Molecular Formula - Online Tutor, Practice Problems & Exam Prep

Understanding the relationship between empirical and molecular formulas is crucial in chemistry. The empirical formula represents the simplest whole-number ratio of elements in a compound, while the molecular formula indicates the actual numbers of atoms present. To convert an empirical formula to a molecular formula, one can use the "n factor," which is multiplied by the subscripts in the empirical formula. For instance, glucose with an empirical formula of CH₂O and an n factor of 6 yields the molecular formula:

$C_6{H}_{12}{O}_6$

If the n factor is unknown, additional steps are required to calculate it, which involves determining the compound's empirical formula and molar mass. This process allows chemists to accurately identify the composition of substances and understand their properties.

Molecular Formula

Recall that the molecular formula gives the actual number of different elements in a given compound, meaning it gives us the real formula for the compound. The empirical formula, on the other hand, just gives us the relative number of those different elements in that same compound. Up to this point, we've learned to calculate the empirical formula when given either the masses or the percentages of those elements within that compound.

But how exactly would we determine the molecular formula for those same compounds? Well, here we have a chart, and in this chart we're talking about 3 compounds, glucose, octane, and salicylic acid. We're going to say that their empirical formulas are given as C1H2O1. Remember that there are subscripts of one and one. There we have C4H9 and C7H6O3. In order to determine their molecular formulas, we can use N factors.

N factors are just numbers that we can multiply the empirical formula by in order to determine the molecular formula. So for example, glucose has an N factor of 6. That means I will multiply each of the subscripts of the empirical formula by 6 and that'll give me my molecular formula. So that would be 6 carbons, 12 hydrogens, and six oxygens. So the molecular formula of glucose is C6H12O6.

Octane has an N factor of 2, so multiplying those subscripts gives me C8H18. And then finally, salicylic acid has a factor of 1, which means that the molecular formula and empirical formula are the same. So just remember, when you're given the N factor, just multiply the empirical formula by it and you'll determine what your molecular formula will be.

But what happens when we don't have the N factor? How do we determine it? In the next series of videos, we'll go through the different steps required to find the different N factors for any given compound once you've determined its empirical formula.

Molecular Formula

Everyone, in this video, we're going to calculate the molecular formula. Now, once the empirical formula is determined, the molecular formula can be obtained if the molar mass is also known. Here in this example question, it says Ibuprofen, which has a molar mass of 206.3 grams per mole, works by reducing the production of prostaglandins. This is the chemical component responsible for pain. Now, if its percent composition is 75.70% Carbon, 8.80% Hydrogen, and 15.50% Oxygen, determine its molecular formula. Alright, step 1 here is just us repeating the steps necessary to determine the empirical formula of the compound.

Now, remember the step is we're going to use our element symbols. We're dealing with Carbon (C), Hydrogen (H), Oxygen (O). We're going to assume we're dealing with 100 grams of our compound, so each one of these percentages can be changed into grams. So we're going to have here 75.70 grams, we're going to have here 8.80 grams, and then 15.50 grams. Now that I have each of their grams, I'm going to change them into moles. And remember, we do this by dividing them by their atomic masses, which we find in the periodic table, and then 1 mole of oxygen is 16 grams. Here, the grams would cancel out. And remember, to avoid rounding errors, we make sure that we have at least 4 decimal places for each one of the mole totals. So when we plug this in, we're going to get 6.3031 moles of carbon, 8.7302 moles of hydrogen, and 0.9688 moles of oxygen.

Next, we divide all our mole totals by the smallest mole answer. The smallest mole out of these is 0.9688. So all of them are going to get divided by that number. This is going to give us 1 oxygen, 9 hydrogen, and 6.5 carbon. Now remember, we need to have a 0.1 or 0.9 in order to round. Here it's 6.5 so we can't round. That means we need to multiply this by a number to make it a whole number. We have to multiply this by 2. And if I multiply my Carbon by 2, I'd have to multiply the Hydrogen by 2 and the Oxygen by 2. So that's going to give me 13 carbons, 18 hydrogens, 2 oxygens. That together would give me my empirical formula, which is C₁₃H₁₈O₂. Alright, so we've determined what our empirical formula is, and just to match things up, it's C₁₃H₁₈O₂.

Alright, we've done all of step 1 so far. Step 2. We now need to calculate the empirical mass of the compound. Now, your empirical mass comes from the empirical formula. So we have 13 carbons, actually 18 hydrogens and 2 oxygens. Multiply them each by their atomic masses from the periodic table. When we do that, we're going to get 156.13 grams carbon, 18.144 grams hydrogen, and 32 grams oxygen. Add them all up together, that gives us 206.274 grams per mole. So this is the empirical mass of our empirical formula.

Now, step 3. We divide the molar mass, which was given to us in the very beginning of the question of the molecular formula, by the empirical mass to determine our n factor. So here we're going to say n equals our molar mass, which again was given to us in the very beginning divided by our empirical mass, which I'll just abbreviate as EM. So if we look at the very beginning of the question, they tell us that our molar mass of Ibuprofen is 206.3 grams per mole. We just calculated the empirical mass as 206.274 grams per mole. So here, n equals just 1. Alright.

