1. Chapter : Chemical Bonding
Cartoon courtesy of NearingZero.net

3. Chemical Bonds Lect1
Forces that hold groups of atoms together and make them function as a unit.
Ionic bonds – transfer of electrons Metal + Nonmetal Ex) NaCl Li2O
Covalent bonds – sharing of electrons. 2 nonmetals Ex) H2O CO2
Metallic bonds- electrons are free to move throughout the material. Metals

4. Covalent Bonds Nonpolar-Covalent bonds (H2)
Electrons are equally shared
Electronegativity difference of 0 to 0.3
Polar-Covalent bonds (HCl)
Electrons are unequally shared
Electronegativity difference between .3 and 1.7

5. Using Electronegativity differences
CO = 0.8
C= look on table pg 198
O= thru 3.3= Ionic Bond

6. Covalent Bonding
Molecule- is the smallest unit quanitity of matter which can exist by itself and retain all the properties of the original substance.
Examples H2O & O2
Diatomic molecule- is a molecule containing 2 identical atoms. (H2 N2 O2 F2 Cl2 Br2 I2)
H NO F

7. Chemical Formula- represents the relative # of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. Example: H2O H=2 O=1
Molecular compound (Covalent Compounds) - simplest formula unit are molecules. Have low melting & boiling pts.
Molecular formula- shows the types and numbers of atoms combined in a single molecule.

8. Empirical formula
Shows the lowest, simplified ratio of elements in a compound
Molecule
Molecular formula
Empirical formula 1
C2H4
CH2 2
C4H8 3
C3H8 ?

9. Naming Covalent Compounds
Two words, with prefixes
Prefixes tell you how many.
mono, di, tri, tetra, penta, hexa, septa, octa, nona, deca
First element whole name with the appropriate prefix, except mono
Second element, -ide ending with appropriate prefix
Practice

10. Naming Covalent Compounds
CO Carbon Dioxide
CO Carbon Monoxide
CCl Carbon terachloride
N2O Dinitrogen tetraoxide
XeF Xenon hexaflouride
P2O9 Diphosphorus Nonaoxide
H2O Dihydrogen Monoxide

11. Covalent compounds The name tells you how to write the formula duh
Sulfur dioxide SO2
diflourine monoxide F2O
nitrogen trichloride NCl3
diphosphorus pentoxide P2O5

12. Bond Length- is the average distance between 2 bonded atoms.
Bond Energy- is the energy required to break a bond.
It gives us information about the strength of a bonding interaction.

13. I. Lewis Diagrams (p. 202 – 213)
Lecture 2

14. The Octet Rule
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an (8) octet of electrons in its valence shell.
8 is Great! H He

15. Lewis Dot
Shows how valence electrons are arranged among atoms in a molecule.
Reflects central idea that stability of a compound relates to noble gas electron configuration.
CH4
H2O
I started it for you
Try one yourself NH3

16. A. Octet Rule
Remember…
Most atoms form bonds in order to have 8 valence electrons.

17. H O H N O F B F F F F F S F A. Octet Rule Very unstable!!
Exceptions:
Hydrogen  2 valence e-
Groups 1,2,3 get 2,4,6 valence e-
Expanded octet  more than 8 valence e- (e.g. S, P, Xe)
Radicals  odd # of valence e-

18. B. Drawing Lewis Diagrams
Find total # of valence e-.
Arrange atoms - singular atom is usually in the middle.
Form bonds between atoms (2 e-).
Distribute remaining e- to give each atom an octet (recall exceptions).
If there aren’t enough e- to go around, form double or triple bonds.

19. B. Drawing Lewis Diagrams
CF4
1 C × 4e- = 4e-
4 F × 7e- = 28e-
32e- F
F C F
- 8e-
24e-

20. B. Drawing Lewis Diagrams
BeCl2
1 Be × 2e- = 2e-
2 Cl × 7e- = 14e-
16e-
Cl Be Cl
- 4e-
12e-

21. B. Drawing Lewis Diagrams
CO2
1 C × 4e- = 4e-
2 O × 6e- = 12e-
16e-
O C O
- 4e-
12e-

22. C. Polyatomic Ions To find total # of valence e-:
Add 1e- for each negative charge.
Subtract 1e- for each positive charge.
Place brackets around the ion and label the charge.

