1. Chapter 7 Chemical Formulas and Bonds
7-1 Ionic Bonding
Octet Rule
Atoms in compounds tend to have a noble gas configuration. Typically 8 valence electrons.
Ionic Bonds
Force of attraction between oppositely charged ions.
Formula units(f.u.) – chemical formula for an ionic compound.

2. Cations / Anions Cations Anions Positively charged ions
Drops to lower energy level
Noble gas or pseudo noble gas configurations
Transition and Group 4A elements have multiple charges except Zinc and Silver. (Zn 2+ and Ag 1+ )
Anions
Negatively charged ions
Fills the valence shell
Noble gas configuration
Halide, ions formed from the halogen atoms.

3. Lewis Dot Diagrams
Using electron dot notation, sketch the exchange of valence electrons in atoms as they form ionic bonds.
Ca + Cl
Na + S

4. Class work
Show the electron transfer for the following elements and write the formula unit for the binary ionic compound:
1) Zn + I
2) Al + S
3) K + P
4) Mg + N
5) Ba + Se
6) Fe O

5. Chapter 7.2 Covalent Bonding
Force of attraction for electrons, that results in a pair of electrons being shared by two atoms.
Molecule (Molecular Substance)
Group of atoms that are held together by a covalent bond.
Molecular Formula
Describes the composition of atoms in a single molecule of a compound.
Number and type of each atom in the molecule.

6. Structural Formula Lewis structures
Formula that specifies which atoms are bonded to each other in a molecule.
Places atoms in a pattern in which they can share electrons to satisfy the octet rule.
Every shared pair of electrons are depicted as a dash ( — ).
Shared Pair, are 2 electrons that are shared between 2 atoms. (bonding pairs)
All unshared electrons must be paired.
Unshared Pair, are 2 electrons that are NOT shared between 2 atoms. (non-bonding pairs)

7. Single Covalent Bonds
Draw the Lewis Structure for the molecule that has 1 N and 3 H.

8. Systematic over Trial-Error
Count the total number of valence electrons (N) needed to account for the atoms (based on the column of the atom in the periodic table) and charge (add one electrons for each negative charge; subtract one electron for each positive charge).
Draw the framework with single bonds. Some knowledge of the way the atoms are connected may be required.
Using lone pairs, complete octets around the noncentral atoms.
Count the number of electrons depicted (two for each bond and two for each lone pair). If this number is less than N, then add electrons to the central atom until the total number of electrons depicted is N.
If the octet rule is not satisfied for the central atom and lone-pair electrons are nearby, use those electrons to make double or triple bonds to the central atom.

9. Practice Single Covalent Bonds
Write the Lewis Structure for the following formulas:
Cl2
OF2
C2H6
C3H7Cl

10. Multiple Covalent Bonds
Some structures need more than a single pair of electrons to reach the octet.
Double covalent bonds
Covalent bond between 2 atoms in which there are a total of 4 electrons being shared. (2pair)
Triple covalent bonds
Covalent bond between 2 atoms in which there are a total of 6 electrons being shared. (3pair)
CO2 and HCN

11. Polyatomic Ion Structures
Polyatomic ions, are covalently bonded atoms that form a charge due to the gain or loss of electrons.
When drawing the structures of polyatomic ions:
(–) ions must have extra electrons in the structure that is equivalent to its charge.
(+) ions must lose the number of electrons in the structure that is equivalent to its charge.

12. Example
NH4 1+, the ammonium ion.
SO4 2-, sulfate ion.

13. Classwork Lewis StructureCO
SO3
ClO3 1-
NO2 1-

14. Exceptions to the Octet
Incomplete Octet
Some atoms that have less than eight electrons.
Hydrogen only needs two and Boron needs six electrons.
Expanded Octet
Atoms that can have more than eight electrons.
Common expanded octets:
Cl, Br, I, S, and P
IMPORTANT!!!
Only one atom may exceed the standard eight valence.
Must be the central atom only!!

15. Draw the Lewis Structure for the following:
BF3
IF4 1-

16. Coordinate Covalent Bond
Like a single covalent bond, but a single atom is sharing 2 electrons with another atom, in which it doesn’t increase the # of electrons in the atom sharing the 2 electrons.
Helps to show ownership of electrons.
Indicated by using an arrow rather than a dashed line. () The arrow must point away from the atom that has less electrons than what it should.

