Astrip of magnesium metal having a mass of 1.30 g dissolves in 150. ml of 5.00 m hcl (specific gravity = 1.10); products of the reaction are magnesium chloride and hydrogen gas. the hcl is initially at 22.0°c, and the resulting solution reaches a final temperature of 52.0°c. the heat capacity of the calorimeter in which the reaction occurs is 345 j/°c. calculate δh (in kj/mol) for the reaction under the conditions of the experiment, assuming the specific heat of the final solution is the same as that for water (4.184 j/g°c).

Explanation:

The given data is as follows.

  Mass of Mg = 1.30 g,   Molarity = 5 for HCl,        Volume = 150 mL

   Density = 1.1 g/mL

  Change in temperature (dT) =

              = 345

Hence, the energy balance will be as follows.

Moles of Mg =

                 =

                 = 0.05349 mol of Mg

                      =

                      = 10350  J

                         =

                         = 20710.8

              = (20710.8 + 10350) J

                  = -31060.8 J

Therefore, calculate the value of  of the reaction as follows.          

                      = -580684.24 J/mol

or,                        (as 1 kJ = 1000 J)

Thus, we can conclude that value of  of the reaction under given conditions is -580.68 kJ/mol.

the ideal gas law equation is as follows

pv = nrt

where p - pressure

v - volume

n - number of moles

r - universal gas constant

t - temperature

so if the pressure, volume and temperature are already known

we are left with n and r

since r is the universal gas constant that has a known fixed value then r too is known

so we are left with 'n'

once we know temperature volume and pressure

we can find the number of moles of gas present using the ideal gas law equation

this is not my own answer, this is an answer for the same question by the user blahblahmali. i do not claim ownership of this account.