If 15.0 kJ are released when 1.40 g of F2 reacts with an excess of SiH4, complete the thermochemical equation below.
SiH4(g) + 2 F2(g) SiF4(g) + 4 HF(g)
ÄHrxn = ?kJ
I worked this problem for you before.
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To find the enthalpy change (ΔHrxn) for the given thermochemical equation, you need to use the given information about the amount of energy released (15.0 kJ) and the mass of fluorine gas (F2) reacted (1.40 g).
Here are the steps to find ΔHrxn:
Step 1: Calculate the moles of F2 reacted.
To do this, you need to convert the mass of F2 to moles using its molar mass.
The molar mass of F2 = 2(atomic mass of F) = 2(19.0 g/mol) = 38.0 g/mol.
Number of moles of F2 = mass of F2 / molar mass of F2
Number of moles of F2 = 1.40 g / 38.0 g/mol ≈ 0.037 moles
Step 2: Use the balanced equation to find the moles of SiH4 reacted.
From the balanced equation: 1 mole of SiH4 reacts with 2 moles of F2.
Since the balanced equation shows a 1:2 ratio between SiH4 and F2, the moles of SiH4 reacted will be the same as the moles of F2 reacted.
Number of moles of SiH4 = 0.037 moles
Step 3: Calculate the enthalpy change (ΔHrxn)
The enthalpy change can be calculated by using the formula:
ΔHrxn = q / n
Where:
ΔHrxn = enthalpy change (in kJ)
q = heat released or absorbed (in kJ)
n = number of moles of the reactant in the balanced equation
ΔHrxn = (-15.0 kJ) / (0.037 moles)
ΔHrxn = -405.4 kJ/mol
So, the enthalpy change (ΔHrxn) for the reaction is approximately -405.4 kJ/mol.