1. Bonding

2. Electrons

3. Electron Configurations
Electron Configurations are a short hand way to show where all the electrons are located in an atom

4. How to Write and Electron Configuration
Determine the number of electrons
For a neutral atom = the atomic number
For an ion= atomic number - the oxidation number
Fill the orbitals according to the orbital filling diagram until all electrons have been placed

5. Examples
Lithium:
Phosphorous:
Chlorine:
Cl-1:
Mg+2:

6. Valence Electrons
Valence electrons are the electrons located in the outermost energy level of the atom
The number of valence electrons can be predicted by the position on the periodic table
Highlight the valence electrons in the examples above

7. Predicting Valence Electrons

8. Lewis Dot Structures
Diagrams that use the chemical symbol and dots to represent the valence electrons for an atom
How to draw Lewis Structure for an atom:
Write the symbol
Place one dot on each side of the symbol to represent the valence electrons until all valence electrons are placed

9. Examples:
Lithium
Phosphorous
Chlorine
Magnesium
Argon

10. The Octet Rule
Atoms will loose, gain, or share electrons in order to acquire a full outer shell of 8 valence electrons
Exceptions include those of period 1, Lithium, Beryllium, and Boron which prefer to have a full outer shell of 2
Once the atoms has bonded in a way to provide a full outer shell, we say it has a stable octet

11. Electronegativity
Electronegativity is the attraction an atom has for electrons in a chemical bond
The electronegativity of an atom determines what will happen to its electrons when bonding

12. Ionic Compounds
Involve a transfer of electrons from a metal to a nonmetal
Results in a metal cation and an a nonmetal anion that then bond in order to neutralize the charge between the two
Ionic compounds have an electronegativity difference greater than or equal to 2.0

13. Covalent Compounds Also known as Molecular Compounds or Molecules
Involve a sharing of electrons between two nonmetals
In a covalent compound there is no neutralization of charge
Atoms share electrons so that all the atoms can obtain a stable octet
Covalent compounds have an electronegativity difference less than 2.0

14. Lewis Structure for Covalent Compounds
Count the total number of valence electrons for ALL atoms in the molecule
If the molecule is charged subtract the oxidation number
Draw the skeletal structure for the molecule
Put the most electronegative or single atom in the center
Hydrogen is always on the outside
Draw a single bond between each of the atoms in the compound (two dots, or one solid line)
Subtract the bonding electrons from the total number
Place the remaining electrons to fill the octets of the outer atoms
Place any remaining electron pairs on the central atom
If the central atom does not have a complete octet move atoms to form multiple bonds until it does

15. Simple Covalent Compounds
Water: H2O
Methane: CH4

16. The Lewis Challenge
Draw the Lewis Structure for the following compounds based on the conditions below:
You must satisfy the octet rule for all atoms involved.
Some atoms may share more than one pair of electrons.
All valence electrons for all atoms must be accounted for.
F2 N2 HCN

17. Covalent Compounds with Multiple Bonds
Single Bonds- two atoms share one pair of electrons
Double Bonds- two atoms share two pairs of electrons
Triple Bonds- two atoms share three pairs of electrons

18. Examples
Carbon Dioxide: CO2
Nitrogen: N2

19. Covalent Compounds with Charges (Polyatomic Ions)
For a charged covalent compound, draw the structure the same way then place the entire structure in brackets with the charge on the outside

20. Examples
Nitrate: NO3-1
Carbonate: CO3 -2

21. VSEPR and Molecular Geometry

22. VSEPR VSEPR stands for valence shell electron pair repulsion
A model for predicting the geometry of molecules based on the arrangement of atoms that minimizes the repulsion of bonding and nonbonding pairs of electrons

23. Bonding pairs= pairs of electrons that are participating in a bond
Nonbonding pairs= pairs of electrons not participating in a bond (also sometimes called lone pairs)

24. Predicting Molecular Geometry

25. Linear
Compound
Dot Diagram
VSEPR Model
HCl

26. Bent/Angular
Compound
Dot Diagram
VSEPR Model
H2O

27. Trigonal Planar
Compound
Dot Diagram
VSEPR Model
AlCl3

28. Trigonal Pyramidal
Compound
Dot Diagram
VSEPR Model
NF3

29. Tetrahedral
Compound
Dot Diagram
VSEPR Model
CH4

30. Molecular Polarity

31. Polar Molecules
Polar molecules are molecules in which electrons are not shared evenly
The charges referred to in a polar molecule are much weaker than those in an ionic compound

32. Polar molecules are sometimes called a dipole because they have two ends with slightly different charges
A slightly negative end
Labeled δ-
Always associated with the more electronegative atom
A slightly positive end
Labeled δ+

33. Polarity can be determined by Electron Placement
Polarity can be predicted based on the number of bonding and nonbonding pairs
Any molecule with a high concentration of nonbonding electrons around an electronegative atom will be a polar molecule
(Delete the line that says: “polar molecules will be one of three molecular geometries”)

34. Examples
Compound
HCl
H2O
NH3
CHCl3
Dot Diagram
Molecular Geometry

35. Bond Polarity
Nonpolar molecules can still have polarity in their bonds
The more electronegative atom will be where the electrons spend more of their time, giving that end of the bond a slightly negative charge
Bond polarity can be calculated using electronegativity.

