thermodymanics. thermodynamics is a branch of science that focuses on energy changes that accompany...
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Thermodymanics
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Thermodynamics is a branch of science that focuses on energy changes that accompany chemical and physical changes.
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Objective: To calculate heat capacity.
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Heat (q): The energy transferred between objects that are at different temperatures.
Unit: joules (J)
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Molar Heat Capacity: Energy (heat) needed to increase the temperature of 1 mol of substance by 1 K.
q = nC∆T
q = heatn = # of molesC = molar heat capacity∆T = change in temperature
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The molar heat capacity of water is larger than the molar heat capacity of land. This means that water does not heat up as easily as land does. As a result, oceans can help keep coastal areas cool during the summer.
The filling of a fruit pie has a larger heat capacity than the crust. This means that fruit filling will retain heat better and the crust will cool much quicker. As a result, eating the fruit filling can cause burns (even though it may appear that the pie is cool).
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Determine the energy (heat) needed to increase the temperature of 10.0 mol of Hg by 7.5 K. The value of C for Hg is 27.8 J/K۰mol.
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Determine the energy (heat) needed to increase the temperature of 10.0 mol of Hg by 7.5 K. The value of C for Hg is 27.8 J/K۰mol.
q = ?n = 10.0 molC = 27.8 J/K۰mol∆T = 7.5 K
q = nC∆Tq = (10.0 mol)(27.8 J/K۰mol)(7.5 K)
q= 2.1 x 103 J
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The molar heat capacity of tungsten is 24.2 J/K ۰mol. Calculate the energy as heat needed to increase the temperature of 0.404 mol of W by 10.0 K.
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The molar heat capacity of tungsten is 24.2 J/K ۰mol. Calculate the energy as heat needed to increase the temperature of 0.404 mol of W by 10.0 K.
q = ?n = 0.404 molC = 24.2 J/K۰mol∆T = 10.0 K
q = nC∆Tq = (0.404 mol)(24.2 J/K۰mol)(10.0 K)
q= 97.8J
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Suppose a sample of NaCl increased in temperature by 2.5 K when the sample absorbed 1.7 x 102 J energy (heat). Calculate the number of moles of NaCl if the molar heat capacity is 50.5 J/K۰mol.
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Suppose a sample of NaCl increased in temperature by 2.5 K when the sample absorbed 1.7 x 102 J energy (heat). Calculate the number of moles of NaCl if the molar heat capacity is 50.5 J/K۰mol.
q = 1.7 x 102 J n = ?C = 50.5 J/K۰mol ∆T = 2.5 K
q = nC∆T1.7 x 102 J = n(50.5 J/K۰mol)(2.5 K)
n= 1.3 mol
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Calculate the energy as heat needed to increase the temperature of 0.80 mol of nitrogen, N2, by 9.5 K. The molar heat capacity of nitrogen is 29.1 J/K۰mol.
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Calculate the energy as heat needed to increase the temperature of 0.80 mol of nitrogen, N2, by 9.5 K. The molar heat capacity of nitrogen is 29.1 J/K۰mol.
q = ?n = 0.80 molC = 29.1 J/K۰mol ∆T = 9.5 K
q = nC∆Tq = (0.80 mol)(29.1 J/K۰mol)(9.5 K)
q= 2.2 x 102 J
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A 0.07 mol sample of octane, C8H18, absorbed 3.5 x 103 J of energy. Calculate the temperature increase of octane if the molar heat capacity of octane is 254.0 J/K۰mol.
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A 0.07 mol sample of octane, C8H18, absorbed 3.5 x 103 J of energy. Calculate the temperature increase of octane if the molar heat capacity of octane is 254.0 J/K۰mol.
q = 3.5 x 103 J n = 0.07 molC = 254.0 J/K۰mol ∆T = ?
q = nC∆T3.5 x 103 J = (0.07 mol)(254.0 J/K۰mol) ∆T
∆T = 2.0 x 102 K
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Objective: To calculate the change in enthalpy.
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The total energy of a system is impossible to measure.
However, we can measure the change in enthalpy of a system.
Enthalpy: The total energy content of a sample.
H = Enthalpy ∆H = Change in Enthalpy
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∆H can be measured with a calorimeter.
