to be hydrated or ANhydrated? (LAB 4C)

So, over the course of two classes, we've once again did a lab!! This time, it was all about finding the empirical formula of a hydrate!

Now to give you a better understanding of this lab, here are some background information:

• A hydrate is a compound that contains a definite number of water molecules (H2O) 
• Without containing water, it becomes an anhydrous (without water) salt
• The formula of a hydrate is AB•xH2O where AB is the anhydrous salt and x is the number of moles of water per mole of the anhydrous salt

**always use crucible tongs and never touch or eat the substance!**


Here is the hydrate that we used:

Now in order to find the empirical formula, we have to find the mass of water that's in the hydrate by finding how much water was given off by burning the compound.
BUT WAIT! Before that, we must:
1. Find mass of empty crucible
2. Find mass of crucible + hydrate (before it's burned)
3. Subtract the mass in step 2 from the mass in step 1 to get mass of HYDRATE
Then, we heat the hydrate in the crucible until the substance changes to a white-ish colour. After letting it cool (SAFETY FIRST OF COURSE! - don't want to get your fingers burned) we weigh the crucible + anhydrous salt (as the water has been dried off). The mass should now be lighter than it was before. We do this process twice to make sure the water is completely gone! (the two masses should be within 0.03g)

(anhydrous salt of copper sulphate)

Next:
4. Using lowest mass of the crucible + anhydrous, subtract the mass of step 1 from this to get the mass of ANHYDROUS salt
5. Subtract the mass of the anhydrous salt from the mass of the hydrate to get mass of water that WAS in the compound


Given that the molar mass of the anhydrous salt (CuSO4) is 159.6g/mol, we can find the # of moles of the anhydrous.
mass of anhydrous salt x 1mol/159.6g = moles of anhydrous salt

Recall (from our quiz) that we have to find the ratio of moles of water to moles of anhydrous. Therefore, we find the moles of water from the molar mass of H2O:
mass of water given off x 1mol/18.0g = moles of H2O


Then, divide the two moles by the smallest molar amount (does this sound familiar now? :D)  and VOILA! You should get a whole-number ratio.


Thus, the empirical formula of this hydrate is CuSO4•5H2O (copper (II) sulphate pentahydrate)

Also, if you try adding water into the crucible with the anhydrous, an exothermic reaction will occur along with the changing of the color of the substance to the original hydrate color. Neat huh!

(teehee, goodluck!)