The Mole Now that we know how to write and name chemical compounds, we need to understand how chemists use these formulas quantitatively. As chemists, we need to know how many atoms or molecules are reacting so that we can chemical formulas can be established. Up until this point, our studies have focused on single atoms or single molecules. The problem with this is that we cannot work with single entities in the laboratory. Thus, we need a way to connect the amount of particles present with quantities we can actually work with. In other words, we need to define a unit of matter that contains a known number of particles this unit is called the mole. A mole is defined as the amount of a substance that contains as many basic entities (atoms, molecules, or particles) as there are atoms in exactly 12 g of the carbon-12 isotope. In 12 g of carbon-12, there are exactly 6.0221415 10 23 atoms. So, 1 mole = 6.02 10 23 particles Although it might sound confusing initially, saying you have 1 mole of something is just like saying you have a dozen of something. Example, if I say I have a dozen eggs, how many eggs do I have? A dozen chickens? A dozen roses? A dozen homework assignments??? Well, if you had 1 mole of chickens, you d have 6.02 10 23 chickens. Or a mole of roses - 6.02 10 23 roses. And a mole of homework assignments?? This number, 6.02 10 23, is known as Avogadro s number. So how does the mole relate to quantities we can use in the lab?
Remember way back when, we discussed atomic mass in terms of atomic mass units (amu, or u). We said that, the atomic mass (weight) found on the Periodic Table is in units of amu. Well conveniently, 1 amu = 1 g/mol. Therefore, the atomic mass found on the Periodic Table will also tell us the mass in grams of 1 mole of that element. We call this the molar mass of that element. Ex: Molar Mass (M) is in units of g/mol Molar mass of (C) = mass of 1 mol of C atoms = 12.011g/mol = mass of 6.02 10 23 C atoms Molar mass of (O) = mass of 1 mol of O atoms = 15.994 g/mol = mass of 6.02 10 23 O atoms So how exactly does this help us? We can now convert from moles to mass and mass to moles, moles to particles and particles to moles! grams moles particles Moles grams moles Conversion Factors: = grams Grams moles grams = moles Moles particles moles = particles
We can use these conversion factors to solve problems. Example: What mass, in grams, is represented by 0.35 mol of aluminum? How many aluminum atoms are in this sample? If you measure out 16.5 g of carbon, how many moles of carbon do you have? How many atoms? So far, so good for elements. What about compounds? We discussed, yesterday, that in one molecule, you have the number of atoms present indicated by subscripts: CH 4
This is good, except it is inconvenient. So instead, we use the mole. If we have 1 mole of this compound, we have 1 mole of carbon atoms and 1 mole of hydrogen atoms. Just as we can obtain the molar mass of an element, we must be able to calculate the molar mass of compounds. To do so, we take into account the molar masses of the elements and their quantity in a particular molecule. Example: What is the molar mass of CH4? The molar mass of a compound is the sum of the molar masses of the elements present in that compound multiplied by their respective quantities. Let s try another: Calculate the molar mass of MgCO3.
Great, now let s say you have 1 mole of MgCO3, how many molecules are present? how many oxygen atoms are present? Working with grams, moles, and particles is an important aspect of chemistry. It requires that you practice working with these quantities and learning how to use the conversion factors to solve problems.
Describing Compound Formulas Now that we understand the importance of the mole and molar mass, we can use this information to determine the formula of an unknown compound. Any sample of a pure compound always consists of the same elements combined in the proportion by mass. Therefore, one way we can describe the molecular composition is by mass percent. Mass percent is the mass of each element in the compound relative to the total mass of the compound. This sounds complicated, but it really is just like any percentage you have seen before: like parts of a whole. Example: Let s say you want to find the mass percent of N and H in the compound NH3 This means that in a 100.0 g sample of NH3, there are 82.244 g of N and 17.755 g of H.
Let s try another example: Atoms, Molecules, and the Mole What is the mass percent of each element present in propane, C3H8? What mass of carbon is contained in 454 g of propane?