Question #2cc3c

1 Answer
Jul 7, 2017

Answer:

Here's what I got.

Explanation:

The idea here is that the heat given off by the combustion of the sample will be equal to the heat absorbed by the calorimeter.

The calorimeter has a heat capacity of #"3024 J" ""^@"C"^(-1)#, which means that it takes #"3024 J"# of heat to increase its temperature by #1^@"C"#.

In your case, the temperature of the calorimeter increased by #3.337^@"C"#, which implies that the calorimeter absorbed

#3.337 color(red)(cancel(color(black)(""^@"C"))) * "3024 J"/(1color(red)(cancel(color(black)(""^@"C")))) = "10,091.1 J"#

Convert this to kilojoules

#"10,091.1" color(red)(cancel(color(black)("J"))) * "1 kJ"/(10^3color(red)(cancel(color(black)("J")))) = "10.091 kJ"#

Now, you know that the calorimeter absorbed #"10.091 kJ"# of heat, so it must mean that the reaction gave off #"10.091 kJ"# of heat.

Moreover, you know that this much heat was given off when #"0.4075 g"# of magnesium underwent combustion. You can thus say that the heat released when #"1 g"# of magnesium undergoes combustion is equal to

#1 color(red)(cancel(color(black)("g"))) * "10.091 kJ"/(0.4075 color(red)(cancel(color(black)("g")))) = "24.76 kJ"#

Since this heat is being given off, you can say that the enthalpy change of combustion of magnesium will be equal to

#DeltaH_"comb" = color(darkgreen)(ul(color(black)(- "24.76 kJ g"^(-1))))#

The minus sign is used to symbolize heat given off.

Finally, to convert this to kilojoules per mole, use the molar mass of magnesium

#-24.76 color(white)(.)"kJ"/color(red)(cancel(color(Black)("g"))) * (24.305 color(red)(cancel(color(Black)("g"))))/("1 mole Mg") = color(darkgreen)(ul(color(black)(-"601.8 kJ mol"^(-1))))#

The answers are rounded to four sig figs.