A 2.20 g-sample of a compound containing carbon, hydrogen, and oxygen is burned and it produces 4.61 g CO2 and 0.94 g H2O. What is the empirical formula of this compound?
Try converting everything into moles and comparing what you started with and what you ended up with. That'll help you figure out the ratio between the products and reactants. This in turn will help you determine how to write out the balanced chemical equation that you can look at to determine how much of each element comprises the compound.
Convert 4.61 g CO2 to g C.
4.61 x (atomic mass C/molar mass CO2) = ?
Convert 0.94 g H2O to g H.
0.94 x (2*atomic mass H/molar mass H2O) = ?
g O = 2.20- g H - g C = ?
Convert grams to mols
mols C = grams/atomic mass C
mols H = grams/atomic mass H
mols O = grams/atomic mass O
Now convert those mols C,H,& O to ratio with the smallest number being 1.00. The easiest way to do that is to divide the smallest number by itself (which makes it 1.000), then divide the other numbers by the same small number. Post your work if you get stuck.
C2H9O10
It's c2h2o
To determine the empirical formula of a compound, we need to find the ratios between the elements present in the compound. In this case, we have a sample that contains carbon (C), hydrogen (H), and oxygen (O).
First, we need to determine the number of moles of each element in the given compounds. To do this, we'll use the molar masses of carbon dioxide (CO2) and water (H2O).
The molar mass of carbon dioxide (CO2) can be calculated by summing the atomic masses of carbon (C) and two oxygen atoms (O):
12.01 g/mol (C) + (2 × 16.00 g/mol (O)) = 44.01 g/mol.
Using the given information that 4.61 g of CO2 are produced during combustion, we can calculate the number of moles of CO2:
Number of moles = mass / molar mass = 4.61 g / 44.01 g/mol ≈ 0.1049 mol.
Similarly, we can calculate the number of moles of water (H2O) using its molar mass:
Molar mass of water (H2O) = 2 × 1.01 g/mol (H) + 16.00 g/mol (O) = 18.02 g/mol.
Using the given information that 0.94 g of H2O are produced during combustion, we can calculate the number of moles of H2O:
Number of moles = mass / molar mass = 0.94 g / 18.02 g/mol ≈ 0.0522 mol.
Now, let's find the moles of carbon and hydrogen present in the compound.
From the balanced chemical equation of the combustion reaction, we know that every 1 mole of CO2 contains 1 mole of carbon. Therefore, the number of moles of carbon in the compound is also 0.1049 mol.
Similarly, from the balanced chemical equation, we know that every 1 mole of H2O contains 2 moles of hydrogen. Therefore, the number of moles of hydrogen in the compound is 2 × 0.0522 mol = 0.1044 mol.
Lastly, to find the moles of oxygen in the compound, we subtract the moles of carbon and hydrogen from the total number of moles:
Moles of oxygen = Total moles - Moles of carbon - Moles of hydrogen
Moles of oxygen = 0.1049 mol - 0.1049 mol - 0.1044 mol = -0.1044 mol
This negative value indicates that there is no oxygen present in the compound. However, this is likely due to experimental error or incomplete combustion. Since we cannot have a compound without oxygen, we need to recalculate this ratio using corrected values.
To do this, we subtract the mass of CO2 and H2O from the initial mass of the compound to find the mass of carbon:
Mass of carbon = Initial mass - Mass of CO2 = 2.20 g - 4.61 g = -2.41 g.
Since the result is negative, there must be an error in the data provided. Please recheck the values given for the mass of CO2 and H2O or verify the initial mass of the compound.