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Raoult's Law

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Question

why the changing in the energy is calculated in terms of delta H but not in delta u for a chemical reaction

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Solution

The solution is as follows:

The change in internal energy cannot be calculated because it is the energy associated with making and breaking of bonds, translation and vibrational energy change associated with the molecules in the system. It cannot be calculated without taking in account the change in pressure and volume associated with the substance.

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Q. For a chemical reactions, the free energy change

(Δ G)

is negative. The reaction is

Q. For which reaction,

Δ H > Δ U

Δ H =

change in enthalpy,

Δ U =

change in internal energy )

For gaseous reactions, if $Δ H$ is the change in enthalpy & $Δ U$ that in internal energy, then

Chemical reaction is invariably associated with the transfer of energy either in the form of heat or light. In the laboratory, heat changes in physical and chemical processes are measured with an instrument called Calorimeter. Heat change in the process is calculated as :

q = m s Δ T

s = s p e c i f i c h e a t

= c Δ T

c = h e a t c a p a c i t y

Heat of reaction at constant volume is measured using bomb Calorimeter.

q_{V} = Δ U =

internal energy change
Heat of reaction at constant pressure is measured using bomb Calorimeter.

q_{P} = Δ H

q_{P} = q_{V} + P Δ V

Δ H = Δ U + Δ n R T

The heat capacity of a bomb calorimeter is

500 J K^{- 1}

0.1 g

of Methane was burnt in this calorimeter, the temperature rose by

2^{\circ} C

Δ U

per mole will be :

Δ H

is the change in enthalpy and

Δ E

the change in internal energy accompanying a gaseous reaction?

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