ATAR Notes: Forum
VCE Stuff => VCE Science => VCE Mathematics/Science/Technology => VCE Subjects + Help => VCE Chemistry => Topic started by: NE2000 on November 11, 2009, 05:36:53 pm
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OK this is a problem I've thought of for a long time (good time to get it checked yeah). My textbook says when a forward reaction dominates a reverse (high K) the products are usually more stable so the system tends towards them. But then exothermic = more stable products. But then not all reactions with high K are exothermic. So what exactly do they mean by more stable?
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OK this is a problem I've thought of for a long time (good time to get it checked yeah). My textbook says when a forward reaction dominates a reverse (high K) the products are usually more stable so the system tends towards them. But then exothermic = more stable products. But then not all reactions with high K are exothermic. So what exactly do they mean by more stable?
I'm not quite sure about the high K = exothemic statement, however, let's think about an enthalpy graph.
Now, the peak of the graph is called the "activated complex", and that is VERY UNSTABLE, and decomposes into the products. In terms of enthalpy, it has high enthalpy and low stability.
Now, let's consider something like C6H12O6 and oxygen to give us water and carbon dioxide. Think about CO2, it has a lower enthalpy than glucose and oxygen. Now as things in and of itself, carbon dioxide is very stable - (how often have you heard of "highly reactive CO2").
Now, we notice a trend. Things with high enthalpy have low stability and vice versa.
And to answer your question, the products of an exo reaction, which is of lower enthalpy than the reactants are MORE STABLE. :)
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Yeah that's the definition I always operated by. Low enthalpy = greater stability.
But then what about the concept of an equilibrium system favouring products with greater stability. Because of the products are highly stable and unreactive then won't react back to form reactants right? This is what I am confused about. It feels like I am missing something.
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Yeah that's the definition I always operated by. Low enthalpy = greater stability.
But then what about the concept of an equilibrium system favouring products with greater stability. Because of the products are highly stable and unreactive then won't react back to form reactants right? This is what I am confused about. It feels like I am missing something.
Entropy plays a big part, and it is highly temperature dependant. You'll learn about it in chemistry (thermodynamics) at uni, at which point everything starts to make sense (a little bit) :)