1. It is really confusing.....

Why we do not use tha actual masses of atoms in our calculations? even if they are small...
I mean for example in chemical industries , how using relative masses does not make difference????  2.

3. I'm not sure what you mean. The masses that we use in calculations (that is to say, what you see on a periodic table) is the exact average mass of the atom. If you have 1 mol of nitrogen, you have exactly 14.0067 grams of nitrogen.  4. There is no such thing as an actual mass for an element. Most elements have several stable isotopes. The listed number of the mass of an element is the average mass of the element based on isotopic abundance. Using that number in calculations is the only practical thing to do, and really if you went along and calculated the mass for each isotope you'd end up with the same number anyway.

And I'd like to draw your attention to something to illustrate. The mass of an atom is measured in AMUs (atomic mass units) where each proton and each neutron are each equal to 1. Electrons are too small to have any effect on mass in any practical measurment. So lets look at boron, B. It has 5 protons, and with 5 neutrons that would make it's atomic mass 10. But it's atomic mass is not 10, it's 10.81. You've never going to find a B that is 10.81 AMUs (it's not possible). The average of its istopes work out to an average 10.81 AMUs.

Does that make sense?  5. Just tell me

How mass of one mole is equal to the relative atomic mass

In the spectrometer , for instance , which is used to determine the RAM of a substance we do not use 1 mole of that substance.....

So how the mass that i determined from this mechine is the mass of 1 mole.......  6. A mass spectrometer does not measure the mass of a mole of a substance, it measures the mass of the actual atom as the molecule breaks apart. Mass spec is also used to determine isotopic abundance and are where atomic masses come from.

Long story short, with mass spec you don't have to use the mole of a substance to determine its atomic mass because it isn't doing a gravimetric analysis or anything like that, it is measuring the mass of the actual molecule via flight times, etc.

Also as in why does the mass in grams = 1 mole of the AMU mass?

It's because that's how a mole is defined. You can actually do it for anything. You could weigh everything in pounds and have mole pounds. You could weigh everything in tons and have mole tons. Ounces, mole ounces. For example 12.01 grams of C would be a mole. 12.01 pounds of C would be a mole pound. It's all very relative, as long as you use the same units you can get the same amount of chemical substances as long as you measure them in the amount relative to their atomic masses.  7. You r right my friend in the spectrometer we measure the mass of the atom....
first we find out the total mass of the isotopes then devided by the number of the particles ( getting from the abundance) ... so we get the averge mass of the atom....

For C12
1 mole = 6.02 times 10^23 atoms = 12 g ,,,,ok??

and here in the spectrometer i am measuring the mass of one atom not ( 6.02 times 10^23 atoms)

So hoe they r the same......  8. Okay?

They work out the same because what your doing in mass spec is measuring AMUs. As I tried to point out earlier, it doesn't have to be measured in grams in order to have a chemical equivalent of a substance, it's just AMUs in grams = Avogadro's number of atoms or a mole. That's the most practical thing to calculate.

However, because AMUs talk about the mass of one atom, and the isotopic abundances give us an idea of how they appear, you can will the same number of atoms as long as you measure the atoms on the same scale, whatever scale you choose.

Does that make sense? Is that what you're asking? I'm just repeating myself now, but I do know what I'm talking about here.  9. I understood every thing .......

Thanks for help.........  Bookmarks
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