Work & Heat
A 50.0 g block of aluminum at 25.0 °C is heated to 75.0 °C. Given that the specific heat of Al = 0.897 J/g·°C, calculate the heat absorbed by the block (in Joules).
2242.5 Joules
How much energy is required to melt 60.0 g of benzene (C6H6) at its melting point given that ΔHfus(C6H6)=10.59 kJ/mol? (in kJ)
q = 8.135 kJ
Calculate the standard entropy change (ΔS°) for the following reaction at 25 °C. Calculate ΔS° for the reaction. Additionally, explain how you would know if this were a spontaneous/nonspontaneous reaction conceptually (without calculations).
2SO2(g)+O2(g)→2SO3(g)
Given standard molar entropies (in J/mol·K):
S∘(SO2)=248.2
S∘(O2)=205.0
S∘(SO3)=256.8
ΔS = -187.8 J/molK
Conceptually: # of gas molecules in equation changes from 3 -> 2
A reaction has the following thermodynamic parameters at 298 K: ΔH° = –45.0 kJ/mol, ΔS° = –120 J/mol·K. Calculate ΔG° at 298 K. Is the reaction spontaneous at room temperature?
ΔG = -9.24 kJ/mol
YES, the reaction is spontaneous as ΔG is "-"
A chemist heats 60.0 mL of methanol from 20.0 °C to 65.0 °C to dissolve a compound. Given that the density of methanol = 0.792 g/mL, and the specific heat of methanol = 2.51 J/g·°C, calculate the total heat absorbed by the methanol (in Joules).
5367.384 Joules
How much energy is needed to vaporize 125.0 g of acetone (C2H3O) at its boiling point? Given that: ΔHvap(CH3COCH3)=29.1 kJ/mol. (in kJ)
q = 62.63 kJ
The formation of liquid water from its elements is given by:
H2(g)+ ½O2(g)→H2O(l) ΔH∘= –285.8 kJ/mol.
Calculate the entropy change of the surroundings (ΔSₛᵤᵣᵣ) at 298 K in J/molK. Is the entropy of the surroundings increasing or decreasing? Explain in terms of heat flow and spontaneity.
ΔS(surroundings) = 959 J/molK
Entropy of system is increasing (+ΔS)
For the reaction:
A(g)+B(l)→C(g)
The ΔH° = +35.0 kJ/mol and ΔS° = +150 J/mol·K. Calculate the temperature at which the reaction becomes spontaneous (specify above or below). Justify your answer using Gibbs Free Energy logic. Predict whether the reaction is spontaneous at 25°C.
ABOVE 233.33K
A 1.00 mol sample of an ideal gas is heated at constant pressure from 300 K to 600 K. Given: Cp = 29.1 J/mol·K, External pressure = 1.00 atm, Calculate the heat absorbed by the gas. Then, determine the change in internal energy ΔU if the gas performed -636.3 Joules of work.
q = 8730 Joules
𝚫U = 8093.7 Joules
Using the following reactions, calculate the enthalpy change (ΔH) for:
2C(s)+2H2(g)→C2H4(g)
C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(l) ΔH= –1411 kJ
C(s) + O2(g) → CO2(g) ΔH= –393.5 kJ
H2(g) + (½)O2(g) → H2O(l) ΔH= –285.8 kJ
52.4 kJ/mol
Consider the synthesis of ammonia:
N2(g)+3H2(g)→2NH3(g)
At 298 K:
ΔH∘= –92.4 kJ/mol
ΔS∘(system) = –198.3 J/mol K
Calculate ΔSₛᵤᵣᵣ and ΔSₜₒₜ at 298 K. Is the reaction spontaneous at 298 K?
ΔS(surroundings) = 310.07 J/molK
ΔS(total) = 111.77 J/molK
Yes, the reaction is spontaneous as ΔS is positive.
The melting of benzene at its melting point (5.5°C) involves the following thermodynamic data: ΔH(freezing) = - 10.59 kJ/mol, ΔS(freezing) = - 37.9 J/mol·K. Calculate ΔG° for melting benzene at 5.5°C. Is melting spontaneous at this temperature?
