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First Law of Thermodynamics
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Let us consider one of the most fundamental principles of the physical world - The Law of Conservation of Energy - It states that energy cannot be created or destroyed.
Now, let us apply this law to various processes in chemistry. It can be written as
ΔU = Q + W
Where ΔU is the change in internal energy
Q is heat
W is work
According to this, there are two kinds of processes that can lead to a change in the internal energy of the system - they are heat and work.
Let us consider the sign convention:
(i) Heat flowing into the system is positive. It increases the internal energy of the system.
(ii) Heat flowing out of the system is negative. It decreases the internal energy of the system.
(iii) Work done by the system is positive. It decreases the internal energy of the system.
Note that: U is a state function. From a given value of ?U, one cannot make out whether this change has come about by adding heat to it or by doing work on it.
System work: When we talk about work done by a thermodynamic system, we are usually talking about the work done by a gas in expanding.
(i) Work done at constant pressure
W = PΔV
It is given by the area under the PV graph
(ii) Work done when the pressure is changing
W = ΔPdV
Note: That work and heat are not state functions. They do depend on the path followed between the initial and final states.
Example
Calculate the internal energy change in each of the following cases:
(i) A system absorbs 5kJ of heat and does 1kJ of work.
(ii) 5kJ of work is done on the system and 1kJ of heat is given out by the system.
(i) Q = +5kJ
W = -1kJ
ΔU = Q + W = 5 + (-1) = 4kJ
(ii) W = +5kJ
Q = -1kJ
ΔU = Q + W = -1 + 5 = 4kJ
In both the cases, the interval energy of the system increases.
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