1. Chemical Bonds, Nomenclature, Lewis Structure and Molecular Shapes
Chapter 5
Chemical Bonds, Nomenclature, Lewis Structure and Molecular Shapes

2. Homework & Quizzes – Chapter 5
Text Homework (not turned in): pages 147 – Problems: 1, 6 – 8, 17, 27, 35, 39, 41 – 66, 68, 69, 72, 86, 88(b&c), 107, 109, 110, 112, 113.
Quiz: Do the graded quiz in Blackboard.

3. I. Chemical Bonds A. Introduction (summary of chapter)
Atoms can combine to produce new larger units called molecules or compounds.
Each molecule has a unique name (two rules to learn).
Molecules held together by chemical bonds (two types).
Bonds result from either transfer of valence electrons (Ionic Bonds) or from sharing of valence electrons – Covalent Bonds.
Valence electrons rearrange to mimic closest Group VIIIA (18) structure.
Molecules resulting from covalent bonding will have predictable shapes.

4. I. Chemical Bonds B. Ionic Bonds
- Metals (except H) loose electrons, form cations; Nonmetals gain electrons to form anions. Both strive for e- configuration of nearest inert gas.
- The resulting opposite ions attract in a ratio which produces a neutral unit. Reduce formula to simplest ratio.
Ionic Bond Definition: bond formed by electrostatic attraction between anions (-) and cations (+).
Write formula with + element first; do not show charges; Final compound is neutral.
Generality: any compound formed from metallic and nonmetallic elements is ionic.

5. I. Chemical Bonds B. Ionic Bonds
Know: Metals combine with nonmetals & form ionic bonds by losing or gaining electrons to mimic closest Inert Gas (VIIIA).
IA Na, K, Li, etc become +1 ions: Na+
IIA - Ca, Mg, etc become +2 ions: Ca+2
IIIA - Al, Ga become +3 ions: Al+3
VA - N, P become -3 ions: N-3
VIA - O, S become -2 ions: O-2
VIIA - F, Cl, Br, I become -1 ions: F-1
Opposite ions attract in a ratio so that the product is neutral.
Inert Gas e- Configurations

6. I. Chemical Bonds B. Ionic Bonds Example
Do not show charges in final formula. NaCl NOT Na+Cl-

7. I. Chemical Bonds B. Ionic Bonds Example

8. I. Chemical Bonds B. Ionic Bonds - Examples
Na+ + Na O Na2O
Ca F- + F CaF2
Mg S MgS
Al+3 + Al O-2 + O-2 + O-2 Al2O3
Give the formulas for the following: Na & Br
Ca & O Ba & I Li & O Al & F Mg & N
Many transition metals form ionic bonds & can have several charges such as Fe+2 = Iron (II); Fe+3 = Iron (III); Cu+2 = Copper (II); Cu+1 = Copper (I)

9. I. Chemical Bonds C. Electron Dot (Lewis) Structures
- A Lewis electron dot structure is a symbol in which the valence electrons are shown as dots.
- Examples: Na Mg: Na Ca2+
H:1- (Called Hydride) :C: :Si:
- How many valence electrons (dots) would
N O F- or Ne have? What about Mg+2?

10. II. Covalent Bonds A. Introduction EN = electronegativity
- Definition of a covalent bond: A bond formed by the sharing of two electrons.
When two atoms of similar EN combine, neither has the “pull” to take electrons away & a sharing of electrons results.
This occurs when NONMETALS, including H, combine with NONMETALS.
Example: H. + H ) H—H = H2
- The atoms share valence electrons to get stable group VIIIA e- configurations.

11. II. Covalent Bonds A. Introduction
Covalent bond = sharing of 2 electrons.
2 shared electrons with (Single Bond).
4 shared electrons with (Double Bond).
6 shared electrons with (Triple Bond).
We frequently show the structure as a Lewis Structure - covalent bonds with lines and nonbonding valence electrons as dots.
- Note: Group IVA usually forms 4 bonds; VA three bonds; VIA two bonds; and VIIA (along with H) one bond.

12. II. Covalent Bonds B. Examples
H F::: ---) H F:::
H. + O H ) H O H
:N + N: ---) :N N:
:::Cl. + .O. + .Cl::: ---) :::Cl O Cl:::
::O: + :C: + :O:: ) ::O = C = O::

13. II. Covalent Bonds C. Lewis Structures 1
II. Covalent Bonds C. Lewis Structures 1. Rules for Drawing Lewis Structures
1. Calculate the total # of valence electrons; take into account charge if the sample is an ion.
2. Place atom that forms most bonds at center (Closest to Group IVA & Lowest if in same group). If there is a charge, then add or subtract the appropriate number of electrons on the central atom.
3. Arrange other atoms around central atom & allow sharing so that each atom has stable electron configuration. Show bonding pairs as dashes & nonbonding valence e- as dots.
4. Double check: a) each atom has a stable electron configuration & b) have the same total number of valence electrons as in step 1.

