1. Bonding

2. Electronegativity – attraction to another atom’s electrons
Chemical Bonds
A force of attraction between 2 atoms
The “glue” that holds chemical compounds together
**Involves VALENCE electrons**
Electronegativity – attraction to another atom’s electrons
Table S
Outer electrons!

3. 3 Types of Bonds Ionic Bonds – metal + nonmetal atoms
“transfer of electrons”
Electronegativity Difference (E.N.D.)
greater than 1.7
Metallic Bonds – metal + metal atoms
“sea of electrons”
Covalent Bonds – nonmetal + nonmetal atoms
“sharing electrons”
less than 1.7
Catch!
Tug-of-War

4. Energy of Bonds Exothermic – Energy Released
Forming Bonds
Atoms are more stable together, and so, RELEASE ENERGY to be at a lower, more stable energy state
Endothermic – Energy Absorbed
Breaking Bonds
Overcome the attraction between the atoms to break a bond, so we ADD ENERGY
The stronger the bond, the closer together the atoms are, the more energy is needed to break the bond

5. X Lewis Dot Structures Use dots to represent electrons
Diagram of the valence electrons in an atom or ion
Use dots to represent electrons X 3 7
Each consecutive dot will go opposite the one before it, before getting paired up. 1 6 5 2
Octet Rule – 8 Electron Rule
Atoms will gain, lose, or share electrons to have 8 valence electrons
(except H, He, Li, Be, B) 8 4

7. Electronegativity Difference
E.N.D. > 1.7
One atom is far more electronegative than the other, results in a transfer of electrons
0.4 < E.N.D. < 1.7
One atom is slightly more electronegative than the other, results in an UNEQUAL sharing of electrons
E.N.D. < 0.4
Two atoms of equal electronegativity, results in EQUAL sharing of electrons
Ionic Bond
Polar Covalent Bond
Nonpolar Covalent Bond

8. Ionic Bonding E.N.D. > 1.7 Transfer of electrons Metal + Nonmetal
One atom far more electronegative than the other
Transfer of electrons
Metals give up electrons to nonmetals
Metal + Nonmetal
Metal (cation)
Nonmetal (anion)
Lewis Structures
Write the neutral atom electron dot structures
Draw arrows to show the movement of electrons from metal(s) to nonmetal(s)
Draw ion electron dot structures bonded together
Na + Cl → [Na]+[ Cl ]-

10. Ionic Solids – Properties
Solid substances formed from ionic bonds
Crystalline structure
High melting point
Dissolve in water (aqueous solution)
Don’t conduct heat or electricity as SOLID
Conduct electricity as
aqueous solutions or
liquids

11. Metallic Bonding 2 or more metal atoms in a sea of electrons
Strong bonds form between 2 metals in liquid or solid phase
Evidence: high melting & boiling point
Properties
Conductive of heat & electricity
LIQUID
SOLID
Malleable
Ductile
High Melting/Boiling Points

12. Covalent Bonding E.N.D. < 1.7 Sharing of electrons
Nonmetal atoms of similar electronegativity will share electrons to complete octets
Sharing of electrons
Polar Covalent – unequal sharing
Nonpolar Covalent – equal sharing

13. Molecules Smallest unit of a compound made of covalent bonds
Exist in all 3 phases of matter (depending on forces of attraction between molecules)
Properties
Soft
Low melting points
Do not conduct heat/electricity

14. sand (silicon dioxide)
Molecules
Network Solids – bond in a strong network of atoms
Have slightly different properties due to the network
Harder
High melting points
Do not conduct heat/electricity (except graphite)
Coordinate Covalent Solids –
Occur in polyatomic ions
One atom donates BOTH electrons to a bond
Compounds with polyatomic ions contain BOTH IONIC & COVALENT BONDS!!
diamond
graphite
aesbestos
sand (silicon dioxide)

15. Lewis Structures Add up valence electrons
The lowest electronegative element goes in the center (usually there is only 1 atom)
Arrange all the other atoms around the central atom
Draw a single bond from each outside (terminal) element to the central atom
Draw dots to complete octets for outside atoms
Add any extra electrons to complete the octet on the central atom
If we don’t have enough electrons to complete this octet, make a multiple bond (double or triple) – C, N, O, P, S
Count electrons – check for all valence electrons used

17. Molecule Geometry VSEPR (Valence Shell Electron Pair Repulsion)
Electrons want to be as far apart as possible
Electron-pair geometry refers to the number of electron pairs surrounding the central atom
Molecule geometry refers to Lewis Structures and depends upon both bonding & nonbonding electrons
NOTE – double and triple bonds “count” as ONE pair

18. Molecular Geometry (VSEPR)
# Bonds (terminal e- pairs)
Examples
Shape
1, 2
H—F , O=C=O
Linear 3
Trigonal Planar 4
Tetrahedral 5
Trigonal Bipyramidal 6
Octahedral

19. Molecular Geometry 2 1 Bent 3 Trigonal Pyramidal Linear T-Shaped 4
# Bonding Pairs
# Nonbonding Pairs
Shape 2 1
Bent 3
Trigonal Pyramidal
Linear
T-Shaped 4
See Saw 5
Square Pyramidal
Square Planar

20. Molecule Polarity 2 factors affect molecule polarity Bond Polarity
If a molecule contains NO polar bonds, it will be nonpolar
If a molecule contains polar bonds it could be either polar or nonpolar
Molecule Geometry (Shape)
If a molecule is symmetrical there is an equal distribution of electrons and so is NONPOLAR
If a molecule is asymmetrical there is an unequal distribution of electrons and so is POLAR

21. Intermolecular Forces
Forces of attraction between molecules – based on polarity (a.k.a. Van der Waals Forces)
More strongly polar moleculs will have greater Intermolecular Forces (IMFs)
Affect melting and boiling points, surface tension, viscosity

22. Intermolecular Forces
Dipole-Dipole Interactions
The slightly positive end of one molecule is attracted to the slightly negative end of another
Hydrogen Bonding
A specific type of dipole interaction where the slightly positive H is attracted to the slightly negative N, O, or F of another

23. Hydrogen Bonding in Water
WATER IS POLAR!!!!!
Slightly + H is attracted to Slightly – O