1. BONDING
SOL Review

2. Introduction to Bonding

3. Bonding Let’s consider the compound Cesium Fluoride, CsF.
The electro-negativity value (EV) for Cs is 0.70; the EV for F is 4.00.
The difference between the two is 3.30, which falls within the scale of ionic character.
When the electro-negativity difference between two atoms is greater than 1.7 the bond is mostly ionic.

5. Ionic Bonding
Ionic Bonding

6. Ionic Bonding In an Ionic bond:
The electro-negativity difference is extreme,
So the atom with the stronger pull doesn’t really share the electron
Instead the electron is essentially transferred from the atom with the least attraction to the atom with the most attraction

7. Ionic Bonding
When a metal bonds with a nonmetal an: Ionic bond is formed
An ionic bond contains a positive and negative ion.
A positive ion is called a cation.
A negative ion is called an anion.
An Ionic bonding always involves the transfer of an electron from the metal to the nonmetal.
The cation and anion are held together by electrostatic attraction.

8. Characteristics of Ionic Compounds
Ionic compounds do not consist of individual molecules. Instead there is a huge network of positive and negative ions that are packed together in a solid brittle crystal lattice.
Because their bonds are strong, ionic compounds tend to have very high melting and boiling points
-Ionic compounds are electrolytes, which means they can conduct electricity
When forming ionic compounds the positive and negative charges must balance
Ionic crystals cannot conduct electricity because the ions must be able to move.

9. Sea of Electrons

10. nonmetal w/nonmetal = covalent
Bonding
The take home lesson on electro-negativity and bonding is this:
The closer together the atoms are on the P.T., the more evenly their e- interact, and are therefore more likely to form a covalent bond
The farther apart they are on the P.T., the less evenly their e- interact, and are therefore more likely to form an ionic bond.
metal w/nonmetal = ionic
nonmetal w/nonmetal = covalent

12. Covalent Bonding

13. Covalent Bonding In a covalent bond:
The electro-negativity difference between the atoms involved is not extreme
So the interaction between the involved electrons is more like a sharing relationship
It may not be an equal sharing relationship, but at least the electrons are being “shared”.

14. Covalent Bonding Covalent Bonding is between two or more non-metals.
Covalent bonds are formed when electrons are shared between two atoms.
If they share 2 electrons, the form a single bond; 4 electrons is a double bond;
If two atoms share 6 electrons, they form a triple bond.

15. Covalent Bonding
Polar bonds usually involve nitrogen, oxygen or fluorine (NOF)
Non-Polar bonds usually involve carbon-hydrogen bonds
In polar bonds, the electrons are shared unequally
In non-polar bonds, the electrons are shared equally.
Covalent compounds can exist in any state (solid, liquid or gas).  They have low melting and boiling points.

16. Lets look at the molecule Cl2
Covalent Bonds
Lets look at the molecule Cl2 Cl Cl Cl +
Shared
Electrons

17. Notice 8 e- in each valence shell!!!
each atom must have 8 valence e's Cl Cl
Notice 8 e- in each valence shell!!!
Shared electrons are counted with both atoms Cl 2

18. How about the molecule HCl? (Polar Covalent) shared, but not evenly
Covalent Bonds
How about the molecule HCl? H Cl Cl H +
2.1
3.0
(Polar Covalent) shared, but not evenly

19. So what’s the bottom line?
To be stable the two atoms involved in the covalent bond share their electrons in order to achieve the arrangement of a Noble Gas.

20. Drawing Lewis Structures
The Octet Rule
Concept of Formal Charge

21. Lewis Structures
Lewis structures are representations of molecules showing all valence electrons, bonding and nonbonding.

22. The Octet Rule
The most important requirement for the formation of a stable compound is that the atoms achieve a noble gas electron configuration.

23. Writing Lewis Structures
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
PCl3
(7) = 26

24. Writing Lewis Structures
The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.
Keep track of the electrons:
= 20

25. Writing Lewis Structures
Fill the octets of the outer atoms.
Keep track of the electrons:
= 20; = 2

26. Writing Lewis Structures
Fill the octet of the central atom.
Keep track of the electrons:
= 20; = 2; = 0

27. Writing Lewis Structures
If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does. Usually occurs with carbon, nitrogen, oxygen

28. More than One Lewis Structure
Resonance Structures
More than One Lewis Structure

29. Resonance
This is the Lewis structure we would draw for ozone, O3. -

30. Resonance
But this is at odds with the true, observed structure of ozone, in which…
…both O-O bonds are the same length.
…both outer oxygens have a charge of -1/2.

31. Resonance
One Lewis structure cannot accurately depict a molecule like ozone.
We use multiple structures, resonance structures, to describe the molecule.

32. Resonance Just as green is a synthesis of blue and yellow…
…ozone is a synthesis of these two resonance structures.

33. Resonance
In truth, the electrons that form the second C-O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.
They are not localized; they are delocalized.

34. Exceptions to the Octet Rule

35. Exceptions to the Octet Rule
There are three types of ions or molecules that do not follow the octet rule:
Ions or molecules with an odd number of electrons
Ions or molecules with less than an octet
Ions or molecules with more than eight valence electrons (an expanded octet)

36. Odd Number of Electrons
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

37. Fewer Than Eight Electrons
Consider BF3:
Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.
This would not be an accurate picture of the distribution of electrons in BF3.

38. Fewer Than Eight Electrons
Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

39. Fewer Than Eight Electrons
The lesson is: if filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

40. More Than Eight Electrons
The only way PCl5 can exist is if phosphorus has 10 electrons around it.
It is allowed to expand the octet of atoms on the 3rd row or below.
Presumably d orbitals in these atoms participate in bonding.

41. More Than Eight Electrons
Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

42. More Than Eight Electrons
This eliminates the charge on the phosphorus and the charge on one of the oxygens.
The lesson is: when the central atom in on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

43. Hydrate Formation
In the construction of a crystal lattice, depending on the ions involved there can be small “pores” develop between ions in the ionic crystal.
Some ionic compnds have enough space between the ions that water molecules can get trapped in between the ions
Ionic compounds that absorb water into their pores form a special type of ionic compound called a hydrate.

44. Trapped Water
Molecules
Hydrated Crystal

45. Hydrate Formation
Hydrates typically have different properties than their dry versions - A.K.A. anhydrides
Anhydrous CuSO4 is nearly colorless
CuSO4•5 H2O is a bright blue color
When Copper (II) Sulfate is fully hydrated there are 5 water molecules present for every Copper ion.
The hydrated name would be Copper (II) Sulfate Pentahydrate

46. Hydrate Formation
Have you ever bought a new purse or camera and found a small packet of crystals labeled – do not eat?
These crystals are there to absorb water that might lead to mildew or mold
The formula of a hydrate is XAYB • Z H2O (Z is a coefficient indicating how many waters are present per formula unit)

47. Molecular Polarity All above are NOT polar Molecules will be polar if
a) bonds are polar AND
b) the molecule is NOT “symmetric”
All above are NOT polar