Hey kids! Now you can perform a titration lab in your very own home!
The purpose of this lab is to determine the concentration
of a reagent
, using an indicator solution and a reagent whose concentration is known.
First, we must prepare an exact volume of our known solution, to which the unknown will be added. Fill your pipette (look under that node for proper pipetting procedure) and drain it completely into the Erlenmeyer flask. Add a few drops of your indicator solution—how many depends on how much solution you're titrating, but about 10 should do it—and you're ready to go.
Pour your unknown solution into the buret. You needn't fill the buret up to the top, as long as you record the reading, but if you fill it all the way (to 0, since the buret measures liquid dispensed) then the subtraction is easier. Look closely at the bottom of the buret (down where the stopcock is), and flick away any air bubbles that have formed. Clamp the buret over the flask and you're ready to begin.
Turn the stopcock so that the knob is parallel to the buret to begin dispensing solution, swirling the flask to ensure that the reagents are reacting. You will note that the solution in the flask will turn purple (if you're using phenolphthalein; I'm not sure what color you'll get with other indicators) briefly and then return to clear. This is because the phenolphthalein turns pale purple when present in a roughly neutral solution (pH=7 or thereabouts; any lower and the solution is clear, any higher and it turns red). But the purple rapidly vanishes as the flask is swirled, because the overall pH is still too low for the solution to change color.
As more and more titrant is added, the solution will remain pale pink for longer periods of time. Slow down the rate at which you add the solution—just a few drops at a time. When the solutions stays pink (the color you're looking for is somewhere between violet and plum) even after being swirled for a few seconds, you're done. Record the final volume in the buret.
Note: these numbers have all been made up. You have been warned.
Known-concentration solution (1M HCl):
Unknown-concentration solution (?M NaOH):
- initial buret reading: 0.3 mL
- final buret reading: 7.8 mL
total NaOH dispensed: 7.8-0.3 = 7.5 mL
reaction equation: NaOH(aq) + HCl(aq) --> NaCl(aq) + H2O(l)
Since final pH = 7, we see that
moles of HCl = moles of NaOH dispensed
moles of HCl = M· L = 1M· 0.015L = 0.015 mol HCl
0.015 mol NaOH = M· L = M· 0.0075L
M = 0.015 / 0.0075 = 2M NaOH
And, just for grins, let's find the pH (go there for a full explanation of these equations):
reaction for NaOH dissolving in water: NaOH(s) --> Na+(aq) + OH-(aq)
M(NaOH) = M(OH-) = 2
pOH = -log[OH-] = -log(2) = -0.3
pH = 14-pOH = 14.3
Source: my own lab work. This is only one procedure, of course; many variations are possible. If you've used a radically different procedure from mine (or if I'm flat wrong in my calculations, always a possibility), drop me a /msg.