1. Chapter 7 Chemical Formulas & Chemical Compounds
7.1 Chemical Names & Formulas

2. Ions Cation: A positive ion Mg2+, NH4+ Anion: A negative ion
Cl-, SO42-
Ionic Bonding: Force of attraction between oppositely charged ions.

3. Predicting Ionic Charges
Groups :
Many transition elements have more than one possible oxidation state.
Iron (II) = Fe2+
Iron (III) = Fe3+

4. Predicting Ionic Charges
Groups :
Some transition elements have only one possible oxidation state.
Zinc = Zn2+
Silver = Ag1+

5. Formula Writing for Binary Ionic Compounds
criss-cross the oxidation numbers to balance out the charge.
Magnesium Bromide
Mg Br – 1
Mg1Br2
MgBr2
Calcium Sulfide
Ca + 2 S – 2
Ca2S2
CaS

6. Writing Ionic Compound Formulas
Example: Iron (III) chloride
1. Write the formulas for the cation and anion, including CHARGES!
Fe3+
Cl-
2. Check to see if charges are balanced. 3
3. Balance charges, if necessary, using subscripts.
Not balanced!

7. Naming Ionic Compounds
1. Cation first, then anion
2. Monatomic cation = name of the element
Ca2+ = calcium ion
3. Monatomic anion = root + -ide
Cl- = chloride
CaCl2 = calcium chloride

8. Naming Ionic Compounds
Metals with multiple oxidation states
some metal forms more than one cation
use Roman numeral in name
PbCl2
Pb2+ is cation
PbCl2 = lead (II) chloride

9. Elements with Multiple Oxidation Numbers
Manganese II Mn+2
Manganese III Mn+3
Manganese VII Mn+7
Cobalt II Co+2
Cobalt III Co+3
Gold I Au+1
Gold III Au+3
Nickel II Ni+2
Nickel III Ni+3
Nickel IV Ni+4
**Silver Ag+1
**Zinc Zn+2
**Cadmium Cd+2
Copper I Cu+1
Copper II Cu+2
Iron II Fe+2
Iron III Fe+3
Mercury I Hg+1
Mercury II Hg+2
Lead II Pb+2
Lead IV Pb+4
Tin II Sn+2
Tin IV Sn+4
Chromium II Cr+2
Chromium III Cr+3
Chromium VI Cr+6

10. ♥ Poly Atomic Ions to Know and Love ♥
Name Formula Name Formula
Acetate
C2H3O2-1 (CH3COO -1)
Hypochlorite
ClO-1
Dichromate
Cr2O7-2
Chlorite
ClO2-1
Ammonium
NH4+1
Chlorate
ClO3-1
Nitrate
NO3-1
Perchlorate
ClO4-1
Nitrite
NO2-1
Cyanide
CN-1
Hydroxide
OH-1
Carbonate
CO3-2
Phosphate
PO4-3
Chromate
CrO4-2

11. ♥ More Poly Atomic Ions to Know and Love ♥
Name Formula Name Formula
Sulfite
SO3-2
Hydrogen Carbonate
HCO3-1
Sulfate
SO4-2
Hydrogen Phosphate
HPO4-2
Hydrogen Sulfite
HSO3-1
Hydrogen Sulfate
HSO4-1
Permanganate
MnO4-1
Oxalate
C2O4-2
Hydronium
H3O+
Silicate
SiO3-2
Peroxide
O2-2
Phosphite
PO3-3
Bromate
BrO3-1
Arsenate
AsO4-2

12. Naming Compounds with Polyatomic Ions
Formula Name
(NH4)2SO4 ammonium sulfate
ZnCO3 zinc carbonate
NH4Br ammonium bromide
Li2CO3 lithium carbonate
* Polyatomic & monatomic cation names remain the same, monatomic anions change their ending to –ide.

13. Writing Ionic Compound Formulas
Example: Barium nitrate
Ba2+
( )
NO3- 2
1. Write the formulas for the cation and anion, including CHARGES!
Not balanced!
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

14. Writing Ionic Compound Formulas
Example: Ammonium sulfate
( )
NH4+
SO42- 2
1. Write the formulas for the cation and anion, including CHARGES!
Not balanced!
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

15. Writing Ionic Compound Formulas
Example: Aluminum sulfide
Al3+
S2- 2 3
1. Write the formulas for the cation and anion, including CHARGES!
Not balanced!
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

16. Writing Ionic Compound Formulas
Example: Magnesium carbonate
Mg2+
CO32-
1. Write the formulas for the cation and anion, including CHARGES!
2. Check to see if charges are balanced.
They are balanced!
3. Simplify to a formula unit.

17. Writing Ionic Compound Formulas
Example: Zinc hydroxide
Zn2+
( )
OH- 2
1. Write the formulas for the cation and anion, including CHARGES!
Not balanced!
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

18. Writing Ionic Compound Formulas
Example: Aluminum phosphate
Al3+
PO43-
1. Write the formulas for the cation and anion, including CHARGES!
They ARE balanced!
2. Check to see if charges are balanced.

