Chemical Forums
Chemistry Forums for Students => High School Chemistry Forum => Topic started by: 777888 on January 13, 2005, 10:51:55 AM
-
1. Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)
100.0mL of 1.0 M HCl(aq) is placed in a coffee cup. Ti=20.8C. 0.50g of Mg is added and Tf=44.8C. Determine the heat of reaction.
Work:
Energy absorbed by solution=100x4.18x24
(Can I ASSUME it has specific heat capacity of 4.18J/gC ?)
q=10.032kJ
Energy released by system=10.032kJ
delta Hx=-10.032kJ / 0.02057mol=-4.9x10^2kJ/mol Mg (But my answer is wrong, why? Please help me out...)
(Is the question asking for molar enthalpy or enthalpy change?)
I need *delete me* Thank you! :)
-
Do I have to do the limiting reagent thing and how?
Please help...
-
yes. you gotta find use the "limiting agent" thingy because not everything has reacted. deltaH is per mole of reactant used up.
-
yes. you gotta find use the "limiting agent" thingy because not everything has reacted. deltaH is per mole of reactant used up.
I found that Mg is limiting reagent...then are my calculations correct?
thanks :)
-
2. MgO(s)+2HCl(aq)->MgCl2(aq)+H2O(l)
100.0mL of 1.0mol/L HCl(aq) is placed in a coeffee cup(Initial temp.=20.4C). 1.00g MgO(s) is added and the final temp.=29.6C
Work:
Surroundings:
q=(100g)(4.18J/gC)(9.2C)=3.8456kJ
System:
MgO(s) + 2HCl(aq) -> MgCl2(aq)+H2O(l)
1.00g 0.1L
0.0248mol 0.1M
0.1mol
MgO(s) is limiting reagent.
Energy released=-3.8456kJ
Delta Hx = -3.8456kJ/0.02481mol
= -155.00kJ/mol MgO
= -1.6x10^2kJ/mol MgO
Can someone see if calculate it right? Just to make sure I get this stuff...Many thanks :)
-
Looks good to me. You can also check for yourself if you did it properly by looking up the delta Heat of formation for solid MgO, aqueous HCl, Aqueous MgCl2, and liquid water. Then just take the sum of the products' heat of formation and subtract the the sum of the reactants' heat of formation. That should give you the theoretical heat of reaction.
-
Looks good to me. You can also check for yourself if you did it properly by looking up the delta Heat of formation for solid MgO, aqueous HCl, Aqueous MgCl2, and liquid water. Then just take the sum of the products' heat of formation and subtract the the sum of the reactants' heat of formation. That should give you the theoretical heat of reaction.
When using your method to find "the theoretical heat of reaction", would that be in kJ or kJ/mol? If it is in kJ, how can I change it to kJ/mol and compare to my answer?
I have one more question, when it appears like below, what is the kJ/mol refering to? kJ/mol O2? kJ/mol H2O? Or...?
H2(g)+ 1/2 O2(g) -> H2O(l) Delta Hx=-285.8kJ/mol
Thanks for your *delete me*
-
Looks good to me. You can also check for yourself if you did it properly by looking up the delta Heat of formation for solid MgO, aqueous HCl, Aqueous MgCl2, and liquid water. Then just take the sum of the products' heat of formation and subtract the the sum of the reactants' heat of formation. That should give you the theoretical heat of reaction.
Hi judrg, my text book doesn't have standard enthalpies for Mg(s), HCl(aq), MgCl2(aq), and H2(g)! Where can I find them? Thank you!
-
When using your method to find "the theoretical heat of reaction", would that be in kJ or kJ/mol? If it is in kJ, how can I change it to kJ/mol and compare to my answer?
I have one more question, when it appears like below, what is the kJ/mol refering to? kJ/mol O2? kJ/mol H2O? Or...?
H2(g)+ 1/2 O2(g) -> H2O(l) Delta Hx=-285.8kJ/mol
Thanks for your *delete me*
In the equation you gave, it's stating that -285.5 kJ are given off for every mole of water that is formed. There are no standard heats of reaction tables. The heat of reaction has to be calculated from the heats of formation. What you do see tables of are heats of formations of compounds. A pure element in its standard state has no heat of formation value. So H2, O2, Na, Mg, Cl2, etc. etc. have no delta heat of formation.
For your second question, the best place to look for heats of formations are the CRC Handbook, or the internet. (Though the internet can be kind of tough). Since pure elements in their standard states have no enthalpies of formation, the values for Mg(s) and H2(g) are 0. You will just need to look up HCl(aq) and MgCl2(aq). (If you can't find the aqueous heats of formation, you may be able to just use the pure compound heats of formation as I don't think there is a whole lot of difference, unless the dissolution in water releases a good deal of heat).