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प्रश्न
Answer the following question.
State Hess’s law of constant heat summation. Illustrate with an example. State its applications.
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उत्तर
1. Hess’s law of constant heat summation:
Hess’s law of constant heat summation states that “Overall the enthalpy change for a reaction is equal to the sum of enthalpy changes of individual steps in the reaction”.
2. Illustration:
- The enthalpy change for a chemical reaction is the same regardless of the path by which the reaction occurs. Hess’s law is a direct consequence of the fact that enthalpy is a state function. The enthalpy change of a reaction depends only on the initial and final states and not on the path by which the reaction occurs.
- To determine the overall equation of the reaction, reactants and products in the individual steps are added or subtracted like algebraic entities.
- Consider the synthesis of NH3,
- \[\ce{2H_{2(g)} + N_{2(g)} → N2H_{4(g)}, Δ_r H^0_1 = + 95.4 kJ}\]
- \[\ce{N2H_{4(g)} + H_{2(g)} → 2NH_{3(g)}, Δ_r H^0_2 = -187.6 kJ}\]
\[\ce{H_{2(g)} + N_{2(g)} → 2NH_{3(g)}, Δ_rH^0 = - 92.2 kJ}\]
The sum of the enthalpy changes for steps (1) and (2) is equal to the enthalpy change for the overall reaction.
3. Application of Hess’s law:
The Hess's law has been useful to calculate the enthalpy changes for the reactions with their enthalpies being not known experimentally. To calculate heat of formation, combustion, neutralization, ionization, etc.
संबंधित प्रश्न
Answer in brief.
What is the standard enthalpy of combustion? Give an example.
The enthalpy change for the reaction, \[\ce{C2H4_{(g)} + H2_{(g)} -> C2H6_{(g)}}\] is −620 J when 100 mL of ethylene and 100 ml of \[\ce{H2}\] react at 1 bar pressure. Calculate the pressure volume type of work and ΔU for the reaction.
State and explain Hess’s law of constant heat summation.
Define standard enthalpy of formation.
Classify the following into intensive and extensive properties.
Pressure, volume, mass, temperature.
When 2 moles of C2H6(g) are completely burnt, 3129 kJ of heat is liberated. If ∆Hf for CO2(g) and H2O(l) are −395 and −286 kJ per mole respectively, the heat combustion of C2H6(g) is ____________.
A compound that has a high negative heat of formation is normally ____________.
Daily requirement of energy of a person is 'x' kJ. If heat of combustion of food material (Molecular mass = 100 g) is 'y' kJ, his daily consumption of the food in gram would be ____________.
The enthalpy change accompanying a reaction in which 1 mole of the substance in the standard state reacts completely with oxygen or is completely burnt is called as ____________.
Which among the following salts, solubility decreases with increase in temperature?
Standard enthalpy of formation of water is - 286 kJ mol-1. When 1800 mg of water is formed from its constituent elements in their standard states the amount of energy liberated is ______.
An ideal gas expands isothermally and reversibly from 10 m3 to 20 m3 at 300 K performing 5 .187 kJ of work on surrounding. Calculate number of moles of gas undergoing expansion. (R = 8.314 J K-1 mol-1)
Heat of formation of water is - 272 kJ mol-1. What quantity of water is converted to H2 and O2 by 750 kJ of heat?
Which of the following equations has ΔfH° and ΔH° same?
Calculate the standard enthalpy of formation of CH3OH(l) from the following data:
- \[\ce{CH3OH_{(l)} + 3/2 O2_{(g)} -> CO2_{(g)} + 2H2O_{(l)}ΔH^° = - 726 kJ mol^{-1}}\]
- \[\ce{C_{(s)} + O2_{(g)} → CO2_{(g)}Δ_cH^° = – 393 kJ mol^{-1}}\]
- \[\ce{H2_{(g)} + 1/2 O2_{(g)} -> H2O_{(l)}Δ_fH^° = - 286 kJ mol^{-1}}\]
From the following bond energies:
H – H bond energy: 431.37 kJ mol−1
C = C bond energy: 606.10 kJ mol−1
C – C bond energy: 336.49 kJ mol−1
C – H bond energy: 410.50 kJ mol−1
Enthalpy for the given reaction will be:
\[\begin{array}{cc}
\phantom{}\ce{H}\phantom{...}\ce{H}\phantom{...................}\ce{H}\phantom{...}\ce{H}\phantom{....}\\
\phantom{.}|\phantom{....}|\phantom{....................}|\phantom{....}|\phantom{.....}\\
\ce{C = C + H - H -> H - C - C - H}\\
\phantom{.}|\phantom{....}|\phantom{....................}|\phantom{....}|\phantom{.....}\\
\phantom{}\ce{H}\phantom{...}\ce{H}\phantom{...................}\ce{H}\phantom{...}\ce{H}\phantom{....}
\end{array}\]
What is the amount of water formed by the combustion of 1.6 g methane?
What is enthalpy of formation of NH3 if bond enthalpies as (N ≡ N) = - 941 kJ/mol.
\[\ce{(H - H)}\] = 436 kJ/mol and \[\ce{(N - H)}\] = 389 kJ/mol?
When 0.5 gram of sulphur is burnt to form SO2, 4.6 kJ of heat liberated. Calculate enthalpy of formation of SO2(g). (Atomic mass : S = 32, O = 16)
Calculate the standard enthalpy of the reaction, \[\ce{SiO2_{(s)} + 3C_{(graphite)} -> SiC_{(s)} + 2CO_{(g)}}\] from the following reactions:
- \[\ce{Si_{(s)} + O2_{(g)} -> SiO2_{(s)}}\], ΔrH0 = −911 kJ
- \[\ce{2C_{(graphite)} + O2_{(g)} -> 2CO_{(g)}}\], ΔrH0 = −221 kJ
- \[\ce{Si_{(s)} + C_{(graphite)} -> SiC_{(s)}}\], ΔrH0 = −65.3 kJ
Standard enthalpy of combustion of a substance is given. Then Write thermochemical equation.
ΔcH0[C2H5OH(1)] = - 1409 kJ mol-1
The enthalpy of combustion of S (rhombic) is − 297 kJ mo1-1. Calculate the amount of sulphur required to produce 29. 74 kJ of heat.
Draw energy profile diagram and show:
- activated complex
- energy of activation for forward and backward reactions
- enthalpy of reaction
The heat evolved in the combustion of 6.022 x 1021 carbon particles is 3.94 kJ. The heat of combustion of carbon is ______.
Calculate heat evolved for combustion of 13 gm of acetylene (C2H2).
Given: \[\ce{C2H2_{(g)} + 5/2O_{2(g)}-> 2CO_{2(g)} + H2O_{(l)} \Delta_{(c)}H^{0} = - 1300 kJ}\]
