At constant pressure, the combustion of 5.00 g of C2H6 (g) releases 259 kJ of heat. What is ΔH for the reaction?
2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(l)

Question

Peyton Adams · Accepted Answer

The enthalpy change of combustion for that reaction will be #"-3120 kJ"#.

Commence by using the provided balanced chemical equation.

#2C_2H_(6(g)) + 7O_(2(g)) -> 4CO_(2(g)) + 6H_2O_((l))#

You'll see that this reaction involves two moles of ethane.

You were given the amount of heat released by burning 5.00 g of ethane. Using the molar mass of ethane, calculate the number of moles of ethane that must burn in order to release that amount of heat.

#"5.00"cancel("g") * "1 mole"/("30.07"cancel("g")) = "0.1663 moles"#

Since two moles of ethane must burn for the given reaction to occur, the amount of heat released will be

#2cancel("moles") * ("-259 kJ")/(0.1663cancel("moles")) = "-3114.85 kJ"#

When the number of sig figs for 5.00 g and 259 kJ is rounded to three, the result is

#DeltaH_("comb") = color(green)("-3120 kJ")#

Wyatt Adkins · Answer

-3120 kJ is the heat of combustion.

You are aware that 5 grams of C3H2 emit 259 kilojoules of heat.

You need to figure out how much heat 2 mol of C₂H₆ releases.

#2 cancel("mol C₂H₆") × "30.07 g C₂H₆"/(1 cancel("mol C₂H₆")) = "60.14 g C₂H₆"#

#ΔH = 60.14 cancel("g C₂H₆") × "-259 kJ"/(5.00 cancel("g C₂H₆")) = "-3120 kJ"#

Wyatt Atkins · Answer

undefined To find ΔH for the reaction, you need to use the given heat released and the stoichiometry of the reaction. First, calculate the moles of C2H6 using its molar mass. Then, use the stoichiometry of the reaction to find the moles of heat released for the reaction as a whole. Finally, divide the heat released by the moles of the reaction to find the heat released per mole. This will be the ΔH for the reaction.

At constant pressure, the combustion of 5.00 g of C2H6 (g) releases 259 kJ of heat. What is ΔH for the reaction? 2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(l)