36. Bond Polarity
Example: Methane

37. Intra and Intermolecular Forces

38. Intramolecular Forces
Intra= inside
Intramolecular forces are the strong forces within a chemical bond

39. Ionic Bonds
Electrostatic attractions between a cation (+ ion) and an anion (- ion)
Results in the formation of a crystal lattice, where every cation is surrounded by anions and every anion is surrounded by cations
Ionic bonds are the strongest chemical bonds

40. Crystal Lattice Illustration

41. Metallic Bonds
Attractions between metal cations and delocalized electrons
Results in a network similar to the ionic crystal lattice
Metallic bonds are slightly weaker than ionic bonds

42. Electron Sea Model

43. Covalent Bonds Result in the formation of molecules
forces between the electrons of one atom and the nucleus of another
Result in the formation of molecules
Slightly weaker than metallic bonds, with the exception of the covalent bonds found in network covalent compounds (which are stronger than ionic bonds)

44. Intermolecular Forces
Inter= between
Intermolecular forces are the forces between molecules
They are ONLY found in covalent compounds, in between molecules

45. Dipole-Dipole Interactions
Weak attractions between the δ- end of one molecule and the δ- end of another
Example: H-Cl

46. Hydrogen Bonding
Dipole interactions between hydrogen of one molecule and the nitrogen, oxygen, or fluorine of another
Example: Water

47. Dispersion Forces
Attractions between nonpolar compounds, between bonding electrons of one molecule and the nucleus of atoms in another molecule
Example: Methane

48. Strength of Intermolecular Forces
Ionic Metallic Covalent Hydrogen Bonding Dipole-Dipole Interactions Dispersion Forces
DECREASES

49. Properties and Strengths of Compounds

50. Types of Compounds
Network Covalent- compounds that are composed of atoms covalently bonded to each other multiple times
Examples: diamond, sand

51. Ionic- compounds that result from the attraction of positive ions (cations) to negative ions (anions)
Examples: NaCl Mg(OH)2

52. Polar Covalent Compounds- compounds where electrons are shared unevenly
Example: water

53. Non-Polar Covalent Compounds-compounds where electrons are shared evenly
Example: Methane

54. Strengths of Compounds
Type of Compound
Relative Strength
Types of Forces
MP and BP
Solubility
Conductivity
Network Covalent
Ionic
Metallic
Polar Covalent
Non-Polar Covalent

55. Strengths of Compounds
Type of Compound
Relative Strength
Types of Forces
MP and BP
Solubility
Conductivity
Network Covalent
Only covalent bonds
highest
Not soluble in anything No
Ionic
Ionic bonds
Soluble in water and polar solvents
Yes
Metallic
Metallic bonds
N/A
Polar Covalent
Covalent bonds, hydrogen bonds, dipole interactions
Non-Polar Covalent
Covalent bonds, dispersion forces
lowest
Soluble in nonpolar solvents only

56. Organic Molecules

57. Defining Organic Molecules
All organic molecules contain a large amount of carbon
Monomers- single organic subunits
Polymers- long chains of monomers
Importance of Polymers
Biological Molecules- DNA, proteins, carbohydrates, lipids
Industry- plastics

58. Alkanes Most basic organic compounds Also called hydrocarbonds
Made of only carbon and hydrogen with single bonds
General Formula for alkanes: CnH2n+2

59. Names and Formulas for Alkanes
CH4
C6H14
C2H6
C7H16
C3H8
C8H18
C4H10
C9H20
C5H12
C10H22

60. Other Organic Compounds
Variation on alkanes with the addition of functional groups
Functional Group- a group of molecules that when added to an alkane base, changes the physical and chemical properties of the compound

61. Alkene Functional Group: carbon-carbon double bond
Properties: Stronger than alkanes

62. Alkyne Functional Group: carbon-carbon triple bond
Properties: stronger than alkanes and alkenes; always linear

63. Alcohols Functional Group: hydroxyl group
Properties: Makes the molecule polar

64. Carboxylic Acids Functional Group: carboxyl group
Properties: polarity, acidic properties