19. Chemical Formula Chemical Name
More Examples…
Chemical Formula Chemical Name
Cr2O3
Cr2O
CuSO4
Ni(OH)2
Cr2(C2O4)3
Cu2S
CuS
chromium (III) oxide
chromium (I) oxide
copper (II) sulfate
nickel (II) hydroxide
chromium (III) oxalate
copper (I) sulfide
copper (II) sulfide

20. Hydrates
Hydrate – when a water molecule (s) are chemically bonded to the ionic compound.
Normal ionic naming protocol are used, then followed by the word “hydrate.”
Prefixes are added to indicate the number of water molecules when naming hydrates.

21. Hydrate Prefixes # of water molecules prefix 1 mono- 6 hexa- 2 di- 7
hepta- 3
tri- 8
octa- 4
tetra- 9
nona- 5
penta- 10
deca-

22. Hydrates Example: MgBr2 ∙ 6H2O Magnesium bromide hexahydrate
The “ ∙ ” means “loosely bonded”
Hygroscopic - easily absorb water molecules from the air.
Deliquescent- very hygroscopic; takes out water from the air to dissolve completely to form a liquid solution.
Anhydrous – when all of the water has been removed.

23. Naming Binary Covalent Compounds
Compounds between two nonmetals
First element in the formula is named first.
Second element is named as if it were an anion.
Use prefixes
Only use mono on second element
P2O5 =
diphosphorus pentoxide
CO2 =
carbon dioxide
CO =
carbon monoxide
N2O =
dinitrogen monoxide

24. Acids always begin with Hydrogen
Anion
Formula
Name
Cl-1
HCl
Hydrochloric Acid
Br-1
HBr
Hydrobromic Acid
SO4-2
H2SO4
Sulfuric Acid
SO3-2
H2SO3
Sulfurous Acid
NO3-1
HNO3
Nitric Acid
CN-1
HCN
Hydrocyanic Acid
PO4-3
H3PO4
Phosphoric Acid

25. Bases Cation Formula Name Na+1 NaOH Sodium Hydroxide K+1 KOH
Potassium Hydroxide
NH4+1
NH3
Ammonia

26. Organic Compounds
Organic compounds are named using a different set of rules.
The simplest group is the hydrocarbons. These compounds are composed solely of the elements carbon and hydrogen.
Carbon atoms can link to each other in chains and in rings.

27. The stem of the compound name is then chosen from the following table:
Naming Hydrocarbons
The stem of the compound name is then chosen from the following table:
# of carbon atoms
prefix
# of carbon atoms 1
meth- 6
hexa- 2
eth- 7
hepta- 3
prop- 8
octa- 4
but- 9
nona- 5
penta- 10
deca-

28. Hydrocarbons: Alkanes
These molecules have the generic formula: CnH2n+2
They contain all single bonds.
CH4
methane
C2H6
ethane
C3H8
propane
C4H10
butane
C5H12
pentane
C6H14
hexane

29. Hydrocarbons: Alkenes
These molecules have the generic formula:
CnH2n
They contain double bonds between carbon atoms.
C2H4
ethene
C3H6
propene
C4H8
butene
C5H10
pentene
C6H12
hexene

30. Hydrocarbons: Alkynes
These molecules have the generic formula:
CnHn
They contain triple bonds between carbon atoms.
C2H2
ethyne
C3H3
propyne
C4H4
butyne
C5H5
pentyne
C6H6
hexyne

31. Chapter 7 Chemical Formulas & Chemical Compounds
7.2 Oxidation Numbers

32. Oxidation Numbers
Oxidation Number – numbers assigned to atoms composing a compound or ion that indicate the general distribution of electrons among bonded atoms

34. Chapter 7 Chemical Formulas & Chemical Compounds
7.3 Using Chemical Formulas

35. Molar Mass
The mass of 1 mole of a pure substance is called its Molar Mass.
Ex: Molar mass of Iron is g/mol
What is the molar mass of Platinum?
g/mol

36. Molar Mass
The molar mass depends on the particles that compose the compound. If your element exists as a molecule, i.e. BrINClHOF, one mole of these particles contains 2 moles of the element as an atom.
Determine the molar mass of oxygen molecules (O2)
(16.00 g/mol) x (2 atoms) = g/mol
The molar mass of oxygen molecules (O2) is twice the molar mass of oxygen atoms!

37. Formula Mass
The molar mass of a compound is the mass of the atomic mass units of one molecule.
This takes into consideration the number of atoms of each element in a compound.
Formula Mass is calculated the same way as molar mass except it is measured in amu, instead of g/mol.

