1. CHEMICAL BONDING Cocaine Chemistry I – Chapter 8
To play the movies and simulations included, view the presentation in Slide Show Mode.

2. What is a Chemical Bond?
Chemical bonds involve ONLY the valence electrons of atoms.
The nucleus is unaffected, so each atom keeps its identity (carbon stays carbon)
Chemical bonds seek stability in the octet rule – Most atoms want to have 8 valence electrons.
3 ways to achieve this:
Transferring electrons – Ionic bonds
Sharing electrons – Covalent bonds
Floating electrons – Metallic bonds

3. Types of Chemical Bonds
Ionic Bond: Chemical bonding that results from the bonding between metals and nonmetals.
Electrons are transferred from one atom to another, from the metal (electropositive atom) to the nonmetal (electronegative atom.)

4. Atoms that lose valence electrons become positively charged and are known as CATIONS.
Atoms that gain valence electrons become negatively charged and are known as ANIONS.
Cations and Anions come together to form a NEUTRAL 3D crystal

6. Properties of Ionic Compounds
Are hard, rigid, brittle crystalline solids which may or may not dissolve in water.
Have high melting and boiling points.
Good conductors of electricity when melted or when dissolved (aqueous solution); nonconductors when solid.

7. Metallic Bonds
Two metals do not transfer electrons to other metals nor do they share electrons.
Their valence electrons are not held to any specific atom in the solid metal and can move freely from one metal atom.
Think Hot Potato! (Electron Sea Model).

9. Properties of Metals
Because these delocalized electrons are free to move, metals are excellent conductors of electricity and heat.
They have a luster and are malleable, ductile, and usually durable.
Melting points are variable but boiling points are all high.

10. Covalent bond: Chemical bonding that results from the sharing of electrons between two nonmetals
Covalent bonds equally (NONPOLAR) or unequally (POLAR) share the electrons between the atoms.
A group of covalently bonded atoms is called a molecule

11. Properties of Covalent Compounds
Because covalent bonds are generally weaker than ionic bonds, the melting points and boiling points of covalent compounds are generally lower than those of ionic compounds.
Many covalent compounds exist as gases at room temperature.

12. Review of Chemical Bonds
There are 3 forms of bonding:
_________—complete transfer of 1 or more electrons from one atom to another (one loses, the other gains) forming oppositely charged ions that attract one another
_________—some valence electrons shared between atoms
_________ – holds atoms of a metal together
Most bonds are somewhere in between ionic and covalent.

13. Bond and Lone Pairs
In Covalent Bonds valence electrons are distributed as shared or BOND PAIRS , and unshared or LONE PAIRS. • •• H Cl
shared or
bond pair
lone pair (LP)
This is called a LEWIS structure.

14. Steps for Building a Dot Structure
Ammonia, NH3
1. Decide on the central atom; never H. Why?
If there is a choice, the central atom is the atom with lowest electronegativity.
Therefore, N is central on this one
2. Add up the number of valence electrons that can be used.
H = 1 and N = 5
Total = (3 x 1) + 5
= 8 electrons

15. Building a Dot Structure
3. Form a single bond between the central atom and each surrounding atom (each bond takes 2 electrons!) H N
4. Remaining electrons form LONE PAIRS to complete the octet as needed (or duet in the case of H). H •• N
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair (2 electrons).

16. Building a Dot Structure
Check to make sure there are 8 electrons around each atom except H. H should only have 2 electrons. This includes SHARED pairs. H •• N
6. Also, check the number of electrons in your drawing with the number of electrons from step 2. If you have more electrons in the drawing than in step 2, you must make double or triple bonds. If you have less electrons in the drawing than in step 2, you have made a mistake!

17. Carbon Dioxide, CO2 1. Central atom = 2. Valence electrons =
3. Form bonds.
C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons
This leaves 12 electrons (6 pair).
4. Place lone pairs on outer atoms.
5. Check to see that all atoms have 8 electrons around it except for H, which can have 2.

18. Carbon Dioxide, CO2
C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons
Do we have octets?
6. Not every atom has its octet. We must form DOUBLE BONDS between C and O. Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each atom and replaced with another bond.

19. Double and even triple bonds are commonly observed for C, N, P, O, and S
H2CO
SO3
C2F4

20. Now You Try One! Draw Sulfur Dioxide, SO2

21. Electronegativity and Polarity
The difference in electronegativities between bonding atoms may be used to predict the type of bond that forms.
For our purposes we will consider bonds between metals and nonmetals to be ionic and bonds between two nonmetals to be covalent.

22. Although all covalent bonds involve a sharing of one or more pairs of electrons between bonding atoms, most of the time this sharing is not equal.
One of the atoms is more electronegative than the other so it “hogs” the shared bonded pair of electrons more of the time.
This type of covalent bond is known as a polar covalent bond.

