Question

Question #a3d65

Answer 1

The reaction's percent yield will be `"71.8%"`.

The balanced equation for the decomposition of magnesium chloride is

`MgCl_(2(l)) -> Mg_((l)) + 2Cl_((g))`

Every time you have to compare the amount of a product produced given an amount of a reactant, you must use mole ratios. The mole ratio between magnesium chloride and magnesium metal will help you determine how much of the latter would have been produced if the reaction had a `"100%"` yield.

This will be the theoretical yield of the reaction. The actual yield of the reaction was given to you - `"1000 g"` of magnesium chloride produced `"185 g"` of magnesium metal.

Determine the theoretical yield of the reaction by using the `"1:1"` mole ratio between magnesium chloride and magnesium metal. One mole of `MgCl_2` produces one mole of `Mg`.

Use molar masses to go from grams to moles and vice versa.

`"1000 g MgCl"_2 * ("1 mole MgCl"_2)/("94.3 g") = "10.6 moles"` `MgCl_2`

This means that you have

`"10.6 moles MgCl"_2 * "1 mole Mg"/("1 mole MgCl"_2) = "10.6 moles Mg"`

The number of grams produced in theory will be

`"10.6 moles Mg" * ("24.3 g")/("1 mole Mg") = "257.6 g"`

Now, percent yield is defined as the ratio between actual yield and theoretical yield, multiplied by 100. All you have to do from this point on is to plug the values you have into the equation

`"%yield" = "actual"/"theoretical" * 100 = "185 g"/"257.6 g" * 100 = "71.8%"`

Percent yield is all about what you would produce in theory vs . what you produce experimentally. A reaction can NEVER have a percent yield that's greater than `"100%"`. NEVER. A percent yield that exceeds this value is automatically a red flag and a clear sign that you've made mistakes in your calculations.