By using the definitions for the mole, molecular weight, and the information given in a chemical equation, it is possible to make a great number of calculation involving chemical reactions.

Calculate the number of grams of hydrogen in 30.0 g of water.

Calculate the number of moles of hydrogen atoms in 30.0 g of water.

Calculate the number of atoms of hydrogen in 30.0 g of water.

To make mole-mole calculations we first need a balanced chemical equation so that we can find the mole to mole ratios of the reactants and products.

Consider the following problem:

How many moles of CO₂ are produced when 4.5 moles of CH₄ are burned in oxygen?

How many moles of H₂O are produced when 4.5 moles of CH₄ are burned in oxygen?

When making mole to mass or mass to mole calculations, we first need a balanced chemical equation. We must change from grams to moles of a substance or from moles to grams.

Consider the following problems:

Calculate the mass of CO₂ produced when 4.5 moles of CH₄ is burned in oxygen.

Calculate the moles of CO₂ produced when 25.0 grams of CH₄ is burned in oxygen.

To make mass-mass calculations we must first balance the chemical equation. Next we convert from grams of the substance given (using molecular weight) to moles and use the mole-mole ratio to convert to the moles of the desired substance. Last we must convert to grams of the desired substance.

Consider the following examples:

Calculate the number of grams of CO₂ formed when 10.0 g of CH₄ is burned.

What weight of sodium sulfate is produced when 5.00 g of sodium hydroxide is reacted with excess sulfuric acid?

Calculate the number of molecules of CO₂ formed when 10.0g of CH₄ is burned.

We can sum up what has been covered so far in the following flow chart.

grams of Using Moles of Using Moles of Using Grams of

what you ----------->what you ---------->what you -----------> what you

have molecular wt have balanced eq. want molecular wt want

You will notice that there are two entry points and two exit points depending upon what it is that is asked.

When finding the amount of a product formed, thus far we have assumed an excess of one of the reactants. In the above example, 10.0 g of CH₄ is burned completely using an excess of oxygen. However, when we are given the amounts of two reactants, it is likely that one of the reactants will be used up first (the limiting reactant) and that some of the other reactant(s) will be in excess (excess reactant). Consider the following example:

Example:

10.0 g of CH₄ is reacted with 5.00 g of Oxygen. How many grams of CO₂ is formed in the reaction.

Since we have two givens, dimensional analysis demands that we start the conversion twice, once for each given reactant. Start by balancing the chemical equation.

1moleCH₄ 1mole CO₂ 44.0 g CO₂

10.0 g CH₄ x ------------ x ------------------ x --------------- =

16.1g CH₄ 1mole CH₄ 1mole CO₂

1mole O₂ 1mole CO₂ 44.0 g CO₂

5.0 g O₂ x------------- x ---------------- x --------------- =

16.0g O₂ 1mole O2 1mole CO₂

There are two answers but only one of them is correct. Which one is correct?

Note that the calculation that gives the fewest number of moles must have started with the limiting reactant. The limiting reactant determines the amount of reactants consumed as well as the amount of products formed.

As a continuation of the above problem find how much of the excess reactant will remain after the reaction is complete.

Problem:

Iron (III) oxide reacts with carbon to form iron metal and CO₂. How many grams of iron can be formed from 1000 g of iron(III) oxide and 1000 g of carbon?

Here is a take home problem: .200 g of hydrogen react with 1.50 g of oxygen.

a) What is the limiting reagent? Show all calculations and logic steps.

b) What is the excess reagent? Show all calculations and logic steps.

c) Determine the number of grams of water formed.

d) Determine the number of grams of excess reagent remaining.