1. 6-1 Introduction to Chem Bonding
*a link between atoms that results from the mutual attraction of their nuclei for e-

2. -Types of chem bonds 1. Ionic bond: transfer of e- 2. Covalent bond:
sharing of e-’s
-bonds between unlike atoms are never completely ionic and rarely completely covalent.(Pauling)
3. Metallic bonds: between metals

3. *nonpolar-covalent bond:
cov bond which the bonding e-’s are shared equally by the bonded atoms
-resulting in a balanced
distribution of elect charge

4. *Polar:
meaning that they have an uneven distribution of charge
*Polar-cov bond:
the united atoms have an unequal attraction for the shared e-’s

5. -Why chem bonding occurs
-if their PE is lowered by the change, two atoms will form a chem bond.

6. 6-2 Covalent Bonding and Molecular Compounds
*Molecule:
smallest unit quantity of matter which can exist by itself and retains all the prop of the original substances
*Diatomic molecule:
2 like atom

7. *Molecular cmpnd:
chem cmpnd whose simplest units are molecules
*Chem formula:
represents the #’s of atoms of each kind in a chem cmpnd by using atomic symbols and numerical subscripts

8. NaCl CO2 CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular
A. Vocabulary
CHEMICAL FORMULA
IONIC
COVALENT
Formula
Unit
Molecular
Formula
NaCl
CO2

9. *Molecular formula:
shows types and numbers of atoms combined in a single molecule

10. -Formation of covalent bond
*Bond length:
aver dist between 2 bonded atoms
*Bond nrg:
nrg required to break a chem bond and form neutral atoms

11. -The Octet Rule
*Octet rule:
8 e-’s in the highest nrg level

12. -Lewis Structure
(Gilbert Lewis)
*pg 169

13. -Lewis Structure *Unshared pair:
pair of e-’s that is not involved in bonding
*single bond:
cov bond produced by sharing of 1 pair of e-’s between 2 atoms

14. *structural formula:
indicates kind, number, arrangement, and bonds of the atoms in a molecule.

15. -Multiple Covalent Bonds
*Double bond:
cov bond between 2 atoms produce by sharing 2 pairs of e- EX

16. -Multiple Covalent Bonds
*Triple bond:
cov bond between 2 atoms produced by sharing 3 pairs of e- Ex
-Sample problem 6.3

17. -Polyatomic Ions
*charged group of cov bonded atoms Ex

18. 6.3 Ionic Bonding and Ionic Compounds
composed of + and – ions combined so that the + and – charges are equal

19. 6.3 Ionic Bonding and Ionic Compounds
*Formula unit:
simplest collection of atoms from which a cmpnds formula can be estab
-Common ions:
Na+, K+,
Mg+2, Ca+2, Ba+2,
O-2, S-2, F-,Cl-, Br-,I-

20. -Formation of Ionic Bonds
*Lattice nrg:
nrg released when one mole of an ionic crystalline cmpnd is formed from gaseous ions

21. -Ionic vs Covalent bonds
-ionic bonds are stronger than cov bonds
-cov bonds
~lower mp, bp, and hardness

22. -Ionic vs Covalent bonds
-ionic bonds
~hard and brittle
~carry a current when
dissolved in water

23. B. Types of Bonds IONIC COVALENT
Bond Formation
e- are transferred from metal to nonmetal
e- are shared between two nonmetals
Type of Structure
crystal lattice
true molecules
Physical
State
solid
liquid or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical Conductivity
yes (solution or liquid) no
Other
Properties
odorous

24. 6.4 Metallic Bonding
-metals have a low ionization nrg and low electroneg.
-two metals form at best a weak covalent bond

25. 6.4 Metallic Bonding
-metals tend to form what we call a sea of electrons

26. 6.4 Metallic Bonding *Metallic bond:
chem bond resulting from the attraction between + ions and surrounding mobile electrons
-to compare strengths in metals vaporization heats are used

27. B. Types of Bonds METALLIC e- are delocalized among metal atoms
Bond Formation
e- are delocalized among metal atoms
Type of Structure
“electron sea”
Physical
State
solid
Melting
Point
very high
Solubility in
Water no
yes (any form)
Electrical Conductivity
Other
Properties
malleable, ductile, lustrous

28. -Nature of Metals
-the motion of the e- sea helps explain why metals have certain properties
*Malleability
*Ductility
-luster

29. B. Types of Bonds
RETURN

30. B. Types of Bonds
RETURN

31. Ionic Bonding - Crystal Lattice
B. Types of Bonds
Ionic Bonding - Crystal Lattice
RETURN

32. Covalent Bonding - True Molecules
B. Types of Bonds
Covalent Bonding - True Molecules
Diatomic Molecule
RETURN

33. Metallic Bonding - “Electron Sea”
B. Types of Bonds
Metallic Bonding - “Electron Sea”
RETURN

34. C. Bond Polarity
Most bonds are a blend of ionic and covalent characteristics.
Difference in electronegativity determines bond type.

35. C. Bond Polarity Electronegativity
Attraction an atom has for a shared pair of electrons.
higher e-neg atom  -
lower e-neg atom +

36. C. Bond Polarity Electronegativity Trend (p. 151)
Increases up and to the right.

37. C. Bond Polarity Nonpolar Covalent Bond e- are shared equally
symmetrical e- density
usually identical atoms

38. + - C. Bond Polarity Polar Covalent Bond e- are shared unequally
asymmetrical e- density
results in partial charges (dipole) + -

39. C. Bond Polarity
Nonpolar
Polar
Ionic
View Bonding Animations.

40. C. Bond Polarity Examples: Cl2 HCl 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9
NaCl
=0.0
Nonpolar
=0.9
Polar
=2.1
Ionic

41. 6.5 The properties of molecular compounds

42. -VSEPR Theory
*electrostatic repulsion between valence e- pairs surrounding an atom causes these pairs to be oriented as far apart as possible

43. (table 6-5 pg 186)
Linear
Tetrahedral
Trigonal pyramidal
Bent

44. -Hybridization(fig.6-23 p188)
*mixing of two or more atomic orbitals of similar nrg’s on the same atom to give new orbital of equal nrg’s

45. -Hybridization
ex: CH4
-hybridization of s and p yield 4 new identical orbits
*hybrid orbitals: orb of = nrg produced by the combination of 2 or more orb on the same atom

46. sp2 sp3 sp Trigonal-planar (BF3): triangle shape
Linear (CO2) : any double or triple bond
sp2
Trigonal-planar (BF3): triangle shape
sp3
Tetrahedral (CH4) : plus looking sign

47. -Intermolecular Forces
*force of attraction between molecules
-generally weaker than the bonds that join atoms in molecules, ions in ionic cmpnds, or metal atoms in solid metals

48. ~dipole-dipole forces and molecular polarity
-strongest intermol forces act between polar molecules
*dipole: = but opp charges separated by a short distance.
- arrow points toward the neg pole

49. *DD force: forces of attraction between polar molecules
ex: polar vs nonpolar

50. ~hydrogen bonding
*an intermolecular attraction between a H atom and an unshared pair of e-’s on a strongly electroneg. Atom

51. ~hydrogen bonding
-H bond represented by dotted lines connecting the H to the unshared e- pair of the electroneg atom

52. ~london dispersion forces
(Fritz London 1930)
*attraction resulting from the constant motion of e-’s and the certain instantaneous dipoles and induced dipoles

53. ~london dispersion forces
-between all atoms and molecules
-only intermolecular forces acting among noble-gas atoms and nonpolar molecules
-LF increases with increasing mass of atoms and molecules