38. Calculating Formula Mass
Calculate the formula mass of magnesium carbonate, MgCO3.
(16.00) =
84.32 amu

39. Steps for Calculating Molar Mass for Compounds
List the elements
Determine how many atoms of each
Identify the atomic masses from the periodic table
Multiply how many atoms by the respective atomic mass
Add up the totals for the Molar Mass

40. Practice H2O H 2 x 1.008 = 2.016 O 1 x 15.99 = 15.99 18.006 NaCl
Na 1 x = 22.9
Cl 1 x = 35.45
58.35
g/mol
g/mol
K2O
K 2 x = 78.2
O 1 x = 15.99
94.19
C6H12O6
C 6 x = 72.06
H x =
O 6 x = 95.94
g/mol
g/mol

41. Calculating Percentage Composition
Calculate the percentage composition of magnesium carbonate, MgCO3.
(16.00) = amu
100.00

42. *Must also find molar mass of H2O
Mass Percent
So…. In one mole of H2O, how many grams of Hydrogen are there?
2 mol H x g H = g H in 1 mol H2O 1 mol H
What % of Hydrogen, by mass, is in H2O?
2.016 g H x = % H
18 g H20
*Must also find molar mass of H2O
What % of Oxygen, by mass is in H2O?

44. Formulas molecular formula = (empirical formula)n [n = integer]
Empirical formula: the lowest whole number ratio of atoms in a compound.
Molecular formula: the true number of atoms of each element in the formula of a compound.
molecular formula = (empirical formula)n
[n = integer]
molecular formula = C6H6 = (CH)6
empirical formula = CH

45. Formulas
Formulas for ionic compounds are ALWAYS empirical (lowest whole number ratio). Often, these are called formula units.
Examples:
NaCl
MgCl2
Al2(SO4)3
K2CO3

46. Formulas
Formulas for molecular compounds MIGHT be empirical (lowest whole number ratio).
Molecular:
H2O
C6H12O6
C12H22O11
Empirical:
H2O
CH2O
C12H22O11

47. Chapter 7 Chemical Formulas & Chemical Compounds
7.4 Determining Chemical Formulas

48. Empirical Formula Determination
Base calculation on assumption of 100 grams of compound.
Determine moles of each element in 100 grams of compound.
Divide each value of moles by the smallest of the values.
Multiply each number by an integer to obtain all whole numbers.

49. Empirical Formula Determination
Adipic acid contains 49.32% C, 43.84% O, and 6.85% H by mass. What is the empirical formula of adipic acid?

50. Empirical Formula Determination (part 2)
Divide each value of moles by the smallest of the values.
Carbon:
Hydrogen:
Oxygen:

51. Empirical Formula Determination (part 3)
Multiply each number by an integer to obtain all whole numbers.
Carbon: 1.50
Hydrogen: 2.50
Oxygen: 1.00
x 2
x 2
x 2 3 5 2
Empirical formula:
C3H5O2

52. Finding the Molecular Formula
The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?
1. Find the formula mass of C3H5O2
3(12.01 g) + 5(1.01) + 2(16.00) = g

53. Finding the Molecular Formula
The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?
2. Divide the molecular mass by the mass given by the emipirical formula.
3(12.01 g) + 5(1.01) + 2(16.00) = g

54. Finding the Molecular Formula
The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?
3. Multiply the empirical formula by this number to get the molecular formula.
3(12.01 g) + 5(1.01) + 2(16.00) = g
(C3H5O2) x 2 =
C6H10O4

55. Determining Chemical Formulas from Mass Percents
A sample has been analyzed, here are the results:
18.8 % Na
29.0 % Cl
52.2 % O
How can you determine the chemical formula?

56. Step 1: Assume a 100 g sample.
Then, your percent quantities become gram (mass) quantities.
18.8 g Na , g Cl & g O
Step 2: Convert those masses to moles.
18.8 g Na x 1 mol Na = mol Na
23 g Na
29.0 g Cl x 1 mol Cl = mol Cl
35.5 g Cl
52.2 g O x 1 mol O = mol O
16 g O

57. Step 3: Since your empirical formula is in small, whole number ratios, divide your mole amounts by the smallest mole quantity.
0.817 mol Na / = mol Na
0.817 mol Cl / = mol Cl
3.26 mol O / = ≈ 4.00 mol O
Step 4: Use these values as subscripts in your formula
Na1Cl1O4 ≈ NaClO4

58. Step 5: In the event the chemical formula is not the same as the empirical formula, you need the molar mass of the desired compound and you must compare it to the molar mass of the empirical formula.
Step 6: Divide the given molar mass by the empirical molar mass to get the multiple quantity.
Step 7: Multiply each subscript in the formula by that multiple quantity.

59. Ex: MM of molecular formula = 180 g/mol
Using steps 1-4, you found that the empirical formula is CH2O.
Find the molar mass of the empirical formula: MM EF = 30 g/mol
Divide MM MF / MM EF to get a whole number.
Ex: 180 / 30 = 6
C1x6 H2x6 O1x6
C6H12O6

60. Practice Problems:
A sample has been analyzed to be % C, 0.84 % H & 89.12% Cl. Find the Empirical Formula.
A compound’s empirical formula has been determined to be HF. The compound’s molar mass is 40 g/mol. What is its chemical formula?
A compound’s empirical formula has been determined to be CH2. The compound’s molar mass is 42 g/mol. What is its chemical formula?