Stoichiometry
Stoichiometry is the calculation of the amounts of products consumed / reactants produced in a reaction.
The Concept of Moles
A mole is a certain amount of something. To be specific, a mole is of something. One mole is approximately the number of atoms in 12 grams of .
This is all one mole is - a quantity. To convert from, say, a mole of something to the number of molecules in that something, you just multiply by .
Molar Mass
The molar mass of a molecule is defined as the number of grams per mole of that substance.
The molar mass of one atom is simply equal to the atomic mass of that atom. This is because a mole has been chosen in such a way as to make this possible. (The atomic mass of is 12 amu, unified attomic mass units.)
We can then find the molar mass of a molecule quite easily. Take . The molar mass of water can be found by adding two times the molar mass of Hydrogen to one times the molar mass of Oxygen:
Moles to Mass
We can convert moles to mass and vice versa if we know the molar mass of a substance.
The formula is as follows:
where is number of moles, is mass in grams and is molar mass in .
Solving for Products / Reactants
If we know the number of moles of any product or reactant in a chemical equation, we can calculate the number of moles of any other product or reactant.
Take the following reaction:
Suppose that we know that 3 moles of hydrogen gas are produced. We can calculate the quantity of hydrogen chloride consumed using the following:
since the original reaction consumes 6 molecules of hydrogen chloride for every three molecules of hydrogen gas it uses.
Essentially in a chemical reaction:
Ideal Gases
An Ideal Gas is a gas that abides by the following rules (assumptions of ideal gases):
- Gases consists of molecules which are constantly in motion in straight lines.
- All collisions are perfectly elastic.
- Intermolecular forces are negligible.
- The particles' volumes are negligible.
- The molecules obey Newton's laws.
We can use the Ideal Gas Law to calculate specific properties of gases at temperatures, etc.
where is pressure in , is volume in , is moles, is the Universal Gas Constant, or about , and is the temperature in .
In addition, if the reaction is at STP ( and ) we can use a simplified equation to convert between moles and volume:
Example Problem
Q. Aluminium metal is added to hydrochloric acid at degrees Celsius and . of hydrogen gas is evolved. Calculate the mass of aluminium that was added.
First we must write the equation.
We can use the Ideal Gas Law to calculate the number of moles of hydrogen gas that were produced:
We can now calculate the number of moles of aluminium that must have been consumed:
(Remember to use unrounded values in your actual calculations, although you can write them down rounded.)
Finally, calculate the mass of Aluminium consumed:
A. .
Empirical Formula Calculation
To calculate the empirical formula of a substance given information, one can find the percentage by mass of each of the components.
Here's an example from a past examination.
Q.
A pure substance A is a colourless pungent liquid which is found by analysis to contain and .
A sample of A was burnt in an excess of dry oxygen and of carbon dioxide and of water were produced.
In a second experiment when of A was treated, all the chlorine was converted to silver chloride. of silver chloride was produced.
Calculate the empirical formula of A.
We know that A is a substance containing only carbon, hydrogen, chlorine and oxygen.
Write the two equations provided.
and are all variables. I've omitted some products in the equations above because they don't affect calculations.
Using the first equation, we see that we can calculate the amount of carbon and hydrogen in A, since there are no products other than A that contain carbon or hydrogen.
The Method
The method is to keep a table of the percentage by mass of each of the components, which we can then use to calculate the number of moles of each. Initially our table looks like this:
There's nothing there yet because we haven't done anything.
Step 0
Calculate some molar masses.
Step 1
Now we'll do some actual calculations.
We can now add to our table of percentage by mass.
Now we will start using the second equation:
We can calculate the mass of Oxygen trivially, by subtracting each of the others from .
Step 2
Add to your table the relative number of moles, i.e. number of moles as a percentage of A's total number of moles.
Element | ||||
---|---|---|---|---|
Percentage of Mass | ||||
Number of Moles (Use ) |
Finally calculate the ratio of the elements:
Element | ||||
---|---|---|---|---|
Percentage of Mass | ||||
Number of Moles (Use ) |
||||
Ratio |
The ratio is , which doesn't need to be simplified. (You may have to modify it so the ratio is all whole numbers.)
So our final empirical formula contains one Carbon, one Hydrogen, one Chlorine and one Oxygen atom.
A.