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<div class="moz-cite-prefix">On 06/25/13 19:35, wrote:<br>
</div>
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cite="mid:C40B2F181831EF44A88CD735258278030261C418D2@MERCERMAIL.MercerU.local"
type="cite">
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<div>Hey Dr. Pounds, </div>
<div>I'm working on my calculations for lab. I was able to find
the empirical formula with ease, but I'm not sure how to go
about finding the molecular formula. I have added up the formula
mass, but won't I need a molecular mass to compare it to? I only
have the average molar mass that I found via experiments. <br>
</div>
</blockquote>
<br>
That is correct -- you have a MOLAR mass that you determined from
your experiment today (using the mass of the vapor divided by the
moles of vapor determined from <img style="vertical-align: middle"
src="cid:part1.05080701.07050902@mercer.edu"
alt="$n=\frac{PV}{RT}$">). Now, lets for the sake of argument
let's imaging that your empirical formula (detemined from the
elemental analysis) is <img style="vertical-align: middle"
src="cid:part2.02040700.05060500@mercer.edu" alt="$\mathrm{CH_4}$">
and that your molar mass from the expement was 33 g/mol. The molar
mass of <img style="vertical-align: middle"
src="cid:part2.02040700.05060500@mercer.edu" alt="$\mathrm{CH_4}$">
is 16 g/mol -- so your MOLECULAR formula would be <img
style="vertical-align: middle"
src="cid:part4.04060104.00020508@mercer.edu"
alt="$\mathrm{C_2H_8}$"> because this is the molecular formula
that is an whole number multiple of the empirical formula that is
has a molecular mass (32 g/mol) close to your experimentally
determined molecular mass.<br>
<br>
<br>
<pre class="moz-signature" cols="72">--
Andrew J. Pounds, Ph.D. (<a class="moz-txt-link-abbreviated" href="mailto:pounds_aj@mercer.edu">pounds_aj@mercer.edu</a>)
Professor of Chemistry and Computer Science
Mercer University, Macon, GA 31207 (478) 301-5627
<a class="moz-txt-link-freetext" href="http://faculty.mercer.edu/pounds_aj">http://faculty.mercer.edu/pounds_aj</a>
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