When 0.767 g of Ca metal is added to 200.0 mL of 0.500 M HCl(aq), a temperature increase of 125C is observed.
Assume the solution's final volume is 200.0 mL, the density is 1.00 g/mL, and the heat capacity is 4.184 J/gC.
(Note: Pay attention to significant figures. Do not round until the final answer.)
The molar heat of reaction, H rxn, for the reaction of
Ca(s) + 2H+(aq) Ca2+(aq) + H2(g)
is.....
q = mass water x specific heat water x delta T
q/0.767 = J/g
Convert to kJ/mol.
I did what you said and it didn't work Dr Bob. Can you explain step by step?
To find the molar heat of reaction (Hrxn) for the given reaction, we need to calculate the amount of heat transferred during the reaction. We can use the formula:
q = mcΔT
Where:
q = heat transferred (in joules, J)
m = mass of the solution (in grams, g)
c = specific heat capacity of the solution (in J/g°C)
ΔT = change in temperature (in °C)
First, we need to find the mass of the solution. Given that the density of the solution is 1.00 g/mL and the final volume is 200.0 mL, we can calculate the mass as follows:
mass = density × volume
mass = 1.00 g/mL × 200.0 mL
mass = 200.0 g
Next, we can substitute the values into the formula and calculate the heat transferred:
q = (200.0 g) × (4.184 J/g°C) × (125°C)
q ≈ 1048000 J
Now that we have the amount of heat transferred, we can calculate the molar heat of reaction (Hrxn). The balanced equation tells us that 1 mole of Ca reacts with 2 moles of H+ ions. Therefore, the number of moles of Ca reacting can be calculated using the given mass of Ca (0.767 g) and the molar mass of Ca (40.08 g/mol):
moles of Ca = mass of Ca / molar mass of Ca
moles of Ca = 0.767 g / 40.08 g/mol
moles of Ca ≈ 0.0191 mol
Since 1 mole of Ca reacts, the heat transferred is equal to the molar heat of reaction (Hrxn). Therefore:
Hrxn = q / moles of Ca
Hrxn = 1048000 J / 0.0191 mol
Hrxn ≈ 54761592.15 J/mol
Remember that we need to pay attention to significant figures. Given the significant figures in the given values (0.767 g, 200.0 mL, 0.500 M), the final answer should be reported with appropriate significant figures.