Advertisements
Advertisements
प्रश्न
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.
Advertisements
उत्तर
Given:
Enthalpy change (ΔH) = −620 J
Volumes of reactants; \[\ce{C2H4}\] = 100 ml, \[\ce{H2}\] = 100 ml
Pressure (Pext) = 1 bar
To find:
Pressure volume work (W) and change in internal energy (ΔU) for the given reaction
Formulae:
- W = −Pext ΔV
- ΔH = ΔU + Pext ΔV
Calculation:
According to the equation of reaction 1 mole of \[\ce{C2H4}\] reacts with 1 mole of \[\ce{H2}\] to produce 1 mole of \[\ce{C2H6}\]. Hence, 100 ml of \[\ce{C2H4}\] would react with 100 ml of \[\ce{H2}\] to produce 100 ml of \[\ce{C2H6}\].
V1 = 100 mL + 100 mL = 200 mL = 0.2 dm3
V2 = 100 mL = 0.1 dm3
From formula (1),
W = −Pext ΔV
= −1 bar (0.1 dm3 − 0.2 dm3)
= 0.10 dm3 bar
= 0.10 dm3 bar × 100 `"J"/("dm"^3 "bar")` = + 10.00 J
∴ − Pext = ΔV = 10.00 J
∴ Pext = ΔV = −10.00 J
From formula (2),
ΔH = ΔU + Pext ΔV
∴ ΔU = ΔH − Pext ΔV
= −620 − (−10.00 J) = −610 J
Pressure volume work (W) = +10.00 J and ΔU = −610 J
APPEARS IN
संबंधित प्रश्न
Answer in brief.
How will you calculate reaction enthalpy from data on bond enthalpies?
Answer the following question.
State Hess’s law of constant heat summation. Illustrate with an example. State its applications.
Calculate the total heat required
a) to melt 180 g of ice at 0 °C
b) heat it to 100 °C and then
c) vapourise it at that temperature.
[Given: ΔfusH° (ice) = 6.01 kJ mol-1 at 0 °C, ΔvapH° (H2O) = 40.7 kJ mol-1 at 100 °C, Specific heat of water is 4.18 J g-1 K-1]
State and explain Hess’s law of constant heat summation.
Define standard enthalpy of formation.
Define the Bond enthalpy.
Does the following reaction represent a thermochemical equation?
\[\ce{CH_{4(g)} + 2O_{2(g)} -> CO_{2(g)} + 2H2O_{(g)}}\], ∆fH° = –900 kJ mol–1
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 ____________.
When 6.0 g of graphite reacts with dihydrogen to give methane gas, 37.4 kJ of heat is liberated. What is standard enthalpy of formation of CH4 (g)?
lf, \[\ce{C_{(s)} + O2_{(g)} -> CO2_{(g)}}\], ∆H = x .........(i)
\[\ce{CO_{(g)} + 1/2O2_{(g)} -> CO2_{(g)}}\], ∆H = y .......(ii)
Then, the heat of formation of CO is:
Enthalpy of formation of two compounds x and y are −84 kJ and −156 kJ respectively. Which of the following statements is CORRECT?
Combustion of glucose takes place as
\[\ce{C6H12O6_{(s)} + 6O2_{(g)} -> 6CO2_{(g)} + 6H2O_{(g)}}\]; ΔH = −72 kcal mol−1
The energy needed for the production of 1.8 g of glucose by photosynthesis will be ___________.
The enthalpy of formation of nitrogen dioxide is +33.2 kJ mol−1. The enthalpy of the reaction \[\ce{2N2_{(g)} + 4O2_{(g)} -> 4NO2_{(g)}}\]; is ____________.
Which of the following equations has ΔfH° and ΔH° same?
Heat of formation of ethane, ethylene acetylene and carbon dioxide are - 136, - 66, - 228 and - 395 (all in kJ) respectively, most stable among them is ______.
Calculate the standard enthalpy of:
\[\ce{N2H_{4(g)} + H_{2(g)} -> 2NH_{3(g)}}\]
If ΔH0(N – H) = 389 kJ mol–1, ΔH0(H – H) = 435 kJ mol–1, ΔH0(N – N) = 159 kJ mol–1.
\[\ce{A -> B}\], ∆H = −10 kJ mol−1, Ea(f) = 50 kJ mol−1, then Ea of \[\ce{B -> A}\] will be ______.
How many moles of helium gas occupies 22.4 Lat 0°c and at 1 atmospheric pressure?
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?
The enthalpy change that accompanies a reaction in which 1 mole of its standard state is formed from its elements in their standard states
Define and explain the term, enthalpy of reaction.
Standard enthalpy of combustion of a substance is given. Then Write thermochemical equation.
ΔcH0[CH3CHO(l)] = - 1166 kJ mol-1
Heat of combustion of methane is - 890 kJ/mol. On combustion of 12 gm of methane in excess of oxygen, ______ heat is evolved.
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.
Calculate ΔsubH of the H2O from the given data:
\[\ce{H2O_{(s)}->H2O_{(l)},}\] ΔfusH = 6.01kJ mol−1
\[\ce{H2O_{(l)}-> H2O_{(g)},}\] ΔVapH = 45.07 kJ mol−1.
Calculate heat evolved for combustion of 13 gm of acetylene (C2H2).
Given: \[\ce{C2H2_{(g)} + 5/2O_{2(g)}-> 2CO_{2(g)} + H2O_{(l)} \Delta_{(c)}H^{0} = - 1300 kJ}\]
Calculate the standard enthalpy of combustion of methane if the standard enthalpy of formation of methane, carbon dioxide and water are −74.8, −393.5 and −285.8 kJmol−1 respectively.