23. O O Cl O C. Polyatomic Ions ClO4- 1 Cl × 7e- = 7e- 4 O × 6e- = 24e-

24. H H N H C. Polyatomic Ions NH4+ 1 N × 5e- = 5e- 4 H × 1e- = 4e- 9e-

25. D. Resonance Structures
Molecules that can’t be correctly represented by a single Lewis diagram.
Actual structure is an average of all the possibilities.
Show possible structures separated by a double-headed arrow.

26. D. Resonance Structures
O S O O
O S O O
O S O

27. D. Resonance-Occurs when more than one valid Lewis structure can be written for a particular molecule.

28. Structural Formula
Shows shared pair of electrons by a dashed line.

29. Single bond- 1 pair of electrons
Double bond- 2 pair of electrons
Triple bond- 3 pair of electrons
Try a couple: O2 N2

30. Ionic Bonding and Compounds
Ionic Compound- is composed of positive and negative ions combined so that the positive & negative charges are equal. (Metal + nonmetal)
Formula Unit- is the simplest collection of atoms from which a compounds formula can be established.

31. Writing Formulas Write the symbols of each element
Put their charge in their upper right corner
Crisscross the numbers down (Not the charges).
Example:
Write the formula for Magnesium Chloride
Mg Cl
Mg+2 Cl-1
MgCl2
MgCl2

32. Writing Formulas Practice+1 +2 -1
Write the formula for:
Aluminum Bromide
Calcium Oxide
Calcium Nitride
Sodium Chloride

33. Lattice Energy- is the energy released when 1 mole of an ionic crystalline compound is formed from gaseous ions.
Ionic Compounds have high melting points & boiling points, are hard and brittle, have crystalline structure.

34. Polyatomic Ions
Many atoms with a charge.
Example SO4-2

35. Metallic Bonding
Metals- conduct heat, have low ionization energy & electronegativity, give up e-
Metallic Bond- is a chemical bond resulting from the attraction between positive ions and surrounding mobile e-.
Malleability and ductility

36. Molecular Geometry
VSEPR Theory- “Valence- shell, electron-pair repulsion”
states that repulsion between the sets of valence-level electrons surrounding an atom cause these sets to be oriented as far apart as possible.

37. Determining VSEPR H-O-H E2 B B A Determine the VSEPR for H2O
Draw the Lewis Dot
Draw the Structural Formula
Label the central atom as A
Label any atoms attached to the center atom as B
Label any paired electrons on the central atom that are not used in the bond as E
H-O-H E2 B B A
VSEPR AB2E2
Shape Bent (look on chart)

38. VSEPR Chart VSEPR SHAPE AB or AB2 Linear AB2E Bent AB3 Trigonal-Planar
Tetrahedral
AB3E
Trigonal-Pyramidal
AB2E2
AB5
Trigonal-Bipyramidal
AB6
Octahedral

39. Hybridization-The Blending of Orbitals.
Dipole- is created by equal but opposite charges that are separated by a short distance.
Dipole-Dipole Attractions-Attraction between oppositely charged regions of neighboring molecules.
Hydrogen Bonding- Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen. Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests.
London Dispersion Forces- The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. London forces increase with the size of the molecules.

40. “Electronegativity chart”. Table. Aug. 9, 2006. http://www
“Lewis Structures”. Drawings. Aug. 9,
“Oscar”. Photo. Aug. 9,
“Water Structural Formula”. Drawing. Aug. 10,
“Periodic Table of Elements”. Chart. Aug. 9,
“Information”. Aug 11,
Holt, Rinehart and Winston. Modern Chemistry. Harcourt Brace & Company