1. Bonding
Dr. Ron Rusay
Fall 2001
© Copyright 2001 R.J. Rusay

2. Chemical Bonds Definition:
Attractive forces which hold atoms together and provide a particular molecular arrangement of atoms with new chemical properties.

3. Bond Model
1. Description of valence electron arrangement (Lewis structure).
2. Prediction of geometry (VSEPR model).
3. Description of atomic orbital types used to share electrons or hold long pairs.

6. Ionic vs. Covalent Bonding

7. Electron Configurations Noble Gases and The Rule of Eight
REVIEW:
A nonmetal and a metal react to form an ionic compound: Valence electrons of the metal are lost and the nonmetal gains these electrons. (Ionic Bonding)
When two nonmetals react: They share electrons to achieve a Noble Gas Configuration. (Covalent Bonding)

8. Octet Rule: General Comments
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.
3rd row 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.

9. Lewis Electron-Dot Symbols for Elements in Periods 2 & 3

10. Isoelectronic Ions
Ions containing the the same number of electrons are isoelectronic with a Noble Gas
(O2, F, Na+, Mg2+, Al3+)
However the ion sizes differ O2> F > Na+ > Mg2+ > Al3+

11. Ionic Bonds
Result from electrostatic attractions of closely packed, oppositely charged ions.
Form when an atom which can easily lose electrons reacts with one which has a high electron affinity, that is, it can easily gain electrons.
Mg and Cl; K and O

12. Bond Energy
It is the energy required to break a bond, i.e. overcome the force of attraction.
The quantitative value provides information about the strength and nature of the bond.

13. Three Ways of Showing the Formation of LiF, [Li+ and F -] through Electron Transfer

14. Covalent Bond Lengths Interatomic distance.
It is the distance where the bond energy is at a minimum value, and which is the most stable atomic form.
IR Tutor

15. Covalent Bonding & Bond Length

16. Bond Lengths and Covalent Radius

17. Bond Lengths : Bond Strengths
The Periodic Table
Bond Lengths : Bond Strengths

19. Actual vs. Expected Bond Length
Actual = nm
Expected = nm

20. Electronegativity

21.  = (H  X)actual  (H  X)expected
Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself.
 = (H  X)actual  (H  X)expected

23. Periodic Trend

24. Electronegativity Differences & Polar Covalent Bonds
A molecule with a relatively high difference in electronegativities, such as HF, has a center of positive charge and a center of negative charge. It is polar, having an experimentally measureable dipole moment.

25. The spectrum of bond differences: Ionic at one extreme : Non-polar covalent at the other.

26. Bond Shapes

27. 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.

28. Lewis Structure of Water

31. Carbonate Ion CO32- :
The carbonate polyatomic ion has two coordinate bonds, resonance forms and two electrons from a cation!

32. Draw the Lewis Structure for Sulfur trioxide

33. SO3 vs SO2 : The Coordinate Bond

34. Where to draw the sulfur-oxygen double bond? There are three options.
Simplified drawings without free pairs of electrons:
Experimental data shows that each of the three sulfur-oxygen bonds are the same length.

35. Resonance
Occurs when more than one valid Lewis structure can be written for a particular molecule.
These are resonance structures. The actual structure is an average of the resonance structures.

36. Delocalized Electron-Pairs
Resonance:
Delocalized Electron-Pairs
Ozone : O3 .. .. .. .. .. O O .. .. .. .. .. .. O O .. O O II I
Resonance Hybrid Structure .. O .. .. .. .. O O
One pair of electron’s resonates
between the two locations!!

37. Lewis Structures of Simple MoleculesH H H ..
CH4 H C C .. O H
Methane H H
Ethyl Alcohol (Ethanol) .. .. .. .. C F .. O K+ .. .. Cl .. .. .. .. .. O .. O
CF4 ..
KClO3
Potassium Chlorate
Carbon Tetrafluoride

38. Lewis Structures of Simple MoleculesH H H N
. . H H
. .
. . N N
Ammonia H H C H +
. . O
. . H N H
Urea H
Ammonium Ion

39. Formal Charge
Equals the number of valence electrons of the free atom minus [the number of unshared valence electrons in the molecule + 1/2 the number of shared valence electrons in the molecule.

40. Resonance and Formal Charge
Not as good Better

42. VSEPR Model Valence Shell Electron Pair Repulsion

43. VSEPR Model
The molecular structure which surrounds a given atom is determined principally by minimizing electron pair repulsions through maximizing separations.

44. Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Put pairs as far apart as possible.
3. Determine positions of atoms from the way electron pairs are shared.
4. Determine the name of molecular structure from positions of the atoms.

45. Molecular Models Computer Generated Models
Ball and stick models of ammonia, water and methane. For many others see: