1. Covalent Compounds

2. Two Types of Bonds Ionic: Electrons are transferred
Covalent: Electrons are shared
Non-polar covalent: equally shared
Polar Covalent: unevenly shared

3. Bond Polarity

4. Review: What is electronegativity?
ability of an atom to attract electrons
Which element is the most electronegative?
Fluorine
- Has 7 valence e- and wants 8
H F

5. Polar bond :
covalent bond with greater electron density around one of the two atoms H F
electron poor
region
electron rich
region
e- poor
e- rich F H d+ d-

6. 1 18 2

7. Electronegativity Difference
What type of Bond is it?
Electronegativity Difference
Bond Type
0 to 0.3
Nonpolar Covalent
0.4 to 1.6
Polar Covalent
 1.7
Ionic
Increasing difference in electronegativity
Nonpolar
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-

8. Classify the following bonds as ionic, polar covalent,or covalent:
Cs to Cl
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H to S
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
Polar Covalent
Cl to N
Cl – 3.0
N – 3.0
3.0 – 3.0 = 0
Nonpolar Covalent

9. Do you notice a pattern for the combo of elements that are ionic vs covalent?
Ionic bonds form between:
Covalent bonds form between:
Identify the following as ionic, covalent, or both:
CaCl2 BaSO4
CO2 AlPO4
SO3 H2O

10. Properties of Covalent Compounds
Usually soft and squishy
Not soluble in water
Does not conduct electricity
Soluble in organic solvents
Low melting points
Low boiling points

11. Properties of Ionic Compounds
Combination of ions (cation/anion)
Tightly packed solids in a crystal lattice
Hard and Brittle
Usually soluble in water
Conducts electricity when dissolved
High melting points
High boiling points

12. Naming Covalent Compounds

13. NAMING COMPOUNDS Nonmetal – Nonmetal USE PREFIXES!
Change the ending of the second word to -ide
No mono on the first word
Drop any double vowels

14. Covalent Prefixes Number of Atoms Prefix 1 Mono- 2 Di- 3 Tri- 4 Tetra-5
Penta- 6
Hexa- 7
Hepta- 8
Octa- 9
Nona- 10
Deca-

15. Examples CO CO2 SO2 SO3 N2H4 N2O3 Carbon Monoxide Carbon Dioxide
Sulfur Dioxide
Sulfur Trioxide
Dinitrogen Tetrahydride
Dinitrogen Trioxide

16. Examples Si2F6 C3Cl8 disilicon hexafluoride tricarbon octachloride
PBr5 NO
SeF2
H2O
disilicon hexafluoride
tricarbon octachloride
phosphorus pentabromide
nitrogen monoxide
selenium difluoride
dihydrogen monoxide

17. EMPIRICAL AND MOLECULAR FORMULAS

18. Define Empirical Formula:
A chemical formula that gives the simplest whole-number ratio of the elements in the formula.
Which of the following is an empirical formula?
CO C2O4
N2H4 NH2

19. Define Molecular Formula:
A chemical formula that gives the actual number of the elements in the molecular compound.
For the following molecular formulas, write the empirical formula:
Molecular: Empirical:
C2H4
C6H12O6
C9H21O6N3

20. Lewis Structures

21. This is called the octet rule
Eight electrons in the valence shell (filling s and p orbitals) make an atom STABLE
This is called the octet rule
Bond formation follows the octet rule… Chemical compounds tend to form so that each atom: by gaining, losing, or sharing electrons, has an octet of electrons in its valence energy level.

22. Lewis Dot Diagrams
an electron-configuration notation with only the valence electrons of an element are shown, indicated by dots placed around the element’s symbol.
tracks the number of valence electrons
the inner core electrons are not shown

23. Lewis Dot Practice
Li N F Be O Ne

24. Lewis Structures for Compounds
The pair of dots between two symbols represents a shared pair.
How many shared pairs does each fluorine have below?
An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

25. Each dash represents TWO electrons
Lewis Structures
The shared pair of electrons is often replaced by a long dash.
Each dash represents TWO electrons

26. Why should two atoms share electrons?
To get a valence of 8 electrons!
7e-
7e-
8e-
8e- F F + F
Lewis structure of F2
lone pairs F
single covalent bond
single covalent bond F

27. Multiple Covalent Bonds
double bond: covalent bond in which two pairs of electrons are shared between two atoms
shown by two side-by-side pairs of dots or by two parallel dashes

28. Multiple Covalent Bonds
triple bond: covalent bond in which three pairs of electrons are shared between two atoms
shown by three side-by-side pairs of dots or by three parallel dashes

29. Lengths of Covalent Bonds
Bond Type
Bond Length
(pm)
C-C
154
CC
133
CC
120
C-N
143
CN
138
CN
116
Bond Lengths
Triple bond < Double Bond < Single Bond

30. Bond Length and Bond Energy
As atomic size increases, bond length increases, and as a result bond energy decreases
As you increase the number of bonds between two atoms, energy increases, while bond length decreases.
Rubber band example

31. Bond Length and Bond Energy Examples
Which bond is greater in length: Br2 or F2?
The HF bond is 570 pm, the H2 bond is 436 pm, which bond requires more energy to break?
Which bond would require more energy to break C-C single bond or C=C double bond? Which bond is longer?

