Calculating Empirical, Molecular Formula and Percent Composition

There are two kinds of formulas: Empirical and Molecular

Empirical:
- Gives lowest-term ratio of atoms or moles in formula
- All Ionic compounds are empirical

Molecular:
- Gives all atoms that make up molecule
- Either Ionic or Covalent

Eg. C6H12O6 = Molecular Formula
CH2O = Empirical Formula

Example:
A gas has an Empirical Formula of CH2. What is the Molecular Formula if the mass of one mole is 42.0 g?

Have to find a whole number (N) = Molar Mass/Empiricial Mass = 42.0/14.0 = 3 (N)
(CH2) X 3 = C3H6

Determining Empirical Formula Given Mass


Example:

Determine the Empirical Formula of Fe and O given 10.87 g of Fe and 4.66 g of O.

First, grams must be converted to moles.
Fe - 10.87 g X 1 mole/55.8 g = 0.1948 moles
O - 4.66 g X 1 mole/16.0 g = 0.291 moles

Now, you must divide all of the derived numbers by the smallest amount of moles.
Fe - 0.1948/0.1948 = 1
O - 0.291/0.194 = 1.49, which can be rounded to 1.5

Lastly, you must multiply until you reach a whole number
Fe - 1 X 2 = 2
O - 1.5 X 2 = 3

Therefore, the Empirical Formula is Fe2O3.

Percent Composition:
- The % by mass of elements in a compound
- MUST ADD UP TO ABOUT 100% PERCENT (99.9 or 100.1)

Example:
What is the Percent Composition of CO2?

First, calculate the molar mass.
C - 12.0 g X 1 = 12.0 g
O - 16.0 g X 2 = 32.0 g
Therefore, the Molar Mass is 44.0 g.

Now, calculate each element's % of that total, rounding to one decimal place.
C - 12.0 g/44.0 g X 100 = 27.3 %
O - 32.0 g/44.0 g X 100 = 72.7 %

Percent Composition to Empirical Formula

Example:
A substance is 45.8% Sulphur and 54.2% Fluorine. What is the Empirical Formula?

First, assume 100.0 grams of material.
Then, convert to moles.
S - 45.8 g X 1 mole/32.1 g = 1.43 moles
F - 54.2 g X 1 mole/19.0 g = 2.85 moles

Then, divide by the smallest number.
S - 1.43 g/1.43 g = 1
F - 2.85 g/1.43 g = 2

Therefore, the Empirical Formula is SF2.

Mole Conversions

Mole Conversions

Mole Conversions can be from either:

1.      1.  Atoms, molecules, formula units or particles to moles and vice versa

For a quick demonstration:

A)  a)   Convert 5.1 x 10^21 atoms to moles

.

5.1 x10^21 atoms x 1 mole                          = 8.5 x 10^-3

                               6.022 x 10^23 atoms

OR

2.    2.   Grams to moles and vice versa

For a quick demonstration:

A)   a)  Convert 16.3g of H2O to moles.

32.3g x 1 mole   = 1.79mol H2O

              18.0g

For a more advanced worksheet involving mole conversions among other conversions click here

à http://www.csun.edu/~JTE35633/worksheets/Chemistry/11-3MoleConversions.pdf

AND for a simple YouTube explanation of how and why moles are converted as such watch this:

Mass

Mass

"Basically the total weight of something...!"

 (Atomic)                                       Mass :  The mass of atoms as compared to the very special CARBON-12 ATOM in atomic mass units (u)

  (Formula)                                        Mass :  The total mass of every atom in a formula of only an ionic                                 compound in u.

(Molecular)                                  Mass :  The total mass of every atom that combine to make: a molecule of a covalent compound, a molecule of an organic compound or a polyatomic element in u.

(Molar)                                        Mass :  The mass of 1 mole (which has 6.022 x 10^23 particles) of something and is the same value (numerically of course) as atomic, formula or molecular mass but is in GRAMS PER MOLE

What is a mole anyway?

A mole, not to be confused with a small furry rodent, is a unit of measurement used for small particles with a unit conversion factor which allows for say almost weightless atoms be shown in a more manageable format.

To be precise: A mole is 6.022 x 10^23 formula units, particles, molecules or atoms of something.

Which raises the question ???

            Why the random number?

Well this number... called Avogadro'a number wasn't created by Amedeo Avogadro!

THIS IS HIM! OMG!         ----------------------->     

Anyway, apparently he's pretty special, for a french scientist named the number after him in 1909.

By the way...

