Answer to Question #93868 in Physical Chemistry for Kylie

Question

(2) In an experiment (enthalpy) the following
data were collected:

Initial mass of water in the calorimeter: 70.89 g

Initial temperature of the water in the calorimeter: 17.6 °C

Final mass of water in the calorimeter: 86.48 g

Final temperature in the calorimeter: 0.0 °C

(a) Find the mass of ice that melted.
(b) Calculate the number of mol of ice melted.
(c) What was the change in temperature of the initial mass of water
in the calorimeter?
(d) Calculate the energy released by the initial mass of water. (The
specific heat capacity of water is 4.184 J g–1 °C–1.)
(e) What was the energy absorbed by the ice that melted?
(f) Calculate the energy required to melt 1 mol of ice.

Answer 1

(A) The mass of ice that melted =$Final \ mass \ of\ water- \ initial\ mass\ of\ water$ $=$ $86.48-70.89=15.89\ g$

(B) Moles of ice that melted = $\frac{mass \ of \ ice\ melted}{molar\ mass \ of\ ice}=\frac{15.89}{18}=0.8661\ mole$

(C) Change in temprature of water = |final temprature $-$ initial temprature|=$17.6\degree\ C$

(D) Energy released = $m_1c(T_2-T_1)=$ $70.89\times 4.184\times ({-17.6})=-5220.2\ J$

$(-) sign\ means\ that \ energy\ is\ released$

$m_1=initial \ mass \ of water$

$c=4.184\ Jg^{-1}\degree C^{-1}$

(E) Energy absorbed by ice that melted $=m_2L=15.89\times 334=5307.26\ J$

$m_2=mass \ of \ ice\ that\ melted\\L=latent \ heat=334\ Jg^{-1}\degree C^{-1}$

(F) Energy required to melt 1 mole of ice $=ML=18\times$ $334=6012\ J$

$M=molar\ mass \ of \ ice=18\ g\ mol^{-1}$