Answer:
Explanation:
1) Reaction (given):
ΔS = -17.4 J/K
2) Spontaneous reactions (ΔG < 0)
The sign of the change in free energy (ΔG) tells if a reaction is spontaneous or not at a given temperature and constant pressure.
If ΔG is negative ( ΔG < 0) the reaction is spontaneous; if ΔG is positive ( ΔG > 0) the reaction is nonspontaneous.
Then, determine the temperature at which ΔG becomes zero, which is the temperature at which the reaction becomes nonspontaneous.
3) ΔG⁰ = ΔH⁰ - TΔS⁰
ΔG⁰ = -11,000 J - T (-17.4 J/ K) [ -11.0 kJ were converted to J]
0 = - 11,000 J + 17.4 T J/K
11,000 J = 17.4 T J/ k
T = 11,000 kJ / (17.4 J/K) = 632 K ← answer
The following reaction become nonspontaneous at 632.18 K
Gibbs free energy is the maximum possible work given by chemical reactions at constant pressure and temperature. Gibbs free energy can be used to determine the spontaneity of a reaction
If the Gibbs free energy value is <0(negative) then the chemical reaction occurs spontaneously. If the change in free energy is zero, then the chemical reaction is at equilibrium ,if it is >0, the process is not spontaneous
Free energy of reaction (G) is the sum of its enthalpy (H) plus the product of the temperature and the entropy (S) of the system
Can be formulated: (at any temperature)
[tex]\large{\boxed{\bold{\Delta G=\Delta H-T.\Delta S}}}[/tex]
or at (25 Celsius / 298 K, 1 atm= standart)
ΔG°reaction = ΔG°f (products) - ΔG°f (reactants)
Under standard conditions:
∆G° = ∆H° - T∆S°
The value of ∆H° can be calculated from the change in enthalpy of standard formation:
∆H° (reaction) = ∑H° (product) - ∑ H° (reagent)
The value of ΔS° can be calculated from standard entropy data
∆S° (reaction) = ∑S° (product) - ∑ S°(reagent)
Reaction :
FeO(s) + CO(g) → CO2(g) + Fe(s)
ΔH = -11.0 kJ= -11000 J; ΔS = -17.4 J/K
ΔG =-11000 -(T.-17.4 J/K)
0 = -11000 + T17.4
T17.4 = 11000
T =632.18
Delta H solution
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an exothermic reaction
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as endothermic or exothermic
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an exothermic dissolving process
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Keywords: the standard gibbs free energy of formation,nonspontaneous
