The molecular formula is important information for any chemical compound. The molecular formula tells what atoms make up a compound and the number of atoms. You must know the empirical formula to calculate the molecular formula, and you must know that the molecular formula is an integer multiple of the empirical formula.
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Part 1 of 3: Deriving Molecular Formulas from Empirical Formulas
Step 1. Know the relationship between the molecular and empirical formulas
Empirical formulas show the ratio of atoms in a molecule, for example two oxygens for each carbon. The molecular formula tells the number of each of the atoms that make up the molecule. For example, one carbon and two oxygen (carbon dioxide). These two formulas have a comparative relationship (in whole numbers) so that the empirical formula will become the molecular formula when multiplied by the ratio.
Step 2. Calculate the number of moles of gas
This means using the ideal gas law. You can find the number of moles based on the pressure, volume, and temperature obtained from the experimental data. The number of moles can be calculated using the following formula: n = PV/RT.
- In this formula, is the number of moles, P is pressure, V is the volume, T is the temperature in Kelvin, and R is the gas constant.
- Example: n = PV/RT = (0.984 atm * 1 L) / (0.08206 L atm mol-1 K-1 * 318, 15 K) = 0.0377 mol
Step 3. Calculate the molecular weight of the gas
This step can only be done after finding the moles of the constituent gases using the ideal gas law. You should also know the mass mass of the gas in grams. Then, divide the mass of the gas (grams) by the moles of gas to get the molecular weight.
Example: 14.42 g / 0.0377 mol = 382.49 g/mol
Step 4. Add up the atomic weights of all the atoms in the empirical formula
Each atom in the empirical formula has its own atomic weight. This value can be found at the bottom of the atomic grid on the periodic table. Add up the atomic weights to get the empirical formula weight.
Example: (12, 0107 g * 12) + (15, 9994 g * 1) + (1, 00794 g * 30) = 144, 1284 + 15, 9994 + 30, 2382 = 190, 366 g
Step 5. Find the ratio between the molecular and empirical formula weights
To do this, you can find the result of dividing the actual molecular weight by the empirical weight. Knowing the result of this division allows you to find out the result of the division between the molecular formula and the empirical formula. This number must be a whole number. If the comparison is not a whole number, you must round it.
Example: 382, 49 / 190, 366 = 2,009
Step 6. Multiply the empirical formula by the ratio
Multiply the small number in the empirical formula by this ratio. This multiplication yields the molecular formula. Note that for any compound with a "1" ratio, the empirical formula and the molecular formula will be the same.
Example: C12OH30 * 2 = C24O2H60
Part 2 of 3: Finding Empirical Formulas
Step 1. Find the mass of each constituent atom
Sometimes, the mass of the constituent atoms is known or the data will be given as a mass percentage. In this case, use the sample compound sample of 100 g. This allows you to write the mass percentage as the actual mass in grams.
Example: 75, 46 g C, 8, 43 g O, 16, 11 g H
Step 2. Convert mass to moles
You must convert the molecular mass of each element to moles. To do this, you must divide the molecular mass by the atomic mass of each element. You can find the atomic mass at the bottom of the element grid on the periodic table.
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Example:
- 75.46 g C * (1 mol / 12.0107 g) = 6.28 mol C
- 8.43 g O * (1 mol / 15.9994 g) = 0.53 mol O
- 16.11 g H * (1 mol / 1.00794) = 15.98 mol H
Step 3. Divide all mole values by the smallest mole value
You must divide the number of moles for each separate element by the smallest number of moles of all the elements that make up the compound. To do this, you can find the smallest mole ratio. You can use the smallest mole ratio because this calculation gives the non-abundant element a value of “1” and results in the ratio of the other elements in the compound.
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Example: The smallest number of moles is oxygen with 0.53 moles.
- 6.28 mol/0.53 mol = 11.83
- 0.53 mol/0.53 mol = 1
- 15, 98 mol/0.53 mol= 30, 15
Step 4. Round your mole value to a whole number
These numbers will be small numbers in the empirical formula. You should round it to the nearest whole number. After looking up these numbers, you can write down the empirical formula.
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Example: The empirical formula is C12OH30.
- 11, 83 = 12
- 1 = 1
- 30, 15 = 30
Part 3 of 3: Understanding Chemical Formulas
Step 1. Understand the empirical formula
Empirical formulas provide information about the ratio of one atom to another in a molecule. This formula does not give precise information about the number of atoms that make up the molecule. Empirical formulas also do not provide information about the structure and bonds of atoms in molecules.
Step 2. Know the information given by the molecular formula
Like empirical formulas, molecular formulas do not provide information about bonds and molecular structure. However, unlike empirical formulas, molecular formulas provide details about the number of atoms that make up a molecule. The empirical formula and the molecular formula have a comparative relationship (in whole numbers).
Step 3. Understand the structural representation
Structural representations provide more in-depth information than molecular formulas. In addition to showing the number of atoms that make up a molecule, structural representations provide information about the bonds and structure of the molecule. This information is very important for understanding how the molecule will react.
