1. Calculate the concentration of HCL in the volumetric flask and then the number of excess moles of HCL added to the alkali metal carbonate.
2. Calculate the total number of moles of HCL added to the alkali metal carbonate in Part I. From the total number of moles of HCL added and knowing the number of excess moles, calculate the moles of HCL that reacted with M2CO3.
3. Calculate the number of moles of M2CO3 in your original -0.5g sample.
4. Using the exact mass of M2CO3 used in the experiment, calculate the molar mass of M2CO3.
5. Identify the alkali metal, M. Justify Your assertion. (It is not potassium)
6. In Part I the solution of M2CO3 + HCL is transferred to volumetric flask and diluted to precisely 250.0 mL. Why must this be done? In other words, why is it not possible to titrate aliquots taken directly from the beaker?
7. If solid alkali metal carbonate were spilled during the transfer to the beaker (in part I), how would the calculated molar mass of M2CO3 be affected? Explain.
8. If too much NaOH were added to the HCL during the titration (i.e. the end point is overshot), how would this affect the calculated molar mass of M2CO3? Explain.