A 30 g ice cube at 0 °C is put in an equal mass of 98 °C water in an insulated container, and allowed to melt. After a while the resulting mixture is at a final (uniform) temperature. How much heat energy does the ice require to completely melt at 0 °C? What is the uniform temperature of the resulting 'mixture'? Note: assume that the heat capacity of water and ice are independent of temperature, and are 4200 J/kgK and 2100 J/kgK, respectively. The heats of fusion and vapourization of water are 3.33x10° J/kg and 2.35×10° J/kg, respectively. How much heat energy does the ice require to completely melt at 0 °C? Answer = What is the resulting temperature of the total liquid (in °C)? Answer = "C

A 30 g ice cube at 0 °C is put in an equal mass of 98 °C water in an insulated container, and allowed to melt. After a while the resulting mixture is at a final (uniform) temperature. How much heat energy does the ice require to completely melt at 0 °C? What is the uniform temperature of the resulting 'mixture'? Note: assume that the heat capacity of water and ice are independent of temperature, and are 4200 J/kgK and 2100 J/kgK, respectively. The heats of fusion and vapourization of water are 3.33x10° J/kg and 2.35×10° J/kg, respectively. How much heat energy does the ice require to completely melt at 0 °C? Answer = What is the resulting temperature of the total liquid (in °C)? Answer = "C

Chemistry: The Molecular Science

5th Edition

ISBN:9781285199047

Author:John W. Moore, Conrad L. Stanitski

Publisher:John W. Moore, Conrad L. Stanitski

Chapter9: Liquids, Solids, And Materials

Section: Chapter Questions

Problem 35QRT

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Calculate the quantity of heating required to convert the water in four ice cubes (60.1 g each) from H2O(s) at 0 °C to H2O(g) at 100. °C. The enthalpy of fusion of ice is 333 J/g and the enthalpy of vaporization of liquid water is 2260 J/g.
A quantity of ice at 0C is added to 64.3 g of water in a glass at 55C. After the ice melted, the temperature of the water in the glass was 15C. How much ice was added? The heat of fusion of water is 6.01 kJ/mol and the specific heat is 4.18 J/(g C).
The enthalpy of vaporization of water is larger than its enthalpy of fusion. Explain why.
Explain why the enthalpies of vaporization of the following substances increase in the order CH4NH3H2O, even though all three substances have approximately the same molar mass.
The phase diagram for water over a relative narrow pressure and temperature range is given in Figure 9.19. A phase diagram over a considerably wider range of temperature and pressure (kbar) is given nearby. This phase diagram illustrates the polymorphism of ice, the existence of a solid in more than one form. In this case, Roman numerals are used to designate each polymorphic form. For example, Ice I, ordinary ice, is the form that exists under ordinary pressures. The other forms exist only at higher pressures, in some cases extremely high pressure such as Ice VII and Ice VIII. Using the phase diagram, give the approximate P and T conditions at the triple point for Ice III, Ice V, and liquid water. Determine the approximate temperature and pressure for the triple point for Ices VI, VII, and VIII. What is anomalously different about the fusion curves for Ice VI and Ice VII compared to that of Ice I? What phases exist at 8 kbar and 20 °C? At a constant temperature of −10 °C, start at 3 kbar and increase the pressure to 7 kbar. Identify all the phase changes that occur sequentially as these conditions change. Explain why there is no triple point for the combination of Ice VII, Ice VIII, and liquid water.
Follow the step-wise process outlined in Problem 31 to calculate the amount of heat involved in condensing 100.00 g of benzene gas (C6H6) at 80.00C to liquid benzene at 25.00C. Use Tables 8.1 and 8.2 for the specific heat, boiling point, and heat of vaporization of benzene.
Is it possible to liquefy nitrogen at room temperature (about 25 C)? Is it possible to liquefy sulfur dioxide at room temperature? Explain your answers.
A special vessel (see Fig. 10.45) contains ice and supercooled water (both at 10C) connected by vapor space. Describe what happens to the amounts of ice and water as time passes.
Use Figure 11.7 to estimate the boiling point of carbon tetrachloride, CCl4, under an external pressure of 250 mmHg.
The cooling effect of alcohol on the skin is due to its evaporation. Calculate the heat of vaporization of ethanol (ethyl alcohol), C2H5OH. C2H5OH(l)C2H5OH(g);H=? The standard enthalpy of formation of C2H5OH(l) is 277.7 kJ/mol and that of C2H5OH(g) is 235.1 kJ/mol.
The molar heat of fusion of sodium metal is 2.60 kJ/mol, whereas its heat of vaporization is 97.0 kJ/mol. a. Why is the heat of vaporization so much larger than the heat of fusion? b. What quantity of heat would be needed to melt 1.00 g sodium at its normal melting point? c. What quantity of heat would be needed to vaporize 1.00 g sodium at its normal boiling point? d. What quantity of heat would be evolved if 1.00 g sodium vapor condensed at its normal boiling point?
When I mole of benzene is vaporized at a constant pressure of 1.00 atm and at its boiling point of 353.0 K, 30.79 kJ of energy (heat) is absorbed and the volume change is +28.90 L. What are E and H for this process?

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