When 0.243 g of mg metal is combined with enough hcl to make 100 ml of solution in a constant-pressure calorimeter, the following reaction occurs: mg(s) + 2 hcl(aq) → mgcl2(aq) + h2(g) if the temperature of the solution increases from 23.0 ºc to 34.1 ºc as a result of this reaction, calculate δh in kj/mol of mg. assume that the solution has a specific heat of 4.18 j/g−ºc.

ΔH = - 464 kJ/mol  

Explanation:

Step 1: Data given

Mass of Mg = 0.243 grams

volume of solution = 100 mL

Initial temperature = 23.0 °C

Final temperature = 34.1 °C

Specific heat = 4.18 J/g°C

Step 2: Calculate the heat transfer

Q= m*c*ΔT

⇒ m = the mass of the solution = 100 grams

⇒ c = the specific heat of the solution = 4.18 J/g°C

⇒ ΔT = The change in temperature = T2 - T1 = 34.1 - 23.0 = 11.1 °C

Q = 100 g * 4.18 J/g°C * 11.1°C

Q = 4639.8 J

Step 3: Calculate moles of Mg

Moles Mg = 0.243 grams / 24.3 g/mol

Moles Mg = 0.01 moles

Step 4: Calculate  ΔH

ΔH = 4639.8 J / 0.01 moles

ΔH = 463980 J/mol = 464 kJ/mol

(4.6398 kJ) / (0.0099979 mol) = 464 kJ/mol  

Since the reaction is exothermic, the ΔH is negative by convention, so:

ΔH = - 464 kJ/mol  

i believe that it is a exoplanet

a. energy is absorbed in the reaction.

if in a given reaction the energy of products is greater than the energy of reactants it means that reaction is endothermic and during the reaction, the energy is absorbed by reactants. in an endothermic reaction, the products have more stored chemical energy than the reactants. and in this reaction the reactants are more stable than products before the compounds with less energy are always more stable.