Asked by Ronnie
An average hot water tank will hold 300.0 L of water needs to be kept at 65.0C in order to kill any bacteria in the water. If you are filling a new tank with water at a starting temperature of 12.5C, what mass of natural gas
(methane) must be burned in a complete combustion reaction in order to bring your tank up to the correct temperature?
I think I mostly need help with identifying the equation needed to solve this but if anyone is able to fully explain it, it would be greatly appreciated
(methane) must be burned in a complete combustion reaction in order to bring your tank up to the correct temperature?
I think I mostly need help with identifying the equation needed to solve this but if anyone is able to fully explain it, it would be greatly appreciated
Answers
Answered by
DrBob222
You have three problems here.
1. First you need to determine how much heat is needed to raise the temperature of 300.0 L of water at 12.5 C to a final temperature of 65 C.
2. Then you need to write the equation for the combustion of methane and calculate how much heat 1 gram of methane (or some other quantity) produces. Then you can calculate how many grams of methane are needed to produce what you need in #1.
3. You DON'T have the heat produced from the combustion of methane. I don't know if you didn't have that in the problem or if you just didn't type it in. I assume also that the specific heat of water is given or you know it and that you know the density of water at that temperature is 1.00 g/mL.
#1. q = heat needed = mass H2O x specific heat H2O x (Tfinal - Tinitial)
q = 300,000 g x 4.184 J/g*C x (65.0 - 12.5) = approximately 66,000,000 J or 66,000 kJ. You should confirm this.
#2. CH4 + 2O2 ==> CO2 + 2H2O
I looked up the heat of combustion for CH4 on th web and found 889 kJ/mol CH4. You should use the value in your text/notes/discussions in class. So 889 kJ/mol is 889 kJ/16 g CH4. You want 66,000 kJ from part 1 and you have reaction in part 2 that will produce 889 kJ/16 g. So
16 g CH4 x 66,000 kJ/889 kJ = ? g CH4.
Hope this helps. Post your work if you get stuck.
1. First you need to determine how much heat is needed to raise the temperature of 300.0 L of water at 12.5 C to a final temperature of 65 C.
2. Then you need to write the equation for the combustion of methane and calculate how much heat 1 gram of methane (or some other quantity) produces. Then you can calculate how many grams of methane are needed to produce what you need in #1.
3. You DON'T have the heat produced from the combustion of methane. I don't know if you didn't have that in the problem or if you just didn't type it in. I assume also that the specific heat of water is given or you know it and that you know the density of water at that temperature is 1.00 g/mL.
#1. q = heat needed = mass H2O x specific heat H2O x (Tfinal - Tinitial)
q = 300,000 g x 4.184 J/g*C x (65.0 - 12.5) = approximately 66,000,000 J or 66,000 kJ. You should confirm this.
#2. CH4 + 2O2 ==> CO2 + 2H2O
I looked up the heat of combustion for CH4 on th web and found 889 kJ/mol CH4. You should use the value in your text/notes/discussions in class. So 889 kJ/mol is 889 kJ/16 g CH4. You want 66,000 kJ from part 1 and you have reaction in part 2 that will produce 889 kJ/16 g. So
16 g CH4 x 66,000 kJ/889 kJ = ? g CH4.
Hope this helps. Post your work if you get stuck.
Answered by
Naeem R
Thank you DrBob222, really appreciate it!
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