4) The following data are for mercury II chloride (HgCl2 M = 271.50 g/mol), and are taken from Chase, M. W. The NIST-JANAF Thermochemical Tables, 4th Edition, 1998. AH°fus Tofus = 550. K 19.41 kJ/mol AH°vap 58.9 kJ/mol T°vap = 577. K Using the above data, give the phase diagram for HgCl2. Include temperatures in the range 500 K-600 K, and pressures in the range 0.0 atm - 2.0 atm. Include in your phase diagram the correct location (pressure and temperature) of the triple point. (Hint - For the solid- gas and liquid-gas phase boundaries use the Clausius-Clapeyron equation to find values of p and T at several different temperatures).

4) The following data are for mercury II chloride (HgCl2 M = 271.50 g/mol), and are taken from Chase, M. W. The NIST-JANAF Thermochemical Tables, 4th Edition, 1998. AH°fus Tofus = 550. K 19.41 kJ/mol AH°vap 58.9 kJ/mol T°vap = 577. K Using the above data, give the phase diagram for HgCl2. Include temperatures in the range 500 K-600 K, and pressures in the range 0.0 atm - 2.0 atm. Include in your phase diagram the correct location (pressure and temperature) of the triple point. (Hint - For the solid- gas and liquid-gas phase boundaries use the Clausius-Clapeyron equation to find values of p and T at several different temperatures).

Chemistry for Engineering Students

4th Edition

ISBN:9781337398909

Author:Lawrence S. Brown, Tom Holme

Publisher:Lawrence S. Brown, Tom Holme

Chapter9: Energy And Chemistry

Section: Chapter Questions

Problem 9.46PAE: 9.46 The heat of fusion of pure silicon is 43.4 kJ/mol. How much energy would be needed to melt a...

See similar textbooks

Similar questions

9.46 The heat of fusion of pure silicon is 43.4 kJ/mol. How much energy would be needed to melt a 5.24-g sample of silicon at its melting point of 1693 K?
A student is asked to calculate the amount of heat involved in changing 10.0 g of liquid bromine at room temperature (22.5C) to vapor at 59.0C.To do this, one must use Tables 8.1 and 8.2 for information on the specific heat, boiling point, and heat of vaporization of bromine. In addition, the following step-wise process must be followed. (a) Calculate H for: Br2(l,22.5C)Br2(l,59.0C) (b) Calculate H for: Br2(l,59.0C)Br2(g,59.0C) (c) Using Hess's law, calculate H for: Br2(l,22.5C)Br2(g,59.0C)
9.95 How much heat is required to convert 250 g of water from liquid at 23°C: into steam at 107°C? (The necessary heat capacities and enthalpy changes can be found in the chapter or online.)
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 enthalpy of combustion of hard coal averages 35 kJ/g, that of gasoline, 1.28105 kJ/gal. How many kilograms of hard coal provide the same amount of heat as is available from 1.0 gallon of gasoline? Assume that the density of gasoline is 0.692 g/mL (the same as the density of isooctane).
The enthalpy of vaporization of CO2(l) is 9.8 kJ/mol. Would you expect the enthalpy of vaporization of CS2(l) to be 28 kJ/mol, 9.8 kJ/mol, or 8.4 kJ/mol? Discuss the plausibility of each of these answers.
Draw a diagram like Figure 2.11 that illustrates the change in enthalpy for the chemical reaction C s 2H2 g CH4 g Which is exothermic by 74.8 kJ/mol.
It has been determined that the body can generate 5500 kJ of energy during one hour of strenuous exercise. Perspiration is the bodys mechanism for eliminating this heat. What mass of water would have to be evaporated through perspiration to rid the body of the heat generated during 2 hours of exercise? (The heat of vaporization of water is 40.6 kJ/mol.)
If you want to convert 56.0 g ice (at 0 °C) to water at 75.0 °C, calculate how many grams of propane, C3H8, you would have to bum to supply the energy to melt the ice and then warm it to the final temperature (at 1 bar).
Consider the data for substance X given in Exercise 117. When the temperature of 1.000 mole of X(g) is lowered from 100.0C to form X(l) at 50.0C. 28.75 kJ of heat is released. Calculate the specific heat capacity of X(g).
A 0.250-g chunk of sodium metal is cautiously dropped into a mixture of 50.0 g water and 50.0 g ice, both at 0C. The reaction is 2Na(s)+2H2O(l)2NaOH(aq)+H2(g)H=368kJ Assuming no heat loss to the surroundings, will the ice melt? Assuming the final mixture has a specific heat capacity of 4.18 J/gc, calculate the final temperature. The enthalpy of fusion for ice is 6.02 kJ/mol.
Phosphorus exists in multiple solid phases, including two known as red phosphorus and white phosphorus. (a) Based on their respective heats of formation, which form of phosphorus is defined as the standard state? (b) Now consider the phase transition between white and red phosphorous: P4(s, white4 P(s, red). Use data from Appendix E to determine which form of phosphorous is actually more stable at 25C. (Your result should reveal that phosphorus is an exception to the usual convention for defining the standard state.) (c) Is the same form of the solid more stable at all temperatures? If not, what temperatures are needed to make the other form more stable?