ΔG = 0.02917 kJ/mol
NOT spontaneous, ΔG is "+"
A 100.0 g sample of water at 25.0 °C is heated to 100.0 °C and then fully vaporized at 1 atm. However, due to imperfect insulation, 5.00% of the total heat supplied is lost to the surroundings. Given: Cp(liquid)=4.184 J/g°C, ΔHvap=40.7 kJ/mol, Assume ideal gas behavior and use R=8.314 J/mol K. Vaporization occurs at 100 °C (373 K). Calculate the total heat required to raise the temperature of the water to 100 °C (in J). Then, calculate the heat required to fully vaporize the water (in kJ).
Total heat to raise temp: 31,380 J
Heat to vaporize: 225.86 kJ
Calculate the enthalpy change for the combustion of ethanol:
C2H5OH(l)+3O2(g)→2CO2(g)+3H2O(l)
Given:
ΔHf∘(C2H5OH)= –277.0 kJ/mol
ΔHf∘(CO2)= –393.5 kJ/mol
ΔHf∘(H2O(l))= –285.8 kJ/mol
Rank the following processes in order of entropy change (ΔS) (1 being the highest in entropy change), and justify your ranking:
Sublimation of dry ice
Condensation of steam
Dissolution of NaNO₃(s) in water
Decomposition of H₂O₂(l) → H₂O(l) + O₂(g)
#1 Sublimation of dry ice
#2 Decomposition of H2O2
#3 Dissolution of NaNO3(s) in water
#4 Condensation of steam
Ice melts at 0°C (273.15 K) under standard pressure. The process is:
H2O(s)→H2O(l)
ΔH°fus = +6.01 kJ/mol
ΔS°fus = +22.0 J/mol·K
Calculate ΔG° for melting ice at 0°C. Is the process spontaneous, nonspontaneous, or at equilibrium at 0°C? Then, predict if the system would become more spontaneous at –10°C.
ΔG = 0.0007 -> basically zero, so in equilibrium
Not spontaneous at -10.0. ΔG = 0.2207 kJ/mol
A 150.0 g sample of solid methanol (CH₃OH) at –10.0 °C is heated at 1 atm until 80.0 g of the substance has vaporized. Assume no heat loss to the surroundings. Given that the melting point of methanol is -97.8°C, boiling point is 64.7°C, Cp(solid) =2.09 J/g°C, Cp(liquid) = 2.510 J/g°C, ΔHfus=3.18 kJ/mol, ΔHvap=35.2 kJ/mol, determine the total heat absorbed during the entire process (in kJ).
q total = 136.439 kJ
Calculate the enthalpy change (ΔH) for the following gas-phase reaction using average bond enthalpies:
CH4(g)+Cl2(g)→CH3Cl(g)+HCl(g)
C–H = 412
Cl–Cl = 243
C–Cl = 338
H–Cl = 431
114 kJ/mol
The thermal decomposition of calcium carbonate is given by:
CaCO3(s)→CaO(s)+CO2(g)
At 25 °C:
ΔH∘= +178.3 kJ
S∘(CaCO3) = 92.2 J/molK
S∘(CaO) = 39.8 J/molK
S∘(CO2) = 213.8 J/molK
Calculate the total entropy change (ΔSₜₒₜ) in J/molK. Is the reaction spontaneous at 298 K? Justify using entropy logic.
ΔS= -436.922
No, reaction is not spontaneous at this temperature as ΔS is negative.
A disposable heat pack contains magnesium powder that reacts with steam to produce magnesium oxide and hydrogen gas:
Mg(s)+H2O(g)→MgO(s)+H2(g)
This reaction is exothermic and used to warm food in field kits.
Standard enthalpies of formation:
ΔH°f(Mg) = 0 kJ/mol
ΔH°f(H₂O, gas) = –241.8 kJ/mol
ΔH°f(MgO) = –601.6 kJ/mol
What is the change in Gibbs free energy given that the reaction occurs at 341K and has a ΔS(system) of 123 J/molK?
ΔG = -401.743 kJ/mol