mg+2 > Al+3
Isoelectronic Ions
Ions containing the same number of electrons
(O-2 , F-1, Na+1, Mg+2, Al+3)
O-2 > F-1 > Na+1 > Mg+2 > Al+3
largest smallest
The primary reason for the difference in size is proton to electron ratio.
In Oxygen it is 8:10 In Aluminum it is 13:10

42. Models
Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.
A model’s main goal is to give the observer a mental image of a concept.

43. Basic Properties of Models
A model does not equal reality.
Models are oversimplifications, and are therefore often wrong.
Models become more complicated as they age.
We must understand the underlying assumptions in a model so that we don’t misuse it.

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.
Electrons that are part of a bond are called bonding pairs.
Electrons that are solely around one atom are called either non-bonding electrons or lone pairs.

45. Localized Electron Model
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.

46. 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.

47. Drawing Molecular Shapes Some Rules to Follow
1. The atom by itself is the central atom 2. Carbon is always the central atom and has 4 bonds ( CO is a notable exception) 3. Hydrogen is always on the outside with a single bond. 4. The Group VII elements always have a single bond. 5. Outside atoms always will have a full octet

48. Drawing Molecular Shapes Other Useful Facts
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. These are called electron deficient.
Certain 3rd row and heavier elements can exceed the octet rule using empty valence ‘d’ orbitals. These are called expanded octets.

49. VSEPR Model
The structure around a given atom is determined by minimizing electron pair repulsions.
In other words, the atoms around the central atom want to be as far apart as possible.

50. Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Take note of the central atom and which pairs of electrons are bonding pairs and which are lone pairs.
3. Refer to the sheet and determine the shape of the molecule depending on the number of lone pairs and bonding pairs around the central atoms.

51. Molecular Shapes
Effective Pairs: The number of combined bonding pairs and lone pairs. Bonding Pairs: Electron Pairs that are involved in bonding of atoms. Lone Pairs (also called Non-bonding pairs): Electron Pairs that are found around a single atom and don’t take part in bonding.

52. Overview of VSEPR
Here is a PPT inside of a PPT that gives a good overview of VSEPR shapes. Notice how the shapes form and that there is a relationship between the number of effective pairs around the Central atom and its shape.

53. Linear
This short animation gives you a look at the shape of a molecule with two effective pairs

54. Trigonal Pyramidal
This animation shows the basic Trigonal Pyramidal Shape along with bond angles.

55. Tetrahedral

56. Trigonal Bipyramidal
There is no animation for this one but the basic shape has 5 bonding pairs. Other shapes are See Saw and T- Shape.

57. Octahedral

58. Formal Charge
The difference between the number of valence electrons on the free atom and the number assigned to the atom in the molecule.
We need:
# Valence Electrons on the atom before bonding
# VE “belonging” to the atom in the
molecule after bonding

59. Formal Charge
With Formal Charges the more stable structure is the one with the lowest values.
We have an example of three possible structures of CO2 below. Notice the one that is the actual structure is the one with the lowest possible Formal Charges.
Actual

60. Resonance Structures
Occurs when more than one valid Lewis structure can be written for a molecule.
These are resonance structures. The actual structure is an average of the resonance structures. See how the double bonds are different in the two structures.

61. Resonance Structures
Probably the easiest way to determine if a structure will require a resonance drawing is if the multiple bond that is needed, seems to be able to be placed in multiple places, it is a resonance structure.