23. Electronegativity Difference
If the difference in electronegativities is between:
1.7 to 4.0: Ionic
0.5 to 1.7: Polar Covalent
0.0 to 0.4: Non-Polar Covalent
Example: NaCl
Na = 0.8, Cl = 3.0
Difference is 2.2, so
this is an ionic bond!

24. Bond Polarity Check the polarity of HCl: 3.0 – 2.1 = 0.9
Because it is polar, it has positive and negative ends
Cl has a greater share in bonding electrons than does H.
Cl has slight negative charge (-d) and H has slight positive charge (+ d) called DIPOLES

25. Bond Polarity “Like Dissolves Like” Polar dissolves Polar
Nonpolar dissolves Nonpolar

26. Polarity of Whole Molecule
Determining the polarity of a bond is easy, but what about the molecule as a whole?
Strangely enough, you can have nonpolar molecules that contain polar bonds!
SO….
How do you determine if a molecule itself is going to be polar or nonpolar?
Look at the Lewis Structure and check out the central atom!

27. Determining if a Molecule is NonPolar or Polar
Look at what is attached to the central atom
-If every attached structure is the same
-If the attached structures are different
-The molecule is symmetrical
-The molecule is asymmetrical
-There are NO lone pairs on central atom
-There are lone pairs on central atom
And therefore NONPOLAR
And therefore POLAR

28. Polar Bond
δ -
δ + H Cl
Asymetrical
Polar Molecule

29. O = C = O The molecule is nonpolar So…
Polar Bonds
δ -
δ +
δ +
δ - O = C = O
However, the shape is symmetrical
So…
The molecule is nonpolar

30. Nonpolar Bond Br Br
Symmetrical
Nonpolar Molecule

31. MOLECULAR GEOMETRY

32. VSEPR MOLECULAR GEOMETRY Valence Shell Electron Pair Repulsion theory.
Molecule adopts the shape that minimizes the electron pair repulsions.
VSEPR
Valence Shell Electron Pair Repulsion theory.
Most important factor in determining geometry is relative repulsion between electron pairs.

33. VSEPR charts
Use the Lewis structure to determine the geometry of the molecule
How the electrons are arranged determines the bond angles and shapes.
Arrangement focuses on the CENTRAL atom for all data!
Think REGIONS WHERE ELECTRONS ARE LOCATED rather than bonds (for instance, a double bond would only be 1 region)

34. Some Common Geometries
Linear
Tetrahedral
Trigonal Planar

35. Structure Determination by VSEPR
The electron pair geometry is TETRAHEDRAL
1) Looking at the central atom, how many bonding regions are there?
2) How many bonded atoms? How many lone pairs?
Water, H2O
The molecular geometry is BENT.
2 bond pairs
2 lone pairs

36. Structure Determination by VSEPR
Ammonia, NH3
Draw!
The electron pair geometry is Tetrahedral , and the Geometry name is TRIGONAL PYRAMID.

37. Van der Waal Forces of Attraction
In covalent compounds the bonds within a molecule are quite strong but the attraction between different molecules is relatively weak.
These weak forces of attraction between individual molecules are known as intermolecular forces (IMFs) or Van der Waal Forces.
There are three kinds of Intermolecular forces – London Dispersal Forces, Dipole-Dipole Interactions, and Hydrogen Bonds.

38. London Dispersal Forces
London Dispersal Forces are the weakest of the intermolecular forces and are caused by the temporary shifts of electrons in the electron clouds.
These occur when nonpolar molecules are attracted to other nonpolar molecules.

40. Dipole-Dipole Interactions are attractions between oppositely-charged regions of polar molecules which allows one polar molecule to “stick” to another polar molecule.

41. Hydrogen Bonding is a special type of dipole-dipole interaction between a hydrogen on one polar molecule and an nitrogen, oxygen, or fluorine on an other polar molecule.

42. Hydrogen Bonding examples
The best example of hydrogen bonding is water
The strong IMFs create LARGE amounts of surface tension
Explains why you can float a paper clip on the surface of water, but it will sink if it’s already under it.
Aquatic bugs use this to walk on water
Also explains why water is the only substance where the solid form (ice) floats on top of the liquid form.

43. Diagram/Summary of IMFs
Hydrogen bonding – H, and N or O or F
Dipole-Dipole forces – all polar compounds
Polar Molecules
London Dispersion forces – all compounds

44. IMFs and Phase Changes
IMFs determine if the compound would rather be a solid, liquid, or gas at room temperature
Also determine melting and boiling pts of the compound
More kinds of IMFs means harder to change phase
In addition to phase, intermolecular forces also play a role in compressibility, fluidity, viscosity, surface tension, and capillarity.