1. Plan, Ppt09(PS4) % Mass Empirical Formula (Tro, 3.10)
From mass data OR chemical formula (Tro, 3.9)
Empirical Formula (Tro, 3.10)
Concept/definition
From mass data OR % mass data
Note: Use “ratio of moles = ratio of FUs” idea
How to reduce a “non-whole-number” ratio to a whole-number ratio
Difference between Empirical and Molecular Formula (Tro, 3.10)
Ppt09(PS4)

2. % Mass If mass data for a particular sample are provided, use those.
If a formula is provided, assume you have a sample of one mole of the substance
Use molar masses from Periodic Table
Ppt09(PS4)

3. From Green Handout Sheet (PS4 & 5 Practice)
What is the mass percent of P in P4O10?
Pick one mole (for your sample size). It contains:
___ moles of P 4
and
4 mol P x g/mol = g P
___ moles of O 10
10 mol 0 x g/mol = g O
g/mol P4O10

4. Types of Chemical Formulas*
Chemical formula (review):
A formula in which the subscripts indicate the exact composition of a single formula unit of the substance.
Molecular formula: (MF)
Same as chemical formula in my class (though this annoys me)
Empirical formula: (EF)
A type of formula in which the subscripts are in the lowest whole number ratio.
Quick quiz: Glucose is C6H12O6 Is this an EF?
No. 6:12:6 can be reduced to _____
1:2:1
; EF = _____
CH2O
* “Formula” in a chemistry context means something completely different than a mathematical formula (e.g., D = m/V).

5. EF and MF are not always different for a substance
SO3 is both the MF and the EF for sulfur trioxide (1:3 cannot be reduced)
Na2S is the MF and the EF for sodium sulfide
But consider dinitrogen tetroxide:
Chemical formula is _____
Empirical formula is _____
N2O4
 Represents exact # of each atom in one FU
NO2
 Represents only the ratio of each atom
NOTE: Most ionic compound formulas are already EFs. Not so for most molecular compounds!

6. From Green Handout Sheet (PS4 & 5 Practice)
1(b) What is the empirical formula for P4O10?
Answer: P2O5
(In 4: 10, if you divide both #’s by 2, you get 2 : 5)
Ppt09(PS4)

7. Why do we have empirical formulas at all? Don’t we want MFs?
“derived from or guided by experience or experiment.” (dictionary.com)
In the lab, you cannot “see” the molecules of a molecular compound! How can you “know” that the molecules are N2O4?
Empirical formula is what you can get with ONLY “masses of elements” data (see next slide!).
NOTE: To get the molecular formula, you need other info
For us, that info is the molar mass (example later)
Ppt09(PS4)

8. Determining Empirical Formula from mass data of a sample (i. e
Determining Empirical Formula from mass data of a sample (i.e., knowing grams of elements A, B, C, etc.)
masses (of A, B, C, …)
moles (of A, B, C, …) atoms
ratio (of moles)
reduced ratio (whole-number ratio of moles)
Empirical formula (symbols with subscripts)
Ppt09(PS4)

9. From Green Handout Sheet (PS4 & 5 Practice)
2. An 8.26-g sample of a compound is chemically separated to form 3.31 g of sulfur and 4.95 g of oxygen.
What is the empirical formula?
What is the mass % of O in the compound?
(atoms!)
(atoms!)
 3:1 ratio  EF is SO3
Ppt09(PS4)

10. From Green Handout Sheet (PS4 & 5 Practice)
2. An 8.26-g sample of a compound is chemically separated to form 3.31 g of sulfur and 4.95 g of oxygen.
What is the mass % of O in the compound?
(two ways; see board or key)
Ans. 59.9% OR %
2 (follow up): If the molar mass of this compound is (separately) determined to be 160. g/mol, what is the molecular formula? [Note: this example is made-up]
Approach?
Ppt09(PS4)

11. To get MF from EF and MM, find “empirical mass” (EM) (see p
To get MF from EF and MM, find “empirical mass” (EM) (see p.116 in Tro)
Since the MF must be a “multiple” of the EF, the MM must be a multiple of the EM
Last example:
EF was SO3; thus EM was ~80. g/mol
MF must be (SO3)n
Thus, if MM is 160. g/mol, n must be 2.
Generally speaking, you can find n by dividing MM by EM! (Don’t memorize, just think about it!!)
Ppt09(PS4)

12. A Note About Ratios
A ratio is a fraction—take one thing and divide it by another.
If you do the division, you’ll get ONE NUMBER—this number is actually the ratio of the numerator to one of the denominator
See example on prior slide; moles of O : moles of S is basically 3 (to 1).
If the number you get is not extremely close to an integer value, you’ll need to reduce further (next slide)
Ppt09(PS4)

13. Reducing a non-whole-number ratio to a whole-number ratio (see pp 114-115 in Tro)
[NOTE: Never forget that in an empirical formula problem, you need the ratio of moles not grams!!!]
One approach:
Divide all mole values by the smallest. (That value will become a “1”)
If all other resultant values are not extremely close to a whole number, multiply all (new) values by a small whole number (2, 3, 4, etc.) until a whole number ratio is obtained
To take some of the “randomness” out of this approach, look for decimals that equate to simple fractions (0.25, 0.333, 0.5, 0.666, 0.75, etc.; SEE NEXT SLIDE )

14. 2nd Step “Multiplication” Help (Discussed in Tro on p. 115)
Since….
…multiply by ____ to make the number a whole #
0.333 =
0.666 =
0.500 =
0.250 =
0.750 =
0.200 =
0.400 =
0.800 =
3 (1/3 x 3 = 1)
3 (2/3 x 3 = 2)
2 (1/2 x 2 = 1)
4 (1/4 x 4 = 1)
4 (3/4 x 4 = 3)
5 (1/5 x 5 = 1)
5 (2/5 x 5 = 2)
5 (4/5 x 5 = 4)
Ppt09(PS4)

15. 2nd Step “Multiplication” Examples
If after dividing through by the lowest # moles you get:
…multiply (all #’s) by ____
to obtain (after rounding)
C1H2.66
Al1Cr1.51O6
C1.334H2.999O1
X2.01Y1Z2.75 3
C3H8 2
Al2Cr3O12 3
C4H9O3 4
X8Y4Z11
Ppt09(PS4)

16. From Green Handout Sheet (PS4 & 5 Practice)
3. You have determined that in a certain compound, the ratio of moles of C : H : O is: moles : 5.78 moles : 2.17 moles.
What is the empirical formula of the compound?
Multiply all numbers in ratio by 3:
NOTE: If you had rounded first: C3H6O2!
Ppt09(PS4)

17. PS Sign-Posting
The concepts and skills related to problem 9 (and also 12, really) on Written PS4, and numerous problems in Mastering PS4 Part 2, have been covered in this PowerPoint. Give those problems a try now!
Ppt09(PS4)