1. Chapter 6.2 and 6.5
Covalent Compounds

2. Covalent Bonds Sharing Electrons
Covalent bonds form when atoms share one or more pairs of electrons
nucleus of each atom is attracted to electron cloud of other atom
neither atom removes an electron from the other

3. Covalent Bonding

4. Covalent Bonds Sharing Electrons Covalent bonds
space where electrons move is called molecular orbital
made when atomic orbitals overlap

5. Molecules

6. Covalent Bonds Energy and Stability
Noble gases are stable (full octet) (low P.E.)
Other elements are not stable (high P.E.)
covalent bonding decreases potential energy because each atom achieves electron configuration like noble gas

7. Covalent Bonds Energy and Stability
because P.E. decreases when atoms bond, energy is released
i.e., atoms lose P.E. when they bond
loss of P.E. implies higher stability

9. Covalent Bonds Energy and Stability
potential energy determines bond length
at minimum P.E., distance between two bonded atoms is called bond length
bonded atoms vibrate
therefore, bond length is an average length

10. Covalent Bonds Energy and Stability bonds vary in strength
bond energy is the amount of energy required to break the bonds in 1 mol of a chemical compound
bond energy predicts reactivity
bond energy is equal to loss of P.E. during formation

11. Bond Energies and Lengths

12. Bonds and Energy
Single bonds are the longest bonds with the least bond energy
Double bonds are shorter, stronger and have intermediate bond energy
Triple bonds are the shortest, strongest, and have the highest bond energy

13. Covalent Bonds Electronegativity
Atoms share electrons equally or unequally
nonpolar covalent bond: bonding electrons shared equally
polar covalent bond: shared electrons more likely to be found around more electronegative atom

14. Covalent Bonds Electronegativity
Atoms share electrons equally or unequally
difference in electronegativity can be used to predict type of bond (but boundaries are arbitrary)

16. Bond Types
0.3
1.7

17. Practice: Calculate the bond type
N and H
F and F
Ca and Cl
C and O
Polar
Non-polar
Ionic

18. Covalent Bonds Electronegativity
Polar molecules have positive and negative ends
such molecules called dipoles
δ means partial in math and science
positive end—δ+
negative end—δ-
example: Hδ+Fδ-

19. Electronegativity Difference for Hydrogen Halides

20. Covalent Bonds Electronegativity Polarity is related to bond strength
greater electronegativity
means
greater polarity
greater bond strength

21. Covalent Bonds Electronegativity
Bond type determines properties of substances
metallic bonds: electrons can move from one atom to another—good conductors
ionic bonds: hard and difficult to break apart
covalent bonds: low melting/boiling points

22. Properties of Substances with Different Types of Bonds

23. Polarity of Molecules
Two atoms: bond polarity is the molecular polarity
More than 2 atoms: the geometry of the molecule must be considered
If the bonds are non-polar, the molecular is non-polar
Some molecules with polar bonds
can be non-polar

24. More
Sometimes the partial charges cancel each other out because they are directly opposite each other
Consider CO2 and CCl4
The symmetrical distribution of the bonds leads to cancellation of the charges

25. Polar and Non-polar Molecules

26. Drawing and Naming Lewis Electron-Dot Structures
Lewis structures represent valence electrons with dots
position of electrons is symbolic (not literal)
shows only the valence electrons of an atom
dots around atomic symbol represent electrons

27. Lewis Structures of Second-Period Elements

28. Drawing and Naming
Lewis Electron-Dot Structures
Cl2
HCl

29. Drawing and Naming Lewis Electron-Dot Structures 1. Gather information
draw Lewis structure for each atom in compound; place one electron on each side before pairing
determine total number of valence electrons

30. Drawing and Naming Lewis Electron-Dot Structures Drawing
2. Arrange atoms
arrange structure to show bonding
halogens and hydrogen usually make one bond at end of molecule
carbon usually in center

31. Drawing and Naming Lewis Electron-Dot Structures Drawing
3. Distribute the dots so that each atom satisfies octet rule (except H, Be, B)
4. Draw the bonds as long dashes
5. Verify the structure by counting number of valence electrons

32. Drawing and Naming Lewis Electron-Dot Structures Polyatomic Ions
use brackets [] to show overall charge
example:

33. Drawing and Naming Lewis Electron-Dot Structures Multiple Bonds
sharing two pairs of electrons is a double bond
sharing three pairs of electrons makes triple bonds
example:

34. Drawing and Naming Lewis Electron-Dot Structures Resonance Structures
sometimes, multiple structures are possible
show all possibilities
example:

35. Drawing and Naming Naming Covalent Compounds
First name: name of first element in formula
usually least electronegative
requires a prefix if more than one of them
Second name: ends in –ide

36. Naming Covalent Compounds

37. Practice Naming 1. antimony tribromide 2. hexaboron silicide
3. chlorine dioxide
4. iodine pentafluoride
5. dinitrogen trioxide
6. ammonia

38. More Naming Practice
P4S5
2. PI3
3. NO
4. N2F4
5. CO2
6. H2O

39. Molecular Shapes Determining Molecular Shapes
Three-dimensional shape helps determine physical and chemical properties
valence shell electron pair repulsion (VSEPR) theory predicts molecular shapes
based on idea that electrons repel one another

40. Molecular Shapes Lewis structures show which atoms are
connected where, and by how many bonds, but they don't properly show 3-D shapes of molecules.
To find the actual shape of a molecule, first draw the Lewis structure, and then use VSEPR Theory.

41. VSEPR Valence Shell Electron Pair Repulsion theory. MOLECULAR GEOMETRY
Most important factor in determining geometry is relative repulsion between electron pairs.

42. MOLECULAR GEOMETRY
Molecule adopts the shape that minimizes the electron pair repulsions.

43. VSEPR Rules To apply VSEPR theory:
1: Draw the Lewis structure of the molecule and identify the central atom
2: Count the number of electron charge clouds (lone and bonding pairs) surrounding the central atom.
3: Predict molecular shape by assuming that clouds orient so they are as far away from one another as possible

44. VSEPR Shape Predictor Table -
117°

45. VSEPR Shape Predictor Table -

46. Bond Angles
Lone-pairs of electrons behave as if they are slightly bigger than bonded electron pairs and act to distort the geometry about the atomic center so that bond angles are slightly smaller than expected:

47. Bond Angles
Methane, CH4, has a perfect tetrahedral bond angle of 109.5°, while the H-N-H bond angle of ammonia, NH3, is slightly less at 107°, trigonal pyramidal

48. Bond Angles
                                                                            
The oxygen of water has two bonded electron pairs and two nonbonded "lone" electron pairs giving a total VSEPR coordination number of 4.
But the geometry is defined by the relationship between the H-O-H atoms and water is said to be "bent" or "angular" shape of 105°.
105°

49. Molecular Shapes Determining Molecular Shapes and angles.
Let’s try some. CO
CO2
BF3
CH4
CBr4
PCl3

50. Simple Shapes
Bond angle for linear is 180°

51. Trigonal Planar
Bond angle of 120°

52. Tetrahedral
Bond angle of 109.5°

53. Bent
Bond angle of 105°