17. Practice
SO3
NO3 1-

18. Polar Vs. Nonpolar Covalent Bonds
Non-Polar – electrons are equally shared
Balanced Distribution of Electrical Charge
Polar – electrons are not equally shared
Unbalanced Distribution of Electrical Charge
Partial Charge
 +
 -

19. Determining Bond Type
The Difference in electronegativity determines the bond type, page 241, according to the following scale:
Example Na-Cl
Na = 0.9 Cl = 3.0 Difference = 3.0 – 0.9 = 2.1
Ionic
2.0
4.0
0.4
Ionic
Polar Covalent
Non

20. Identifying Bond Types
Determine if the following are going to form; Ionic, Polar Covalent or Non-polar Covalent Bonds:
C-N
Ca-F
Br-Br
H-Br

21. 7.3 Nomenclature Systematic way of writing and naming compounds.
Purpose
So we don’t have to memorize all the common names.

22. Formula Writing for Ionic Compounds
Binary Ionic Compounds
Ionic Bond between two atoms.
Total charge of all compounds must equal zero!
Total (+) = Total (-)
Identify the charges of both + and – ions.
If charges equal no subscripts needed.
If charges are not equal, cross the charges without (+/-) signs as subscripts.
If the subscripts can be reduced, reduce to the lowest whole number ratio.

23. Criss Cross Method Binary Ionic
Ba + I
BaI2

24. Practice Formulas
Write the formula unit for the following Binary Ionic Compounds:
(Zn + O)
(Fe 3+ + Cl)
(Ba + N)
(Sn 4+ + S)
(Cu 1+ + P)

25. Writing Ternary Ionic Compounds
Polyatomic Ions
Combinations of 2 or more non-metals that form common ions.
Most are negatively charged except NH4 1+
Must be placed in ( ) if more than 2 are needed.
Ternary Ionic Compounds
Ionic bond between an element and a polyatomic ion.

26. Criss Cross Method Ternary Ionic
Na and PO4 3-
Na3PO4

27. Practice Writing Ternary Ionic
Write the formula unit for the following Ternary Ionic Compounds:
(Zn + OH-)
(Fe 3+ + ClO3 1- )
(Ba + NO3 1- )
(Sn 2+ + SO3 2- )
(NH P)

28. Naming Ionic Compounds
Binary Ionic Compounds
Metals, retain the name of the element.
Multiple charged metals must use either stock or classical naming system. (STOCK is preferred)
Non-metals, root of name + ide ending.
Chlorine = Chloride
Nitrogen = Nitride Phosphorus = Phosphide
Sulfur = Sulfide Carbon = Carbide
Oxygen = Oxide Selenium = Selenide
Example
MgS
Magnesium Sulfide

29. Naming Binary Ionic
Naming ionic compounds with multiple charged metals.
Use roman numerals after the metal to indicate the charge of the metal.
Cr2O3
Chromium (III) Oxide
PbS
Lead (II) Sulfide

30. Practice Naming Binary Ionic
NaCl
FeCl2
CaF2 KI
Al2O3
SnO

31. Naming Ternary Ionic Compounds
Same rules as the binary ionic compounds.
Except look up the name of the polyatomic ion, on the ion sheet given, and write the name as it appears.
Example
NaNO3
Sodium Nitrate

32. Practice Naming Ternary
Ca3(PO4)2
AgNO2
Zn(OH)2
KC2H3O2

33. Summary of Writing Ionic Compounds
Identify the ending of the compound.
-ide, check to see if the first word is NOT ammonium or the last is NOT Hydroxide or Cyanide.
If not, then it’s a binary ionic compound.
If yes, then it’s a ternary ionic compound.
-ate, or –Ite
Ternary Ionic Compound.

34. Summary of Naming Ionic Compounds
Identify the number of elements.
2 elements, then it’s a binary ionic:
Name the metal + nonmetal –ide
3 or more elements, then it’s a Ternary ionic:
Name the metal + polyion OR
Name the polyion + nonmetal –ide OR
Name the polyion + polyion

35. Practice Write formulas for the following compounds:
Silver Sulfide
Magnesium Nitrate
Copper(II) Nitride
Chromium(III) Sulfite
Name the following compounds:
K2SO4
Fe(NO3)3
Ca3P2
SnO2

36. Writing/Naming Molecular Compounds
Mono- 1
Di- 2
Tri- 3
Tetra- 4
Penta- 5
Hexa- 6
Hepta- 7
Octa- 8
Nona- 9
Deca- 10
Molecular Compounds
Contain covalent bonds.
Writing Formulas
Write the symbol for the name of each atom.
Use a subscript equivalent to the meaning of the prefix.
Table of Prefixes, page 246.