14. II. Covalent Bonds C. Lewis Structures 2. Examples
H I H2O NH4+
H2O CH4 SO2
AlCl NO CN-

15. Bonding Summary Two General Bonding Types
Ionic: Compound containing metallic element. Atoms lose/gain e to look like nearest inert gas. Add together ions such that neutralize charge.
Ia IIa IIIa Va VIa VIIa
2. Covalent: Compound containing nonmetals.
Atoms obtain inert gas configuration by
sharing valence electrons. : :: :::

16. II. Covalent Bonds – Organic Compounds
Can write organic structures several ways.
Example – Butane (Note the five ways of presenting)
Note: Carbon always has four bonds.
C4H CH3CH2CH2CH CH3-CH2-CH2-CH3
H H H H
H – C – C – C – C – H
H H H H

17. II. Covalent Bonds – Organic Compounds
Cyclic Organics: Example of Cyclopropane
Aromatics Contain Benzene, C6H6

18. II. Covalent Bonds – Organic Compounds
C7H6O3 MW = 138g
C9H8O4 MW = 180 g

19. II. Covalent Bonds – Organic Compounds – Aspirin Lab
1) Equation & Conversion Factors:
1 Salicylic Acid Acetic Anhydride ) 1 Aspirin Acetic Acid
1 = molecules or moles; 1 mole = formula weight in grams = 6.0x1023 molecules
2) Lab Calculations (questions 2 & 3):
2.0 g SA x 1 mole SA = mole SA
138 g SA
0.014 mole SA x 1 mole Aspirin = mole Aspirin
1 mole SA
From the coefficients in the balanced chemical equation above.

20. III. Shapes Molecular Shapes play a major role in:
1) Physical Properties
2) Chemical Properties
3) Biochemical Properties
To Obtain the shape of a molecule one draws the Lewis Structure, counts the number of “things” around the central atom, and uses simple geometry to predict the shape.

21. III. Shapes C. Simplified Examples
Bond angle = 180o
Bond angle = 120o
Bond angle = 109o

22. IV. Nomenclature A. Introductions
There are common & systematic names for chemicals. A chemical may have scores of common names.
A systematic name must allow one to both obtain the formula and derive the name from the formula.
There are two general rules for naming inorganic compounds.
Ionic compounds use Rule #1. Molecular or Covalent compounds use Rule #2.

23. IV. Nomenclature B. Ionic Compounds
Rule #1 for ionic compounds: Name the + element, then the – element and change the ending to “ide.”
Examples:
NaCl = Sodium Chloride
Na2O = Sodium Oxide

24. IV. Nomenclature B. Ionic Compounds Rule #1 – “ide” names
Negative atoms have an “ide” ending.
Atom Anion Name
Chlorine Cl1- Chloride
Oxygen O2- Oxide
Fluorine F1- Fluoride
Sulfur S2- Sulfide
Nitrogen N3- Nitride
Iodine I1- Iodide
Bromine Br1- Bromide
Phosphorus P3- Phosphide

25. IV. Nomenclature B. Ionic Compounds Examples
NaCl
Na2O
AlF3
Be3N2
Calcium Sulfide
Barium Iodide
Barium Oxide
Magnesium Nitride
Sodium Chloride
Sodium Oxide
Aluminum Fluoride
Beryllium Nitride
CaS
BaI2
BaO
Mg3N2

26. IV. Nomenclature C. Molecular Compounds Rule #2
When nonmetals & H combine with each other through sharing electrons (covalent bonds), they form molecules; there are no ions.
Rule #2 – When both elements are nonmetals (molecular compounds), then Name the + & the - & change ending to “ide” as before. Use prefixes of di, tri, tetra, penta, etc to tell how many of each element is present.

27. IV. Nomenclature C. Molecular Compounds
CO2 = Carbon Dioxide
CCl4 = Carbon Tetrachloride
N2O = Dinitrogen Oxide
P2S5 = Diphosphorus Pentasulfide
PBr3 = Phosporus Tribromide
BI3 = Boron Triiodide
Notes: (1) Organic compounds like CH4 use their own rules which we won’t cover.
(2) diatomic molecules named with the element name. O2 = Oxygen

28. V. Polyatomic Ions Previous compounds formed from two elements.
Frequently have compounds formed from three or four elements. When this happens, then usually have a polyatomic ion present.
Polyatomic ions: stable ions formed from two or more elements; held together by covalent bonds.
Examples:
SO4-2 = Sulfate NO2- = nitrite PO4-3 = Phosphate

29. V. Polyatomic Ions
Polyatomic ions are held together by covalent bonds, and they form ionic bonds with metals.
Examples: NaNO Na2SO Na3PO4
When have more than one polyatomic ion in a compound then use parentheses around the ion.
Examples: Na2SO Ca(NO2)2 Ca3(PO4)2
Nomenclature: Simply use the polyatomic ion name.
Example: Calcium Nitrite & Calcium Phosphate above
Need to memorize the following polyatomic ions, their names and their charges.

30. V. Polyatomic Ions - Memorize the Names, Formulas and the Charges
Formula Name Formula Name
NH Ammonium (The Only Positive One in this list)
C2H3O2- Acetate CN Cyanide
NO3- Nitrate NO Nitrite
OH- Hydroxide
HCO3- Hydrogen Carbonate
CO3-2 Carbonate Cr2O Dichromate
SO Sulfate SO Sulfite
PO4-3 Phosphate

31. V. Polyatomic Ions – Examples of Naming and Obtaining Formulas
Aluminum Hydroxide
Calcium Cyanide
Barium Sulfate
Ammonium Nitrate
Ba(OH)2
LiNO2
KNO3
NaHCO3
Al2(SO4)3
Al(OH)3
Ca(CN)2
BaSO4
NH4NO3
Barium Hydroxide
Lithium Nitrite
Potassium Nitrate
Sodium Hydrogen Carbonate
Aluminum Sulfate

32. Naming: Mixed Examples
NaF
CS2
NI3
BaI2
K3PO4
Boron Trifluoride
Sodium Sulfite
Sodium Fluoride
Carbon Disulfide
Nitrogen Triiodide
Barium Iodide
Potassium Phosphate
BF3
Na2SO3