Advertisements
Advertisements
प्रश्न
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]
Advertisements
उत्तर
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
To find:
The total heat required to carry out the given reaction using 180 g of ice.
Calculation:
\[\ce{\underset{\text{(ice at 0 °C)}}{H2O_{(s)}} ->[Latent heat][of fusion 0 °C]\underset{\text{(water at 0 °C)}}{H2O_{(l)}}->[Heating]\underset{\text{(water at 100 °C)}}{H2O_{(l)}} ->[Latent][haeat of vaporization 100 °C] \underset{\text{(Stream at 100 °C)}}{H2O_{(g)}}}\]
a) H2O(s) → H2O(l)
0 °C 0 °C
Heat required = Latent heat for 180 g.
1 mol of H2O = 6.01 kJ
1 mol of H2O = 18 g
∴ 180 g of H2O = `(180 "g")/(18 "g mol"^-1)` = 10 moles of H2O
∴ 10 mol of H2O requires = 60.1 kJ
∴ Heat required = 60.1 kJ …(i)
b) H2O(l) → H2O(s)
0 °C 100 °C
Heat required = Mass × Specific heat × ΔT
= 180 g × 4.18 J g-1 K-1 × 100 K
= 75240 J
= 75.240 kJ ....(ii)
c) H2O(l) → H2O(g)
100 °C 100 °C
Heat required = Latent heat of vaporization
1 mol of H2O requires = 40.7 kJ
∴ 1 mol of H2O = 18 g
∴ 180 g of H2O = 10 moles of H2O
∴ Heat required by 10 moles of water = 407 kJ ….(iii)
From (i) , (ii) and (iii),
Total heat required to carry out the given reaction using 180 g of ice
= 60.1 kJ + 75.240 kJ + 407 kJ = + 542.34 kJ
The total heat required to melt 180 g of ice at 0 °C, heat it to 100 °C and then vaporize it at that temperature is + 542.34 kJ.
APPEARS IN
संबंधित प्रश्न
Answer the following in one or two sentences.
What is standard state of a substance?
Answer in brief.
How much heat is evolved when 12 g of CO reacts with NO2? The reaction is:
4CO(g) 2NO2(g) → 4CO2(g) + N2(g), ΔrH° = - 1200 kJ
Answer the following question.
Calculate ΔrH° for the following reaction at 298 K:
1) 2H3BO3(aq) → B2O3(s) + 3H2O(l), ΔrH° = + 14.4 kJ
2) H3BO3(aq) → HBO2(aq) + H2O(l), ΔrH° = - 0.02 kJ
3) H2B4O7(s) → 2B2O3(s) + H2O(l), ΔrH° = + 17.3 kJ
Define the Bond enthalpy.
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 ____________.
The volume of oxygen required for complete combustion of 0.25 mole of methane at STP is ______.
The standard heats of formation for CCl4(g), H2O(g), CO2(g), and HCl(g) are −25.5, −57.8, −94.1 and −22.1 kcal mol−1, respectively.
∆H for the reaction
\[\ce{CCl4_{(g)} + 2H2O_{(g)} -> CO2_{(g)} + 4HCl_{(g)}}\] at 298 K
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 ____________.
The heat evolved in the combustion of benzene is given by
\[\ce{C6H6 + 7 1/2O2 -> 6CO2_{(g)} + 3H2O_{(l)}}\]; ΔH = −3264.6 kJ
Which of the following quantities of heat energy will be evolved when 39 g C6H6 are burnt?
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?
Standard entropies of N2(g), H2(g), and NH3(g) are a1, a2 and a3 J K-1 mol-1 respectively. What is value of ΔS° for formation of NH3(g)?
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.
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:
\[\begin{array}{cc}
\phantom{}\ce{H}\phantom{...}\ce{H}\phantom{...................}\ce{H}\phantom{...}\ce{H}\phantom{....}\\
\phantom{.}|\phantom{....}|\phantom{....................}|\phantom{....}|\phantom{.....}\\
\ce{C = C + H - H -> H - C - C - H}\\
\phantom{.}|\phantom{....}|\phantom{....................}|\phantom{....}|\phantom{.....}\\
\phantom{}\ce{H}\phantom{...}\ce{H}\phantom{...................}\ce{H}\phantom{...}\ce{H}\phantom{....}
\end{array}\]
Identify the invalid equation.
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?
Define and explain the term, enthalpy of reaction.
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
Which of the following reactions defines the enthalpy of formation?
Standard enthalpy of combustion of a substance is given. Then Write thermochemical equation.
ΔcH0[C2H5OH(1)] = - 1409 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.
For the reaction, H2 + I2 ⇌ 2HI; ΔH = 12.4 kcal. The heat of formation of HI, ΔHf = ______.
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: \[\ce{C2H2_{(g)} + 5/2O_{2(g)}-> 2CO_{2(g)} + H2O_{(l)} \Delta_{(c)}H^{0} = - 1300 kJ}\]
