Complete combustion of 3.30 g of a hydrocarbon produced 10.2 g of CO2 and 4.68 g of H2O. What is the empirical formula for the hydrocarbon?
I did the calculation and the ratio comes out to
C: 1 H: 2.24
So I thought the answer is CH2 but apparently its wrong. Anyone help..
With 1:2.24 you can't round to 1:2.
Multiply by a whole number until you get two whole numbers (or at least to close whole numbers).
1:2.24
2:4.48
3:6.72
4:8.96 which I would round to 4:9
for the formula C4H9. Having done that I don't know what to do except look at your data for as far as know there is no C4H9. But I would key in C4H9; the person in charge of making up these problems may not have worked it through.
Well, it seems your calculation took a detour and got a little lost in the chemistry maze. Let me help you find your way back with a dash of humor!
To determine the empirical formula, we need to find the simplest ratio of atoms in the compound. Let's take a closer look at your numbers.
We know that 3.30 g of the hydrocarbon produced 10.2 g of CO2 and 4.68 g of H2O. But wait, isn't CO2 composed of carbon and oxygen, not carbon and hydrogen? Oxygen must have been sneaking into the party without an invitation!
So, let's do some calculations to find out how much carbon and hydrogen we have in our hydrocarbon.
First, let's calculate the amount of carbon in CO2. It's given as 10.2 g, but since the atomic mass of carbon is 12 g/mol and CO2 has a molar mass of 44 g/mol, we can figure out that 10.2 g of CO2 contains 2.77 moles of carbon (10.2 g / 44 g/mol * 1 mole of carbon / 1 mole of CO2).
Now let's calculate the amount of hydrogen in H2O. It's given as 4.68 g, and since the atomic mass of hydrogen is 1 g/mol and H2O has a molar mass of 18 g/mol, we can find that 4.68 g of H2O contains 2.6 moles of hydrogen (4.68 g / 18 g/mol * 2 moles of hydrogen / 1 mole of H2O).
Now, let's compare the moles of carbon and hydrogen. We have 2.77 moles of carbon and 2.6 moles of hydrogen. We need to find the simplest whole-number ratio between these two, and it turns out to be approximately 1:1.
Therefore, the empirical formula of the hydrocarbon is CH, not CH2. It seems that the hydrocarbon had a surprising guest, and that was oxygen, not hydrogen!
Remember, chemistry can sometimes be as tricky as a clown juggling flaming torches. But don't worry, with a bit of laughter, we can solve any problem!
To find the empirical formula of the hydrocarbon, you need to determine the simplest whole number ratio of carbon and hydrogen atoms present in the compound.
First, calculate the moles of carbon dioxide (CO2) produced:
moles of CO2 = mass of CO2 / molar mass of CO2
moles of CO2 = 10.2 g / 44.01 g/mol ≈ 0.232 mol
Then, calculate the moles of water (H2O) produced:
moles of H2O = mass of H2O / molar mass of H2O
moles of H2O = 4.68 g / 18.015 g/mol ≈ 0.259 mol
Now, let's find the moles of carbon and hydrogen in the hydrocarbon. Since the whole compound is combusted to produce CO2 and H2O, the moles of carbon and hydrogen in the hydrocarbon should be equal to the moles of carbon and hydrogen in the products.
Moles of carbon in the hydrocarbon = 0.232 mol
Moles of hydrogen in the hydrocarbon = 0.259 mol * 2 = 0.518 mol
Now, we need to find the simplest whole number ratio between the moles of carbon and hydrogen. Divide both values by the smaller value (0.232 mol in this case) to obtain the ratio:
Moles of carbon in the hydrocarbon = 0.232 mol / 0.232 mol ≈ 1
Moles of hydrogen in the hydrocarbon = 0.518 mol / 0.232 mol ≈ 2.23
Now, round the ratio of carbon to the nearest whole number and hydrogen to the nearest whole number:
Ratio of carbon to hydrogen ≈ 1:2
Hence, the empirical formula for the hydrocarbon is CH2. Your initial calculation was correct, and CH2 is indeed the empirical formula for the hydrocarbon.
To determine the empirical formula of the hydrocarbon, you need to find the simplest whole number ratio between the elements carbon (C) and hydrogen (H) in the given combustion reaction.
Let's start by finding the number of moles for each compound:
1. Convert the mass of CO2 produced to moles using its molar mass. The molar mass of carbon dioxide (CO2) is 44.01 g/mol.
Moles of CO2 = mass of CO2 / molar mass of CO2
= 10.2 g / 44.01 g/mol
2. Convert the mass of H2O produced to moles using its molar mass. The molar mass of water (H2O) is 18.02 g/mol.
Moles of H2O = mass of H2O / molar mass of H2O
= 4.68 g / 18.02 g/mol
3. Next, find the number of moles of carbon (C) and hydrogen (H) in the hydrocarbon.
Moles of C = Moles of CO2 (from step 1)
Moles of H = 2 * Moles of H2O (since each water molecule has 2 hydrogens) (from step 2)
Now, we will divide the moles of each element by the smallest number of moles obtained. This will give us the mole ratio, which we can then use to determine the empirical formula.
4. Divide the moles of C and H by the smallest number of moles obtained.
Moles of C / smallest mole value
Moles of H / smallest mole value
In this case, we have:
Moles of C = 10.2 g / 44.01 g/mol
= 0.2318 mol
Moles of H = 2 * (4.68 g / 18.02 g/mol)
= 0.5172 mol
Since the smallest mole value is 0.2318 mol, divide both moles by this value:
Moles of C / smallest mole value = 0.2318 mol / 0.2318 mol
= 1
Moles of H / smallest mole value = 0.5172 mol / 0.2318 mol
= 2.23
Now we need to round off the ratio to the nearest whole number. Multiplying by a factor to obtain whole numbers, we get:
C: 1 H: 2.23
Since we need whole numbers for the empirical formula, it would be C1H2. However, in empirical formulas, subscripts are typically written as whole numbers. Therefore, multiplying the ratio by 2, we get:
C: 2 H: 4
So, the empirical formula for the hydrocarbon is CH4, which is methane.
Please note that the rounding errors and approximations could lead to a slight deviation from the exact empirical formula.