32. Writing Lewis Structures
Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.
Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.
Complete an octet for all atoms except hydrogen
If structure contains too many electrons, form double and triple bonds on central atom as needed.

33. Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons F N

34. Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e-
4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete
octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24 O C

35. resonance structure:
When there are two or more Lewis structures for a single molecule
What are the resonance structures of the
carbonate (CO32-) ion? O C - O C - O C -

36. Some elements do not follow the octet ruleBe F B
There can also be expanded octets!

37. Molecular Geometry

38. VSEPR THEORY Lewis Dot Diagrams are 2D but we live in a 3D world.
How are molecules actually arranged??
Follows the Valance Shell Electron Pair Repulsion Theory or VSEPR

39. Number of Surround Atoms
AB2 – Linear
Number of Surround Atoms
Number of Lone Pairs
Bond Angle 2
180˚ Cl Be

40. Number of Surround Atoms
AB3 – Trigonal Planar
Number of Surround Atoms
Number of Lone Pairs
Bond Angle 3
120˚

41. Number of Surround Atoms
AB2E1 – Bent
Number of Surround Atoms
Number of Lone Pairs
Bond Angle 2 1
<120˚

42. Number of Surround Atoms
AB4 – Tetrahedral
Number of Surround Atoms
Number of Lone Pairs
Bond Angle 4
109.5˚

43. AB3E1 – Trigonal Pyramidal
Number of Surround Atoms
Number of Lone Pairs
Bond Angle 3 1
107˚

44. Number of Surround Atoms
AB2E2 – Bent
Number of Surround Atoms
Number of Lone Pairs
Bond Angle 2
104.5˚

45. Predicting Molecular Geometry
Draw Lewis structure for molecule.
Count number of lone pairs on the central atom and number of atoms bonded to the central atom.
Use VSEPR to predict the geometry of the molecule.
What are the molecular geometries of SO2 and SF4? C F S O
AB4
AB2E
tetrahedral
bent

46. Intermolecular Forces

47. Intermolecular forces: attractive forces between molecules.
Intramolecular forces: attractive forces within a molecule (the bonds)
Intermolecular Forces
Intramolecular Forces
Intramolecular Forces
intramolecular forces are much stronger than intermolecular forces

48. Dipoles What is a dipole? A polar molecule
Uneven sharing of electrons so there is a separation of charge

49. Dipole-Dipole Forces
Attraction between two polar molecules — — + +

50. Hydrogen Bonding Special type of Dipole – Dipole
Attraction between: Hydrogen and Nitrogen/Oxygen/Fluorine

51. Dipole – Induced Dipole
Attraction between one polar and one nonpolar molecule — +
Electrons shift toward positive end of dipole — — + +

52. London Dispersion Forces
Attraction between two nonpolar molecules
Electrons become uneven and form a dipole — — + +

53. Strength of IMF Hydrogen Bond Dipole – Dipole Dipole – Induced Dipole
London Dispersion Forces
strongest
weakest

54. Which of the following molecules is polar?
H2O, CO2, SO2, and CH4 O H S O
dipole moment
polar molecule
dipole moment
polar molecule C H C O
no dipole moment
nonpolar molecule
no dipole moment
nonpolar molecule

55. CH4 is nonpolar: dispersion forces.
What type(s) of intermolecular forces exist between each of the following molecules?
HBr
HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: dispersion forces. S O
SO2
SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

56. What does IMF effect? Viscosity Surface Tension Cohesion/Adhesion
Boiling Point

57. Stronger IMF  Higher Viscosity
Measures a fluid’s resistance to flow
Stronger IMF  Higher Viscosity

58. Stronger IMF  Higher Surface Tension
result of an imbalance of forces at the surface of a liquid.
Stronger IMF  Higher Surface Tension

60. Stronger IMF  Higher Boiling Point
Point at which liquid particles escape the surface of the liquid into the gas phase
Stronger IMF  Higher Boiling Point

61. Adhesion and Cohesion
Cohesion: intermolecular attraction between like molecules
Adhesion: intermolecular attraction between unlike molecules
Adhesion
Cohesion