A calorimeter is used to measure the heat absorbed or released in a chemical or physical change.
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Enthalpy changes can be used to determine if a process is endothermic or exothermic.
Exothermic Reaction: Negative Enthalpy Change
Endothermic Reaction: Positive Enthalpy Changes
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The standard enthalpy of formation ( ) is the enthalpy change in forming 1 mol of a substance from elements in their standard states.
Note: The value for the standard enthalpy of formation for an element is 0.
The values for the standard enthalpies of formation can be found using a table.
0fH
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ΔHreaction = ΔH products - ΔHreactants
Step 1: Determine for each compound using enthalpy table.
Step 2: Multiply by the coefficients from the balanced equation (# of moles).
Step 3: Set up ΔH equation. Step 4: calculate.
0fH
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ΔHreaction = ΔH products - ΔHreactants
Calculate ΔH for the following reaction and determine if the reaction is exothermic or endothermic.
SO2(g) + NO2(g) SO3(g) + NO(g)
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ΔHreaction = ΔH products - ΔHreactants
SO2(g) + NO2(g) SO3(g) + NO(g)
-296.8(1) 33.1(1) -395.8(1) 90.3(1)
kJ/molkJ/mol kJ/molkJ/mol
ΔH= [(-395.8 kJ/mol)(1) + (90.3 kJ/mol)(1)] – [(296.8
kJ/mol)(1) + (33.1 kJ/mol)(1)]
Δ H = -305.5 kJ – 329.9 kJ
Δ H = -635.4 kJ The reaction is exothermic.
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Calculate the enthalpy change for the following reaction:
2C2H6(g) + 7O2(g) 4 CO2(g) + 6H2O(g)
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Calculate ΔH for the following reaction:CaO(s) + H2O(l) Ca(OH)2(s)
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Calculate the enthalpy change for the combustion of methane gas.
CH4(g) + 2O2(g) CO2(g) + 2H2O(l)
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Objective: To calculate the change in entropy.
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A reaction is more likely to occur if enthalpy (ΔH) is negative.
However, some endothermic reactions can occur easily. Why? Entropy!
Entropy (ΔS): A measure of the randomness or disorder of a system.
A process if more likely to occur if there is an increase in entropy (or if ΔS is positive).
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Entropy is increased by the following factors:Diffusion (process of dispersion)Dilution of a solutionDecreasing the pressure of a gas Increasing temperatureThe number of moles of product is greater
than the number of moles of reactant Increasing the total number of particles in a
systemWhen a reaction produces more gas
particles (opposed to liquid or solids)
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ΔSreaction = ΔSproducts - ΔSreactants
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Find the change in entropy for the following reaction:
2Na (s) + 2HCl (g) 2NaCl (s) + H2(g)
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Find the change in entropy for the following reaction:
2Na (s) + 2H2O (l) 2NaOH (s) + H2(g)
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Objectives:(1)To calculate the change in Gibbs
energy.(2)To determine if a reaction is
spontaneous or nonspontaneous.
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Gibbs Energy: the energy in a system that is available to do useful work.
Gibbs Energy is also called Free Energy.
ΔG = Change in Gibbs Energy
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ΔGreaction = ΔGproducts - ΔGreactants
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Calculate ΔG for the following water-gas reaction:
C(s) + H2O (g) CO (g) + H2 (g)
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Calculate the Gibbs energy change that accompanies the following reaction:
C(s) + O2 (g) CO2 (g)
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Calculate the Gibbs energy change that accompanies the following reaction:
CaCO3 (s) CaO (s) + CO2 (g)
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If ΔG is negative, the forward reaction is spontaneous.
If ΔG is 0, the system is at equilibrium.
If ΔG is positive, the forward reaction is nonspontaneous. NOTE: In this case, the reaction is spontaneous in the
reverse direction.
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ΔG = ΔH - TΔS
Step 1: Organize the information.Step 2: Change the units.Step 3: Step up ΔG equation.Step 4: Calculate.
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Given the changes in enthalpy and entropy are -139 kJ and 277 J/K respectively for a reaction at 25⁰C, calculate the change in Gibbs energy.
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Given the changes in enthalpy and entropy are -139 kJ and 277 J/K respectively for a reaction at 25⁰C, calculate the change in Gibbs energy.