So now what do I do with this n factor? Well, we come down to step 4 here. We're going to say we multiply the subscripts of the empirical formula by the n factor to get the molecular formula. So basically what I'm saying here is that we're going to say that n times your empirical formula equals your molecular formula. Okay. So we're going to say n is 1 times my empirical formula, which again was C₁₃H₁₈O₂. So that gives me C₁₃H₁₈O₂ as my molecular formula. This would be our final answer. We see here that sometimes your empirical formula and your molecular formula are the same. This happens anytime n is equal to 1. But there will be instances where you find your n factor and it'll be another number like 2 or 3, you'll multiply your empirical formula by that new n factor, and you'll see that in those cases your empirical formula and your molecular formula don't match. It just happens to be in this particular question, our empirical formula and molecular formula are the same.

Use the given empirical formula and molar mass to determine the molecular formula.

Empirical Formula: NH₂                      Molar Mass: 32.052 g/mol

• A. NH₄
• B. N₄H
• C. NH₂
• D. N₂H₄

Glyceraldehyde (M = 90.078 g/mol), a simple monosaccharide, is comprised of 39.999% C, 6.714% H, and 53.297% oxygen by atomic weight. What would be its molecular formula?

• A. C₃H₆O₃
• B. CH₂O
• C. C₃HO₃
• D. C₂H₈O₂

Elemental analysis of a pure compound indicated that the compound had 72.2% C, 8.50% H and the remainder as O. If 0.250 moles of the compound weighs 41.55 g, what is the molecular formula of the compound?

• A. C₅H₇O
• B. C₁₀H₁₄O₂
• C. C₆H₈O₂
• D. C₁₂H₁₆O₄

Here’s what students ask on this topic:

How can I find the molecular formula?

To find the molecular formula, you first need to determine the empirical formula, which represents the simplest whole-number ratio of elements in the compound. Then, you'll need the molecular weight (or molar mass) of the compound, which can be found using a mass spectrometer or calculated from the atomic weights of its constituent elements.

Here's a step-by-step guide:

1. Obtain the Percent Composition: If you're given percentages of each element in the compound, assume a 100-gram sample so that the percentages can directly translate into grams.
2. Convert to Moles: Use the atomic masses of the elements (from the periodic table) to convert the mass of each element into moles.
3. Find the Simplest Ratio: Divide the moles of each element by the smallest number of moles to find the simplest whole-number ratio.
4. Determine the Empirical Formula: Use the ratios to write the empirical formula.
5. Calculate the Empirical Formula Mass: Add up the atomic masses of all the atoms in the empirical formula.
6. Find the Ratio of Molecular Mass to Empirical Formula Mass: Divide the molecular weight of the compound by the empirical formula mass.
7. Determine the Molecular Formula: Multiply the subscripts in the empirical formula by the ratio calculated in the previous step to get the molecular formula.

How can the molecular formula be found from the empirical formula?

To find the molecular formula from the empirical formula, you need to determine the molecular mass of the compound and compare it with the empirical formula mass. Here's how you do it:

1. Calculate the empirical formula mass by adding up the atomic masses of all the atoms in the empirical formula.
2. Obtain the molecular mass of the compound. This is often done using experimental data from techniques like mass spectrometry.
3. Divide the molecular mass by the empirical formula mass to find the ratio between them. This ratio should be a whole number or very close to a whole number.
4. Multiply all the subscripts in the empirical formula by this ratio to get the molecular formula.

For example, if the empirical formula is CH₂ and the molecular mass is found to be approximately 56 g/mol, the empirical formula mass (12 for C + 2*(1 for H)) is 14 g/mol. Dividing the molecular mass by the empirical formula mass gives 4. Therefore, the molecular formula is C₄H₈, which is four times the empirical formula.

What is the molecular formula?

A molecular formula represents the actual number of atoms of each element in a molecule of a compound. It is a way of expressing the chemical composition of a molecular substance using the symbols for the elements and numerical subscripts to indicate the number of atoms of each element present in the molecule. For instance, the molecular formula for water is:

$H_{2}O$

which tells us that each molecule of water is composed of two hydrogen atoms and one oxygen atom. In contrast to empirical formulas, which show the simplest whole-number ratio of atoms in a compound, the molecular formula provides specific information about the molecular structure. This can be particularly important for distinguishing between isomers—compounds that have the same empirical formula but different molecular structures and properties.

How do you calculate the molecular formula?

To calculate the molecular formula, you first need to determine the empirical formula, which represents the simplest whole-number ratio of elements in a compound. Here's a step-by-step process:

1. Obtain the percent composition of each element in the compound, either from experimental data or given information.
2. Convert these percentages to grams (assuming a 100-gram sample) to make calculations easier.
3. Convert the grams of each element to moles by dividing by the atomic mass of each element (from the periodic table).
4. Find the simplest whole-number ratio of moles of each element by dividing all the mole amounts by the smallest number of moles calculated.
5. This ratio gives you the empirical formula.

To find the molecular formula, you need the molecular weight (or molar mass) of the compound, which can be determined experimentally or provided.

6. Divide the molecular weight by the empirical formula weight (the sum of all the atomic masses in the empirical formula).
7. This gives you a whole number, which you multiply by the subscripts in the empirical formula to get the molecular formula.

For example, if the empirical formula is CH₂ and the molecular weight is 56 g/mol, and the empirical formula weight is 14 g/mol, then the molecular formula is C₄H₈ (56/14 = 4, so multiply each subscript by 4).