Avogadro's Number and the Mole

Avogadro's Number and the Mole Definition

Let's start by looking at the definition of Avogadro's number and the mole.

Avogadro's number is unitless. We can technically use it to count anything, but it is such a massive number that it is really only practical for small things like atoms or molecules. It is similar to how "a dozen" can refer to 12 things, but we typically use it just to count eggs. When we count atoms/molecules, we use the unit mole (not the animal, though they are cute!).

Avogadro's Number and the Mole Importance

The mole is the standard way to count things in chemistry. It is used in our chemical formulas and equations, and it is used to convert from AMU (atomic mass unit) to the much more practical unit of grams.

The Mole, Avogadro's Number, and Stoichiometry

When we write chemical formulas and equations, the units we are using are mols. Let's look at a reaction.

$$2H_{2\,(g)} + O_{2\,(g)} \rightarrow 2H_2O$$

The coefficient (number in front) for each molecule is equal to the number of mols. If there is no coefficient written, it is 1. So in this case, we have 2 mols of H₂ and H₂O and 1 mol of O₂. The subscript (little number) tells us the number of moles per element. So for O₂, there are 2 mols of atomic oxygen, O, per 1 mol of molecular oxygen, O₂. Now that we know the basics, let's work on a problem.

Chemical reactions must follow the law of conservation of mass. This law says that mass cannot be created or destroyed in a reaction (i.e. our products need to have the same mass as our reactants.) When writing chemical reactions, we have to abide by this law. To achieve this, we balance the number of moles for each element using stoichiometry (stoichiometry is just a fancy word for the ratio between products and reactants). Here's an example:

When we do experiments, we are typically measuring in grams, not mols. So now we need to look at how we convert between these units.

The Atomic and Molar Mass

The way we convert from mols to grams (or vice-versa) is by using an element's atomic mass.

When we are referring to the mass of a whole compound, it is called the molecular weight/molar mass. The molecular weight is the sum of the atomic masses of each element in a compound. Being able to convert from moles to grams is important when experimenting for many reasons.

One reason is that we can calculate the yield of a reaction (mass of products), based on the mass of reactants. It is also important since we measure things in grams, but all of our chemical equations/formulas are in mols.

Using Avogadro's Number and the Mole Concept for Conversions

It's helpful to consider these conversions to be a ladder or staircase. You can either climb up or down, but you need to take each step to reach your destination.

By following the "steps" we can convert between units. StudySmarter Original.

For some elements/compounds, the number of atoms is equal to the number of molecules. For example, in NaCl, both Na and Cl have no subscripts (so a subscript of 1), so these values are equal.

Percent Composition

Mols and molar mass can also be used to calculate percent composition.

Here's an example:

We can also use percent composition to calculate the mass of an element. So for our example above, H₂O has an 11.2% composition of H:

$$\frac{11.2}{100}*18.02\frac{g}{mol}=2.02\frac{g}{mol}$$

When calculating percent composition, you should make sure your percentages add to 100. For the previous example, If we forgot to multiply the mass of H by 2, the total percentage would not have equaled 100.