What is the final temperature when 1.192 kj of energy is added to 10.0g of ice at 0.00C?

You have from the First Law of Thermodymamics :

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Q = [ Delta U ] + [ W sub B ]

Q = 1.192 kj

Delta U = ( m ) ( L sub F ) + ( m ) ( C sub V ) ( TF - TI )

Delta U = ( 10.0 ) ( 0.3337 ) + ( 10.0 ) ( 0.00418 ) ( TF - 0.0 )

Delta U = 3.337 + ( 0.0418 ) ( TF - 0.0 )

W sub B = 0.0 kj ...... no work is done

Q = Delta U

1.192 = 3.337 + ( TF - 0.0 )

TF = 0.0 deg C <-----------------------------------------------------------------

The heat added is not enough to melt all of the ice and you will have a mixture of ice and water

at 0.0 deg C.

Q greater than 3.337 kj would be needed to melt all of the ice and start increasing the

temperature above 0.0 deg C.

The equation bearing on specific warmth to temperature replace is Q = m*c*deltaT the place Q is your power (a million.57 kJ), m is the mass of your cloth (10.0g), c is the specific warmth of water (4.186 J/gC), and delta T is the replace in temperature in celcius. you are able to now plug in those values, remedy for deltaT, and make certain the ideal temp (i.e. 0 + deltaT).

Enthalpy = -mass(kg) x heat capacity x change in temperature

-1.192 = -0.01 x 4.18 x T

-1.192 = -0.0418 x T

28.52 = T

Since initial Temperature is 0 then final temp is

= 0 + 28.52

= 28.52 Degrees Celsius