Answer:
-3108 kJ
Explanation:
Because the combustion is releasing heat, the reaction is exothermic, and because of that, the value of the enthalpy must be negative. When a reaction absorbs heat it is endothermic and the enthalpy is positive.
By the reaction given, there are 2 moles of C₂H₆ reacting. The molar mass of the compound is:
2* 12 g/mol of C + 6* 1 g/mol of H = 30 g/mol
So, 2 moles have 60 g (2mol * 30g/mol).
5.00 g --------------- -259 kJ
60 g --------------- δh
By a simple direct three rule:
5δh = -15540
δh = -3108 kJ
3120 kJ
Given:
At constant pressure, the combustion of 5.00 g of C₂H₅ releases 259 kJ of heat.
Question:
What is ΔH for the reaction given below?
[tex]\boxed{ \ 2C_2H_6_{(g)} + 7O_2_{(g)} \rightarrow 4CO_2_{(g)} + 6H_2O_{(l)} \ }[/tex].
The Process:
Step-1
Step-2
Let us convert mass to mol of 5 g of C₂H₅.
[tex]\boxed{ \ n = \frac{mass}{Mr} \ } \rightarrow \boxed{ \ n = \frac{5}{30.07} \ } \rightarrow \boxed{ \ n = 0.166 \ moles \ }[/tex]
Remember, the combustion of 0.166 moles of C₂H₅ releases 259 kJ of heat.
Step-3
Combustion reactions that release heat into the environment, including the type of exothermic reaction. The exothermic reaction has a negative ΔH because the product enthalpy is lower than the reactant enthalpy of the system.
And now, we will solve the problem of what is ΔH for the reaction given below.
[tex]\boxed{ \ 2C_2H_6_{(g)} + 7O_2_{(g)} \rightarrow 4CO_2_{(g)} + 6H_2O_{(l)} \ }[/tex]
Notice that 2 moles of C₂H₅ take part in this reaction.
[tex]\boxed{ \ \Delta H = \frac{2 \ moles}{0.166 \ moles} \times (-259 \ kJ) \ }[/tex]
[tex]\boxed{ \ \Delta H = 12.048 \times (-259 \ kJ) \ }[/tex]
[tex]\boxed{ \ \Delta H \approx - 3120 \ kJ \ }[/tex]
Thus, the combustion of C₂H₅ based on the equation of reaction above releases heat of 3120 kJ.
