895/variants/431895/take/9/ TEXT ANSWER Question 10 A calorimeter holds 50 g water at 22.0°C. A sample of hot iron is added to the water. The final temperature of the water and iron is 28.0°C. What is the change in enthalpy associated with the change in the water's temperature? Note: The specific heat of water is 4.18 g°C Use the formula AH = - cm AT Show your work. Normal A A 3x S I = = = = 外 Iロび Enter your answer here Oo00 OO of 10 Total Questions Answered All Changes Saved Continue > 11:10 AM Vordsworth, Shelle.. OEdio | Course Stude. 书 5/5/2022

895/variants/431895/take/9/ TEXT ANSWER Question 10 A calorimeter holds 50 g water at 22.0°C. A sample of hot iron is added to the water. The final temperature of the water and iron is 28.0°C. What is the change in enthalpy associated with the change in the water's temperature? Note: The specific heat of water is 4.18 g°C Use the formula AH = - cm AT Show your work. Normal A A 3x S I = = = = 外 Iロび Enter your answer here Oo00 OO of 10 Total Questions Answered All Changes Saved Continue > 11:10 AM Vordsworth, Shelle.. OEdio | Course Stude. 书 5/5/2022

Principles of Modern Chemistry

8th Edition

ISBN:9781305079113

Author:David W. Oxtoby, H. Pat Gillis, Laurie J. Butler

Publisher:David W. Oxtoby, H. Pat Gillis, Laurie J. Butler

Chapter12: Thermodynamic Processes And Thermochemistry

Section: Chapter Questions

Problem 14P

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In a coffee-cup calorimeter, 100.0 mL of 1.0 M NaOH and 100.0 mL of 1.0 M HCI are mixed. Both solutions were originally at 24.6C. After the reaction, the final temperature is 31.3C. Assuming that all the solutions have a density of 1.0 g/cm3 and a specific heat capacity of 4.18 J/Cg, calculate the enthalpy change for the neutralization of HCl by NaOH. Assume that no heat is lost to the surroundings or to the calorimeter.
Ethyl alcohol, C2H5OH, is the intoxicating agent in liquor. Burning 1.00 g of ethyl alcohol in an excess of oxygen at 23.28C and constant volume releases 29.52 kJ of heat. When 7.40 g of ethyl alcohol is burned in oxygen under the same conditions in a bomb calorimeter, the temperature of the bomb and water rises from 23.28C to 48.04C. The bomb holds 0.750 kg of water. (a) What is q for the combustion of the ethyl alcohol in the bomb calorimeter? (b) What is qH2O? (c) What is qcal? (d) What is the heat capacity of the calorimeter?
A 21.3-mL sample of 0.977 M NaOH is mixed with 29.5 mL of 0.918 M HCl in a coffee-cup calorimeter (see Section 6.6 of your text for a description of a coffee-cup calorimeter). The enthalpy of the reaction, written with the lowest whole-number coefficients, is 55.8 kJ. Both solutions are at 19.6C prior to mixing and reacting. What is the final temperature of the reaction mixture? When solving this problem, assume that no heat is lost from the calorimeter to the surroundings, the density of all solutions is 1.00 g/mL, the specific heat of all solutions is the same as that of water, and volumes are additive.
A sample of sucrose, C12H22O11, is contaminated by sodium chloride. When the contaminated sample is burned in a bomb calorimeter, sodium chloride does not burn. What is the percentage of sucrose in the sample if a temperature increase of 1.67C is observed when 3.000 g of the sample are burned in the calorimeter? Sucrose gives off 5.64103kJ/mol when burned. The heat capacity of the calorimeter and water is 22.51 kJ/C.
A bomb calorimetric experiment was run to determine the enthalpy of combustion of ethanol. The reaction is C2H5OH() + 3 O2(g) 2 CO2(g) + 3 H2O() The bomb had a heat capacity of 550 J/K, and the calorimeter contained 650 g of water. Burning 4.20 g of ethanol, C2HsOH() resulted in a rise in temperature from 18.5 C to 22.3 C. Calculate U for the combustion of ethanol, in kJ/mosl.
