[Julie Smith]

hi, I've been working on this problem for a while now and can't seem to get anywhere with it, hopefully someone can help

so the background to the question is that to determine the % iron in a sample a titration can be formed with a standardized dichromate solution (Cr2O7)the overall charge is -2. So the iron sample is completely converted to the ferrous ion and then the ferrous ion is oxidized by reacting it with the dichromate solution until the end point is reached.

so the question is: a 2.7234g sample of an unknown compound is isolated and analyzed. If this unknown compound is found to contain 0.3420% Fe, and its assumed that 4 ferrous ions are incorporated in each molecule of this compound, calculate the molar mass of this compound.

My attempt (which I know is all wrong):

4Fe(y)(X)2  ... the X represents the compound binding to the ferrous Fe and the y is the charge of that compound.

y(55.9 g/mol of iron)/2.7234g of the unknown compound x 100=0.3420
y= 0.0001666

4Fe(2+) + 4X(-y)=4Fe(y)X(2)

ok I'm going to stop there becuase I'm all messed up, any help would be appreciated!

[Borek]

Let's ignore this chromate part, it is completely irrelevant.

a 2.7234g sample of an unknown compound is isolated and analyzed. If this unknown compound is found to contain 0.3420% Fe, and its assumed that 4 ferrous ions are incorporated in each molecule of this compound, calculate the molar mass of this compound.

What is mass of iron in the sample?

How many moles of that iron?

If there are 4 atoms of iron per molecule - how many moles of the compound?

Mass was given, number of moles is known - you are home.

[Julie Smith]

so is it something like this:

to find the grams of the iron in the compound I go 2.7234g x 0.3420% = 0.9314g Fe?

and then 0.9314g Fe x 1mol Fe/55.9g Fe x1mol of compound/4 mols Fe = 0.00417 mols of compound?

and then I take the 2.7234g/0.00417mol = 653g/mol?

please tell me if I'm on the right track

[Julie Smith]

ok so I just realized that in my calculations to get the grams of Fe should be 2.7234g x 0.003420 = 0.00931g of Fe

but then if I carry out the same steps as I previously did just replacing the grams of Fe with this #, the molar mass is 6530 instead of 653g/mol, which is wayyy off I'm assuming

[Borek]

You are still off by order of magnitude, but it will not change anything, final result will look even worse (something like 65300). Generally approach is OK, just your math fails somewhere.

Are you sure it is 0.3420% and not 34.20%?

[Julie Smith]

Are you sure it is 0.3420% and not 34.20%?

ya, I'm sure its 0.3420% of Fe in the unknown compound. So do I need to change anything?

[Julie Smith]

does anyone happen to know the molar mass of hemoglobin? If i'm not mistaken i think it has 0.342% iron in it.

[Borek]

ya, I'm sure its 0.3420% of Fe in the unknown compound. So do I need to change anything?

Just check your math.

does anyone happen to know the molar mass of hemoglobin? If i'm not mistaken i think it has 0.342% iron in it.

Good shot, close to that.

http://en.wikipedia.org/wiki/Haemoglobin

[Julie Smith]

Just check your math.

ya i did mess up on the math. Thanks for the *delete me*

[Julie Smith]

so there were actually2 parts to the question. The first was the one I previously posted:

Quote from: Julie Smith on September 30, 2008, 09:43:01 PM
a 2.7234g sample of an unknown compound is isolated and analyzed. If this unknown compound is found to contain 0.3420% Fe, and its assumed that 4 ferrous ions are incorporated in each molecule of this compound, calculate the molar mass of this compound.

and the answer for the molar mass i found was 65300g/mol.

the question for part b is:

calculate the [ ] if the dichromate solution required to titrate the 2.7234g sample if you wish to reach the end point when 15.02mL of the dichromate solution have been added.

For the balanced reaction there's a 6:1 mole ration between the Fe(2+) and the dichromate.

here's my attempt:

2.7234g of unknown compound x 1mol of unknown compound/65300g unknown compound x 4 mol Fe(2+)/1mol unknown x 1mol dichromate/6mols Fe(2+)
=.0000276mol dichromate/0.01502L dichromate
= .00183 M

is that right?

[Borek]

when 15.02mL of the dichromate solution have been added.

Not sure I get it - does it mean "how much more you must add if 15.02 was already added"?

If it is the same sample you started with, you know its mass and percentage iron content, so you can use these numbers to find out mass and number of moles of iron. In fact you have already did it correctly in you very first post (0.0001668). This way it is much simpler.