37. Example Write the formula for Carbon Monoxide.CO
No prefix on the first word, indicates 1 atom.
Prefix mono-, indicate 1 atom
(Used only on second word)
Write the formula for Carbon Dioxide.
CO2

38. Practice Writing Formulas
Write the formulas for the following compounds:
Dinitrogen Pentoxide
Phosphorus Trichloride
Sulfur Trioxide
Carbon Tetrachloride

39. Naming Molecular Compounds
Use the same prefixes on page 246.
First word
Prefix + Name of element, only uses a prefix if there are 2 or more atoms.
Second word
MUST HAVE a prefix followed by name ending in –ide.

40. Practice Naming Write the name for the following formulas: P4O10 N2O5
CF4
N2O3
SO2

41. Binary Acids / Oxyacids
All acids contain the element H.
H is always written 1st in any formula representing acids.
Binary Acids
Contain Hydrogen + a non-metal.
Oxyacids
Contain Hydrogen + a polyatomic ion

42. Binary Acids Naming: Example Must begin with the prefix Hydro-
Must end in the suffix –ic.
Place the nonmetal between the prefix and suffix.
Example
HCl
Hydrochloric Acid

43. Binary Acids Writing: Example Always place the H first.
Identify the nonmetal in the acid.
Cross the charges on the H 1+ and the nonmetal as subscripts.
Example
Hydrosulfuric Acid
H +1 S 2-
H2S

44. Binary Acids Practice Write the name for each of the following acids:HI
H2S
H3N
Write the formula for the following acids:
Hydrofluoric acid
Hydrophosphoric acid
Hydroselenic acid

45. Oxyacids Naming: Example Identify the polyatomic ion after the H.
Write the name of the non-metal in the polyatomic ion.
If the polyatomic ion ends in:
-ate  change to –ic
-ite  change to –ous
Example
H2SO4
Sulfate  Sulfuric Acid
H2SO3
Sulfite  Sulfurous Acid

46. Oxyacids Writing: Example Identify the root of the polyatomic ion.
If the acid ends in:
- ic  use the –ate ending
- ous  use the –ite ending
Place the H in front of the polyatomic ion and cross the charges if needed.
Example
Phosphoric Acid
Phosphoric = Phosphate (PO4) 3-
H3PO4

47. Oxyacids Practice Write the name for the following acids:
HNO3
H2CO3
HC2H3O2
Write the formula for the following acids:
Nitrous acid
Perchloric acid
Phosphorous acid

48. Hydrates Ionic compounds that contain absorbed water. Example
MgSO4  5H2O
Magnesium Sulfate Pentahydrate

49. Summary of Naming Compounds
1st element H, then its an acid:
2 or more elements after the H
Oxyacid
1 element after the H
Binary acid
2 elements, 1st element is NOT hydrogen.
Both are non-metals
Molecular compound, use prefixes.
1st is a metal
Binary Ionic, check to see if the 1st element is multiple charged.
3 or more elements, 1st element is NOT hydrogen.
Ternary Ionic
Contains at least 1 polyatomic ion.
1st element a metal, check to see if it is a multiple charge.

50. Summary of Writing Formulas
Most formulas use only 2 words.
Check the 2nd word ending first!
Ending is –ic or –ous
Compound is an acid.
Has –ic ending or –ous without a Hydro- prefix
Oxyacid
Contains a Hydrogen + Polyatomic Ion.
Has –ic ending with a Hydro- prefix
Binary Acid
Contains a Hydrogen + Non-metal.

51. Writing Formulas Continued
Ending is –ide.
If it ends in Hydroxide or Cyanide.
Ternary Ionic, check charges and cross.
Reminder!! More than 1 polyatomic ion, use ( ).
If it isn’t Hydroxide or Cyanide.
Binary Compound
Check if prefixes are used on the 2nd word.
Prefix  then Binary molecular, just write subscripts = to prefixes.
No Prefix  Binary Ionic, check charges and cross if needed.
If it ends in –ite or –ate.
More than 1 polyatomic ion needed, use ( ).