Understanding Absolute Value in Thermodynamics

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In thermodynamics, the expression |ΔH|>|TΔS| indicates that the absolute value of the change in enthalpy must exceed the product of temperature and the change in entropy for a synthesis in a cell to be valid. This relationship implies that ΔG, the Gibbs free energy change, is negative (ΔG<0), which is necessary for spontaneity. The use of absolute values emphasizes that only the magnitude of the changes matters, regardless of their sign. Without absolute values, the interpretation could lead to confusion about whether the changes are positive or negative. Understanding this concept is crucial for grasping thermodynamic processes in chemical reactions.
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Hi everyone,I have a problem.I get a short thermodynamics course.Until now,I have thought,that I know,what absolute value means.
And here comes my problem.In lecture is writen,that if we want to make a synthesis in cell valid,we must to have:|ΔH|>|TΔS|.So what this expression means?Does it mean,ΔG<0,in other words?And how I know,that we are dealing with negative or positive numbers in this expression?And if yes,why this expression is expressed like absolute value?If we have it without it,what would happen?I asked teacher,but I m still confused.
Thanks for answers and time,you lost with my question.
 
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With absolute values, only the magnitude of the change matters. It doesn't matter if the change is positive or negative. So H must change more than TS.
 

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