Heating of Magnesium Lab

• After experiment (after cooled):

• Grainy substance, white/grey chalky residue & streaks on bottom of crucible, some of the substance was stuck onto the crucible, residue created a spiral design on the bottom of crucible

Quantitative Observations:

Calculations:

Empirical Formula Work

Mass of MgO

Mass of magnesium = 0.136 g ± 0.01g

Mass of MgO = (Mass of MgO + crucible) – mass of crucible

= 25.650 ± 0.01g – 25.429g ± 0.01g

= 0.221 ± 0.02g

Mass of Oxygen= Mass of MgO – Mass of Magnesium ribbon

                  = 0.221g ± 0.02g – 0.136 g ± 0.01g

                  = 0.085g ± 0.03g

Percent composition of a 100g sample

Magnesium                                                 Oxygen 

Mass of Mg

Mass of MgO

0.136g ± 0.01g                                   Oxygen = 100% - percent composition of Mg 

0.221g ± 0.02g                                                   =   100% - 38%  

= 62% → 62g of 100g sample              = 38% → 38g of 100g sample

Molar Ratio

Moles (n) =       mass (m)

        Molar mass (MM)

Magnesium                                Oxygen

n= 62g/24.3050g/molecule                        n= 38g/15.9994g/molecule

n= 2.6 moles                                                           n= 2.4 moles

=2.6 moles: 2.4 moles

= 1.1:1              Empirical Formula

=Mg1.1O       

Molar Equivalency

M.E. =       moles (n)

                 Balancing #

Magnesium                                Oxygen

M.E. = 2.6 moles / 2 units                        M.E. = 2.4 moles / 1 unit

M.E. = 1.3 moles / unit                                                           n= 2.4 moles / unit

Limiting Reagent                                                Excess Reagent

Conclusion Questions

1. The type of reaction that occurred between the magnesium ribbon and oxygen is a synthesis reaction.
2. Mg1.1O      

      3)  2Mg(s) +    O2 (g)                 ∆             2MgO(s)    

4) In this experiment, the oxygen acted as the excess reagent as it had a higher value of moles per unit than the magnesium.  This occurred because during the experiment there was a limited supply of magnesium (i.e. 10 cm) but there was an abundant amount of oxygen in the surrounding environment which could be used up in the process.

5) Percent Error

2Mg(s)                              +              O2 (s)              ∆                              2MgO(s)          

Mass= 0.136g                                                                     mass= moles (n) x molar mass

Molar mass= 24.3050 g/mol                                              n= M.E. x bal #

n= m / mm                                                                          n= 0.0028moles / units x 2 units  

n= 0.136g / 24.3050 g/mol                                                 n= .0055 moles  

n= 0.0056 mol                          

                                                                                          m= n x mm

ME= moles/ Bal #                                                             m=0.0056moles x 40.30 g/mol

      = 0.0056 mol / 2 units                                                 actual mass= 0.23 g      

      = 0.0028 mol/units

Experimental value – Actual value x 100

                Actual value

= 0.221g - 0.22g x 100

           0.22g

= 0.45% error

6)Percent Yield

Yield= Experimental yield   x 100

              Theoretical yield

Experimental yield→Mass of MgO= 0.221g

Theoretical yield-→Mass of MgO= 0.22g

= (0.221g/0.22g) * 100

= 100.5% yield

In conclusion the empirical formula for MgO proves that the lowest ratios of the atoms of oxygen and magnesium have close to the same ratio. Any discrepancies were possibly a result of an error in some of the equipment that was used to weigh the magnesium and oxygen. This could explain why the empirical formula was 1:1.04 as there is a chance that the preciseness of the scale might have been incorrect, also next time the lab is done, the mass of the lid after the experiment should be taken into account as the magnesium oxide might have been evident upon it as well as there were grey streaks after the experiment on the lid.  Furthermore, if we did the experiment more times (held more trials) we could’ve got an average or at least made sure the behavior we noticed stayed consistent.