First Law of Thermodynamics for a Closed System

Conservation of Energy

Internal energy (U) is associated with microscopic motions and forces. Since this energy cannot be seen, it is usually separated from the macroscopic i.e., measurable mechanical energy in order to express the total energy of the system as:

E_total = U + Kinetic Energy (E_K) + Potential Energy (E_P)

Let us consider a closed system to which energy in the form of heat and work is supplied from its surroundings:

The first law of thermodynamics states that the total energy of the universe is constant, i.e.,

E_total (universe) = constant    or     Δ E_total (universe) = 0

 

The universe is composed of the system and its surroundings; thus, the above expression can be expressed as follows:

E_total (system) + E_total (surroundings) = 0

The increase in the total energy of the system is given by:

ΔE_total (system) = ΔU + ΔE_K + ΔE_P

Conversely, the decrease in the total energy of the surroundings is given by:

ΔE_total (surroundings) = -Q – W (where Q is heat into or out of the surroundings and W is work done on or by the surroundings)

Combining the above two equations we obtain the first law of thermodynamics for a closed system:

ΔU + ΔE_K + ΔE_P = Q + W

The differential form of this equation is as follows:

dU + dE_K + dE_P = dQ + dW

(Note: Since Q and W are path functions, these quantities in differential form are expressed as δQ and δW in some examples)

If the changes in kinetic and potential energies are negligible the equation can be simplified as:

ΔU = Q + W

The term W includes expansion and non-expansion type of work. Expansion (or contraction) work is related to the change in the volume of system. While non-expansion work includes shaft work (work done on the system by a rotating mechanical device), chemical work, electrical work, etc.

 

 

 

 

Multiple Reactions

 

How to Obtain Rate Data, Reaction Rate Constant and Reaction Order

 

Classification of Reactions

 Classification based on Reactant and Products

Homogeneous Reaction -  These are reactions that occur in only one phase whether it be gas, liquid or solid. Examples include the combination of common household gas and oxygen to produce a flame, the reactions between aqueous solutions of acids and bases.  

Heterogeneous Reaction - These are reactions that require the presence of at least two phase in order to proceed. It doesn't matter if the reaction occurs in one, two, or more phases at the interface, or whether the reactants and products are distributed among the phases or contained in only one phase. The only important factor is that at least two phases are present for the reaction to proceed. 

Classification based on Presence of Catalyst

Catalytic Reaction - In this reactions, a catalyst is used to accelerate the rate by which a specific chemical reaction proceeds. The catalyst provides an alternative, lower energy pathway for the reaction.

The catalyst interacts with a reactant and forms an intermediate/transient compound which after it forms, breaks apart leaving the original catalyst species unaltered. 

The catalyst is not affected by the reaction as far as the chemical structure or mass at reaction completion. 

Non-Catalytic Reaction - These are chemical reactions in which a catalyst is not used. Thus for these reactions the reaction rate does not increase by any external influence. Since catalysts reduce the energy required for a reaction to occur (activation energy), non-catalytic reactions require more energy than catalytic reactions for the reaction to occur. 

Classification of Chemical Reactions Useful in Reactor Design