In the 1880s, Frederick Trouton noted that the enthalpy of vaporization of 1 mol pure liquid is approximately 88 times the boiling point, Tb, of the liquid on the Kelvin scale. This relationship is called Troutons rule and is represented by the thermochemical equation liquid gas H = 88 Tb, joules Combined with an empirical formula from chemical analysis, Troutons rule can be used to find the molecular formula of a compound, as illustrated here. A compound that contains only carbon and hydrogen is 85.6% C and 14.4% H. Its enthalpy of vaporization is 389 J/g, and it boils at a temperature of 322 K. (a) What is the empirical formula of this compound? (b) Use Troutons rule to calculate the approximate enthalpy or vaporization or one mole of the compound. Combine the enthalpy of vaporization per mole with that same quantity per gram to obtain an approximate molar mass of the compound. (c) Use the results of parts (a) and (b) to find the molecular formula of this compound. Remember that the molecular mass must be exactly a whole-number multiple of the empirical formula mass, so considerable rounding may be needed.
Peanuts and peanut oil are organic materials and bum in air. How many burning peanuts does it take to provide the energy to boil a cup of water (250 mL of water)? To solve this problem, we assume each peanut, with an average mass of 0.73 g, is 49% peanut oil and 21% starch; the remainder is noncombustible We further assume peanut oil is palmitic acid, C16H32O2, with an enthalpy of formation of 848.4 kJ/mol. Starch is a long chain of C6H10O5 units, each unit having an enthalpy of formation of 960 kJ.
9.102 A runner generates 418 kJ of energy per kilometer from the cellular oxidation of food. The runner's body must dissipate this heat or the body will overheat. Suppose that sweat evaporation is the only important cooling mechanism. If you estimate the enthalpy of evaporation of water as 44 kJ/mol and assume that sweat can he treated as water, describe how you would estimate the volume of sweat that would have to be evaporated if the runner runs a 10-km race.
You should use care when dissolving H2SO4 in water because the process is highly exothermic. To measure the enthalpy change, 5.2 g of concentrated H2SO4() was added (with stirring) to 135 g of water in a coffee-cup calorimeter. This resulted in an increase in temperature from 20.2 C to 28.8 C. Calculate the enthalpy change for the process H2SO4() H2SO4(aq), in kf/mol.
Magnesium sulfate is often used in first-aid hot packs, giving off heat when dissolved in water. A coffee-cup calorimeter at 25C contains 15.0 mL of water at 25C. A 2.00-g sample of MgSO4 is dissolved in the water and 1.51 kJ of heat are evolved. (You can make the following assumptions about the solution: volume=15.0 mL, density=1.00 g/mL, specific heat=4.18J/gC.) (a) Write a balanced equation for the solution process. (b) Is the process exothermic? (c) What is qH2O? (d) What is the final temperature of the solution? (e) What are the initial and final temperatures in F?
A 29.1-mL sample of 1.05 M KOH is mixed with 20.9 mL of 1.07 M HBr in a coffee-cup calorimeter (see Section 6.6 of your text for a description of a coffee-cup calorimeter). The enthalpy of the reaction, written with the lowest whole-number coefficients, is 55.8 kJ. Both solutions are at 21.8C prior to mixing and reacting. What is the final temperature of the reaction mixture? When solving this problem, assume that no heat is lost from the calorimeter to the surroundings, the density of all solutions is 1.00 g/mL, and volumes are additive.
When one mol of KOH is neutralized by sulfuric acid, q=56 kJ. (This is called the heat of neutralization.) At 23.7C, 25.0 mL of 0.475 M H2SO4 is neutralized by 0.613 M KOH in a coffee-cup calorimeter. Assume that the specific heat of all solutions is 4.18J/gC, that the density of all solutions is 1.00 g/mL, and that volumes are additive. (a) How many mL of KOH is required to neutralize H2SO4? (b) What is the final temperature of the solution?