This number is just like any ordinary way of saying there is so much of something.  For instance, a dozen is a simple way of saying there is twelve of something just as a mole is a simple way of saying there is 6.022 x 10^23 of something!

See this website for super in depth information and discussion on this iconic number!

http://www.chem1.com/acad/webtext/intro/MOL.html

Lab 2E Finding the Thickness of Aluminum

Finding the thickness of aluminum foil combines two different concepts to measure the thickness:
Volume of a solid: V = L x W x H (length x width x height)
Density of a substance: D = m / V (mass / Volume)

Height being the thickness of the aluminum sheet being measured. Therefore in the experiment you determine the thickness in the volume equation.
Aluminum foil 17cm x 15cm (1.04g)
Start with density: (density of aluminum = 2.70g/cm3)
2.70 = 1.04g/V
    V = 0.39

Then using the volume in the Volume equation:
0.39 = 17 x 15 H
0.39 = 255 x H
H = 0.39 / 255
H = 1.53x10(-3)cm

Then find the average of 3 pieces of aluminum foil:
1.53x10(-3)cm +
1.48x10(-3)cm +
1.52x10(-3)cm +
/3 = 1.51x10(-3)cm

Then find the % off the acceptable thickness of aluminum foil (1.55x10(-3)cm)

finding the average by (1.51x10(-3)cm) - (1.55x10(-3)cm) / 1.55x10(-3)cm = 2.5%

Graphing

Today, we did some graphing using Excel. It is a very useful tool. You can alter the graph, making it as large as you want, labelling the axis, changing the colours and styles, as well as many other things. You can also choose to put a line-of-best-fit as well as displaying the equation. Using these tools, you can figure out if the numbers are linear or not.
An example of a linear relationship is:                                              


An example of a non-linear relationship is:

Plotting data makes it easier to read. For substances, for example, you can determine their melting point, boiling point, etc. It is also easy to determine which data is reliable, and which aren't on the line-of-best-fit and may have been a mistake or inaccurately measured.

Density

What is density?
Density is a physical property of matter. The formula to find density is:

Density = mass / volume

Ways you can tell how dense an object is through weight and sight suggested in these pictures:

-If the density of an object is greater than that of water the object will sink.

-If the density of an object is lower than that of water the object will float.

an example of how to find the density of an iron bar:
1200 g / 1.25 L = 960 g / L

Measurement and Uncertainty

In science, we take measurements. These, however, are never exact. Measurements are only our best estimate, which is subject to some uncertainty.
The only numbers that are exact are when you are counting whole things. Ex. seven apples, 15 pencils, 48 dogs, etc.

There are two types of Uncertainty: Absolute Uncertainty and Relative Uncertainty.

Absolute Uncertainty

- This uncertainty is expressed in units of measurement, not as a ratio
- Two methods can be used

Method One:
- Take at least three measurements
- Calculate the average of these numbers
- The Absolute Uncertainty of these measurements is the difference between the average and the highest or lowest reasonable measurement
- Always remove any measurements that are unreasonable

Ex.
Trial #    Mass of an Object
   1                 27.3 g
   2                 27.5 g
   3                 27.4 g
   4                 27.9 g
   5                 27.4 g

You would remove 27.9 g, since it is unreasonable. You would then add up the other measurements and divide by four, calculating the average:
109.6 g / 4 = 27.4 g
You would then find the difference between 27.4 g and the highest or lowest measurement:
27.4 g - 27.3 g = 0.1 g
27.4 g - 27.5 g - 0.1 g
Therefore, the Absolute Uncertainty of these measurements is 0.1 g. You would record the mass as:
27.4 +- 0.1 g

Method Two:
-Determine the uncertainty of the instrument being used 
-Measure as precisely as possible
-Estimate to 10% of the smallest division made on the instrument
Ex.
On a ruler, the smallest division is 0.1 of a cm, or 1 mm. Therefore, you should measure to the nearest tenth of a mm. If you estimate a measurement to be 25.34 cm on a ruler, you would record the number as 25.34 +- 0.01 cm. If you estimate a measurement to be 6.7oC on a thermostat, you would record the number as 6.7 +- 0.1 oC.

Relative Uncertainty
-This can be expressed as in percent or by using significant figures.
-Percent is most commonly used
-Relative Uncertainty = Absolute Uncertainty / Estimated Measurement
-To get a percent, you would multiply that number by 100.
Ex. Estimated measurement = 27.23 and the Absolute Uncertainty = 0.01
(0.01 / 27.23) X 100 = 0.03672
Therefore, the Relative Uncertainty is 0.04 %