ΔH = -139 kJΔS = 277 J/K / (1000 J/kJ) = 0.277 kJ/KT = 25 ⁰C + 273 = 298 KΔG = ?
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Given the changes in enthalpy and entropy are -139 kJ and 277 J/K respectively for a reaction at 25⁰C, calculate the change in Gibbs energy.
ΔH = -139 kJΔS = 277 J/K / (1000 J/kJ) = 0.277 kJ/KT = 25 ⁰C + 273 = 298 KΔG = ?
ΔG = ΔH - T ΔSΔG = (-139 kJ) – (298K)(0.277 kJ/K)ΔG = (-139 kJ) – (82.546 kJ)ΔG = -221.55 kJ The reaction is spontaneous.
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A reaction has a ΔH of -76 kJ and a ΔS of -117 J/K. Calculate the ΔG at 298 K. Is the reaction spontaneous?
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A reaction has a ΔH of 11 kJ and a ΔS of 49 J/K. Calculate ΔG at 298 K. Is the reaction spontaneous?
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Objective: To calculate the melting and boiling point of a substance.
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Tmp: melting point temperature Tbp: boiling point temperature ΔHfus: molar enthalpy of fusion ΔSfus: molar entropy of fusion ΔHvap: molar enthalpy of vaporization ΔSvap: molar entropy of vaporization
fus
fusmp S
HT
vap
vapbp S
HT
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Step 1: Organize the information. Step 2: Change the units. Step 3: Step up the equation. Step 4: Calculate.
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The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol·K. Calculate the melting point and boiling point of Hg.
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The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol·K. Calculate the melting point and boiling point of Hg.
Tmp = ?
ΔHfus = 2.295 kJ/mol
ΔSfus = 9.79 J/mol·K = 0.00979 kJ/mol·K
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The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol·K. Calculate the melting point and boiling point of Hg.
Tmp = ?
ΔHfus = 2.295 kJ/mol
ΔSfus = 9.79 J/mol·K = 0.00979 kJ/mol·K
Tmp = (2.295 kJ/mol) / (0.00979 kJ/mol·K)
Tmp = 234 K
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The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol·K. Calculate the melting point and boiling point of Hg.
Tmp = ?
ΔHfus = 2.295 kJ/mol
ΔSfus = 9.79 J/mol·K = 0.00979 kJ/mol·K
Tmp = (2.295 kJ/mol) / (0.00979 kJ/mol·K)
Tmp = 234 K
Tbp = ?
ΔHvap = 59.2 kJ/mol
ΔSvap = 93.8 J/mol·K = 0.0938 kJ/mol·K
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The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol·K. Calculate the melting point and boiling point of Hg.
Tmp = ?
ΔHfus = 2.295 kJ/mol
ΔSfus = 9.79 J/mol·K = 0.00979 kJ/mol·K
Tmp = (2.295 kJ/mol) / (0.00979 kJ/mol·K)
Tmp = 234 K
Tbp = ?
ΔHvap = 59.2 kJ/mol
ΔSvap = 93.8 J/mol·K = 0.0938 kJ/mol·K
Tbp = (59.2 kJ/mol) / (0.0938 kJ/mol·K )
Tbp = 631 K
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For ethanol, the molar enthalpy of fusion is 4.931 kJ/mol and the molar entropy of fusion is 31.6 J/mol·K. The molar enthalpy of vaporization at the boiling point is 42.32 kJ/mol and the molar entropy of vaporization is 109.9 J/mol·K. Calculate the melting and boiling points for ethanol.
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For sulfur dioxide, the molar enthalpy of fusion is 8.62 kJ/mol and the molar entropy of fusion is 43.1 J/mol·K. The enthalpy of vaporization at the boiling point is 24.9 kJ/mol and the molar entropy of vaporization is 94.5 J/mol·K. Calculate the melting and boiling points for sulfur dioxide.
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For ammonia, ΔHfus is 5.66 kJ/mol and Δsfus is 29.0 J/mol·K. ΔHvap is 23.33 kJ/mol and ΔSvap is 97.2 J/mol·K. Calculate the melting and boiling points for ammonia.