1. Chapter 7 Chemical Formulas and Compounds
Nomenclature
Systematic way of writing and naming compounds.
Purpose
So we don’t have to memorize all the common names.

2. 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.

3. Criss Cross Method Binary Ionic
Ba + I
BaI2

4. 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)

5. 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.

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

7. 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

8. 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

9. Naming Binary Ionic PbS
Naming ionic compounds with multiple charged metals.
Use roman numerals after the metal to indicate the charge of the metal.
Cr2O3
PbS

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

11. 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

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

13. 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.

14. 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

15. 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

16. 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.

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

18. Naming Molecular Compounds
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.

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

20. 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

21. 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

22. 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

23. 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

24. 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

25. 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

26. 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

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

28. 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.

29. 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.

30. 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 ( ).

31. Chapter 7.2 Oxidation Numbers
To show the general distribution of electrons among the bonded atoms in a molecule or polyatomic ion, oxidation states are assigned to the atoms in the compound or ion.
Similar to ionic charges, but they don’t have an exact physical meaning.
Useful when explaining chemical reactions and balanced equations.

32. Oxidation Rules The oxidation number of a free element is always 0.
The oxidation number of a monatomic ion equals the charge of the ion.  
The usual oxidation number of hydrogen is +1 with nonmetals and -1 with metals.
The oxidation number of oxygen in compounds is usually -2. Unless its peroxide which will be -1 or with Fluorine.  
Fluorine is always -1, highest electronegative element.
The sum of the oxidation numbers of all of the atoms in a neutral compound is 0.  
The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion.

33. Rules for Oxidation Numbers
Page 232 list 8 rules you should follow.
Note that all elements in the pure state have oxidation numbers equal to zero.
Identify the oxidation states the practice problems page 234.

34. Stock system over Prefix naming
Instead of using the prefixes mono to deca, the oxidation numbers can be used as Roman numerals like the stock naming system.
Ex.
SO2 – sulfur dioxide
Sulfur (IV) Oxide
SO3 – sulfur trioxide
Sulfur (VI) Oxide

35. 7.3 Using Chemical Formulas
Atomic Mass
The average mass of all naturally occurring isotopes of an element.
Formula Mass
Sum of the average atomic masses for all atoms in a molecule or formula unit.

36. Calculating Formula Mass
Determine the formula mass of acetic acid, HC2H3O2.
Use the periodic table to find the atomic mass of each element. Round the mass to the nearest whole number, with the exceptions of Copper and Chlorine.
Cu = 63.5amu Cl = 35.5amu

37. Practice: Formula Mass
Determine the formula mass for the following:
H2SO4
(NH4)2CO3
CaClO3

38. Molar Mass Mass in grams of 1 mole of any substance.
Equivalent to the formula mass of a compound and to the atomic mass of an element.
1 mol of S = 32g S
1 mol of H2SO4 = 98g H2SO4

39. Determining Molar Mass
Determine the formula mass and the molar mass of C6H12O6.
C: 6 x 12amu = 72amu
H: 12 x 1amu = 12amu
O: 6 x 16amu = 96amu +
Formula mass = 180 amu
Molar mass = 180 g/mol

40. Molar Conversions 1 mol = Formula mass = Molar Mass
1 mol = 6.02 x particles
The mole is the central unit in converting the amount of substances in chemistry.

41. Mass – Mole Conversions
Use 1 mol = Formula Mass conversions.
Mass to Moles Conversion
How many moles are in 250.g of NaCl?
First determine the molar mass of NaCl
Moles to Mass Conversion
How many grams are in .55 moles of CO2?
First determine the molar mass of CO2.

42. Mole – Particle Conversion
Use 1 mol = 6.02 x particles
Particle to Mole Conversion
How many moles are equivalent to 550. molecules of SO3?
Mole to Particle Conversion
How many atoms are in .525 moles of Ca?

43. Molar Conversions 1mol = Formula mass(g) 1mol = 6.02 x 1023 particles
Atoms
Molecules
Formula Units(f.u.)
Determine the number of particles in 45.0g of Ca(NO3)2.

44. Percent Composition
The percent by mass of each element in a compound.
Example: Determine the % composition of each element in NaCl.

45. Practice % Composition
Determine the % composition of each element in Na2SO4.

46. Hydrates Hydrates – compounds that contain water.
Example : MgSO4  5H2O
The ()really is equivalent to (+).
Determine the formula mass of this compound.

47. % composition of hydrates
Determine the % of water in MgSO4  5H2O.
Determine the % of water in CaCO3  10H2O.

48. % Composition of Data
Determine the percent composition of a 5.50g compound that contains sulfur and 3.30g of silver.
1st determine the amount of sulfur:
(5.50g Ag + S) – (3.30g Ag)= 2.20g S

49. Using Percent composition.
Determine the mass of an element in a given quantity of a compound.
Multiply the % of the element in decimal form by the total mass of the compound.
Example:
Determine the mass of Ag in 120g of AgCl.

50. Practice % by mass 1. Determine the mass of Ca, in 400g of CaCO3.
2. Determine the mass of O, in 125g of H2SO4.
3. Determine the mass of K, in 50g of KClO34H2O.

51. Determining Chemical formulas
Empirical Formula (Simplest Formula)– Smallest whole number ratio of atoms in a chemical formula.

52. Calculating an Empirical Formula:
1. Take the given % and change to grams.
ex % = 25.65g
2. Convert each value to moles. (divide by molar mass)
3. Divide all answers by the smallest value.
4. If all answers are whole numbers, round to the nearest (+/- .1 ). Write the formula with those subscripts.
5. If not, then multiply each by a whole number to give a whole number value for all elements.
ex. 1.5 x 2 = 3, or 1.4 x 5 = 7

53. Sample Problem
Determine the simplest formula for a compound that contains 26.56% K, 35.41% Cr, 38.08% O.

54. Practice Problem
Find the simplest formula for a compound that contains 32.38% Na, 22.65% S, 44.99% O.

55. Calculating the Molecular Formulas
Molecular Formula – is a multiple of an empirical formula.
Molecular mass is always in the given problem.
Molecular Formula = (Empirical Formula)x

56. Sample Problem
Determine the molecular formula of a compound having a simplest formula of CH and a formula mass of g
CH, empirical mass = 13g

57. Sample Problem #2
A compound with a formula mass of 34g is found to consist of .44g H and 6.92g O. What is its molecular formula?
1st determine the empirical formula and its mass.