Magnesium reacts with aqueous HCI according to the following balanced equation:

Question

Magnesium reacts with aqueous HCI according to the following balanced equation: Mg(s) + 2HCl(aq) rightarrow H_2(g) + MgCl_2(aq) When 0.0500 g of temperature of the ammonium is added to 100.0 g solution in a coffee cup calorimeter, the temperature of the solution increases from 22.21degree C to 24.46degree C. a) Is this reaction endothermic or exothermic? b) What the heat gained by the solution in this reaction (ie q_solution) in joules? c) How many moles of magnesium are consumed when the reaction proceeds to complete d) What is the Delta H rxn a written above in kJ/mol?

Explanation / Answer

The question here seems confusing.

You have provided a reaction in which Mg is reacting, but the statement written below the reaction is altogether different. Here I am assuming that 0.05 g of Mg is added to 100 g of HCl solution.

(a)

The reaction is exothermic because heat is released which causes increase in the temp.

(b)

qsol = M*C*dT = (100*0.05)*4.20*(24.46-22.21) = 47.25 J

(c)

Moles of Mg consumed = Mass taken / MW = 0.05/24.3 = 0.002 moles

(d)

dHrxn = Heat evolved/Moles of Mg = 47.25/0.002 = 23.625 kJ/mol

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