A geochemist in the field takes a small sample of the crystals of mineral compound X from a rock pool lined with more crystals of X. He notes the temperature of the pool, 30.° C, and caps the sample carefully. Back in the lab, the geochemist dissolves the crystals in 750. mL of distilled water. He then filters this solution and evaporates all the water under vacuum. Crystals of X are left behind. The researcher washes, dries and weighs the crystals. They weigh 0.098 kg. O yes Using only the information above, can you calculate the solubility of X in water at 30.° C ? no O-0 If you said yes, calculate it. Be sure your answer has a unit symbol and 2 significant digits.

A geochemist in the field takes a small sample of the crystals of mineral compound X from a rock pool lined with more crystals of X. He notes the temperature of the pool, 30.° C, and caps the sample carefully. Back in the lab, the geochemist dissolves the crystals in 750. mL of distilled water. He then filters this solution and evaporates all the water under vacuum. Crystals of X are left behind. The researcher washes, dries and weighs the crystals. They weigh 0.098 kg. O yes Using only the information above, can you calculate the solubility of X in water at 30.° C ? no O-0 If you said yes, calculate it. Be sure your answer has a unit symbol and 2 significant digits.

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter6: Thermochemistry

Section: Chapter Questions

Problem 68E: In a coffee-cup calorimeter, 1.60 g NH4NO3 is mixed with 75.0 g water at an initial temperature of...

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In a coffee-cup calorimeter, 1.60 g NH4NO3 is mixed with 75.0 g water at an initial temperature of 25.00C. After dissolution of the salt, the final temperature of the calorimeter contents is 23.34C. Assuming the solution has a heat capacity of 4.18 J/C g and assuming no heat loss to the calorimeter, calculate the enthalpy change for the dissolution of NH4NO3 in units of kJ/mol.
From data in Appendix J, estimate (a) the boiling point of titanium(IV) chloride. (b) the boiling point of carbon disulfide, CS2, which is a liquid at 25 C and 1 bar.
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?
As O2(l) is cooled at 1 atm, it freezes at 54.5 K to form solid I. At a lower temperature, solid I rearranges to solid II, which has a different crystal structure. Thermal measurements show that H for the I II phase transition is 743.1 J/mol, and S for the same transition is 17.0 J/K mol. At what temperature are solids 1 and II in equilibrium?
A 5.00 g-sample of KOH at 25.0 C was added to 100.0 g of H2O(l) at room temperature inside an insulated cup calorimeter, and the contents were stirred. After all the KOH(s) dissolved, the temperature of the solution had increased. Based on the information given, which of the following best justifies the claim that the dissolution of KOH is a thermodynamically favorable process? A. The forces between the ions and the water molecules are stronger than the forces between water molecules, thus, delta H < 0. Also, the ions become less dispersed as KOH(s) dissolves, thus delta S > 0. Therefore, delta G < 0. B. The energy require to break the bonds between the ions in the solid is less than that released as the ion-dipole attractions form during solvation, thus delta H < 0. Also, the ions become widely dispersed as KOH(s) dissolved, thus delta S > 0. Therefore, delta G < 0. C. The average kinetic energy of the particles increases, resulting in delta H > 0. Also, the…
The enthalpy of fusion of aluminum is 10.7 kj/mole. How many grams of aluminum can be melted by adding 81.4 KJ of energy to the metal at its melting point?
During the industrial synthesis of methanol (CH3OH), carbon monoxide (CO) reacts with hydrogen gas. For methanol, ΔH= -100.4 kJ/mol. a. Write the balanced chemical equation for this reaction. b. Is it preferable to maintain this system at high or low pressure? Explain.
If 94.0 g of ice (ΔHfus = 6.02 kJ/mol, at exactly 0°C) is added to a 240.-mL glass of room temperature (25°C) ethanol (c = 2.46 J/g°C. MW = 47.07 g/mol, density=0.789 g/mL) what will be the equilibrium temperature of the solution? Will it contain any solid ice? If so, what mass of ice remains?
The normal boiling point of naphthalene, C10H8, is 218 °C. What is the its enthalpy of vaporization using Trouton’s rule?
Unit 3: The normal boiling point of methanol is 64.7 °C and the molar enthalpy of vaporization is 71.8 kJ/mol. Assuming that the vaporization is an equilibrium process, what is the value of ΔS (in J mol-1 K-1) when 1.00 mol of CH3OH (l) vaporizes at 64.7 °C ?
) The salt potassium perchlorate is soluble in water. When 2.84 g of KClO4 is dissolved in 106.00 g of water, the temperature of the solution decreases from 25.00 to 22.54 °C. Based on this observation, calculate the enthalpy of dissolution of KClO4 (in kJ/mol).Assume that the specific heat of the solution is 4.184 J/g °C and that the heat absorbed by the calorimeter is negligible.
The salt cesium sulfate is soluble in water. When 22.7 g of Cs2SO4 is dissolved in 120.00 g of water, the temperature of the solution decreases from 25.00 to 22.96 °C. Based on this observation, calculate the enthalpy of dissolution of Cs2SO4 (in kJ/mol). Assume that the specific heat of the solution is 4.184 J/g °C and that the heat absorbed by the calorimeter is negligible.