If no heat is lost to the surroundings, what is the final temperature of the mixture?

asked by @student63603 • about 1 year ago • Chemistry • 5 pts

Four ice cubes at exactly 0 ∘C with a total mass of 50.0 g are combined with 135 g of water at 90 ∘C in an insulated container. (ΔH∘fus=6.02 kJ/mol, cwater=4.18J/g⋅∘C).

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1 answer

Hey there. So for this question we first need to calculate the amount of heat the ice cubes must absorb in order to melt. Remember that during a phase change that q = m x delta H.

Heat to melt the ice cubes:

50.0 g x (1 mole / 18.016 g) x (6.02 kJ / mol) x (1000 J / 1 kJ) = 16707.4 kJ

The heat of the 135 g of water:

q = m x c x delta T = (135 g) (4.18 J / g x C) (90 - 0 C) = 50787 J

The amount of heat remaining:

50787 J - 16707.4 J = 34079.6 J

Now that we have the difference in heat absorbed versus released we can deal with the final temperature of the solution.

qsolution = m x c x delta T 34079.6 J = (50.0 g + 135.0 g) x (4.18 J / g x C) (Tfinal - 0) 34079.6 J = 773.3 (Tfinal - 0) 44.07 Celsius = Tfinal

Hope that helps.

answered by @jules • about 1 year ago