I would really like to understand how to do these problems, because my exam is coming up in a few weeks. I still cannot comprehend how you can tell if a system is is positive or negative according to reaction.
For #1, I presume it is E if it is spontaneous at all temperatures. For #2, I think ΔS have to be negative since the reaction contains all gas phases. But I have no idea if you can tell if ΔH is spontaneous or nonspontaneous..And for #3, I presume it is A(?) basing on Gibbs Free Energy.
1. The following reaction is spontaneous at all temperatures:
CaC2(s) + 2H2O(l) ---> Ca(OH)2(s) + C2H2(g)
Which of the following statements is true?
A) ΔG is positive at all temperatures.
B) ΔG is positive and ΔS is positive.
C) ΔH is negative and ΔS is negative.
D) ΔH is negative and ΔS is positive.
E) ΔH is positive and ΔS is negative.
2. Consider the following reaction:
C3H8(g) + 5O2(g) ---> 3CO2(g) + 4H2O(g)
One would predict that
A) ΔH is positive and ΔS is positive for the reaction.
B) ΔH is negative and ΔS is negative for the reaction.
C) ΔH is negative and ΔS is positive for the reaction.
D) ΔH is positive and ΔS is negative for the reaction.
E) ΔG is positive at all temperatures.
3. For the reaction system that is at equilibrium, which of the following must always be true?
A) ΔG = 0
B) ΔH = 0
C) ΔU = 0
D) ΔS = 0
E) q = 0
Did you read my response to your earlier post with this same question? If not you should go back to read it. If so, then please go through and explain, in detail, exactly what you don't understand about my answers. If I answer this post I'll just rewrite what I've already written earlier.
To determine if a system is positive or negative according to a reaction, you can use the concept of Gibbs Free Energy (ΔG), enthalpy (ΔH), and entropy (ΔS). The sign of ΔG indicates whether a reaction is spontaneous (negative ΔG) or non-spontaneous (positive ΔG) at a given temperature. Let's go through each problem and how to approach them:
1. In the first problem, the question asks which statement is true for the given reaction. To determine the spontaneity, you need to look at the sign of ΔG. If a reaction is spontaneous at all temperatures, it means ΔG is negative. So, you can eliminate options A, B, C, and D. The correct answer is actually E, ΔH is positive and ΔS is negative. This is because the reaction is exothermic (releasing heat) and the surroundings become more ordered (decrease in entropy).
2. For the second problem, you need to predict the signs of ΔH and ΔS for the given reaction. This can be done by observing the phases of the reactants and products. Since all the species in the reaction are gaseous, the change in entropy (ΔS) will be positive. However, you cannot determine the sign of enthalpy (ΔH) just by looking at the phases. In this case, the correct answer is C, ΔH is negative and ΔS is positive, as gas-phase combustion reactions typically release energy (exothermic) and have an increase in entropy due to gas formation.
3. The third problem asks which statement must always be true for a reaction system at equilibrium. At equilibrium, ΔG is equal to zero (ΔG = 0), so option A is incorrect. ΔH represents the change in enthalpy, which may or may not be zero at equilibrium, depending on the specific reaction. ΔU represents the change in internal energy, and it can be zero but not necessarily at equilibrium. ΔS represents the change in entropy and it can be zero, but it depends on the reaction conditions or system. The only option that must always be true at equilibrium is option E, q = 0, meaning at equilibrium there is no heat transfer between the system and its surroundings.
Remember, for a deeper understanding of these concepts, studying the thermodynamics laws and equations will be helpful overall.