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Fed [463]
2 years ago
9

In a constant‑pressure calorimeter, 70.0 mL of 0.770 M H2SO4 is added to 70.0 mL of 0.420 M NaOH. The reaction caused the temper

ature of the solution to rise from 23.14 ∘C to 26.00 ∘C. If the solution has the same density and specific heat as water ( 1.00 g/mL and 4.184 J/(g⋅°C), respectively), what is ΔH for this reaction (per mole of H2O produced)
Chemistry
1 answer:
schepotkina [342]2 years ago
3 0

Answer : The enthalpy of neutralization is, -113.9 KJ/mole

Explanation :

First we have to calculate the moles of H_2SO_4 and NaOH.

\text{Moles of }H_2SO_4=\text{Concentration of }H_2SO_4\times \text{Volume of solution}=0.770mole/L\times 0.070L=0.0539mole

\text{Moles of NaOH}=\text{Concentration of NaOH}\times \text{Volume of solution}=0.420mole/L\times 0.070L=0.0294mole

The balanced chemical reaction will be,

H_2SO_4+2NaOH\rightarrow Na_2SO_4+2H_2O

From the balanced reaction we conclude that,

As, 2 mole of NaOH neutralizes by 1 mole of H_2SO_4

As, 0.0294 mole of NaOH neutralizes by \frac{0.0294}{2}=0.0147 mole of H_2SO_4

Thus, the number of neutralized moles = 0.0147 mole

Now we have to calculate the mass of water.

As we know that the density of water is 1 g/ml. So, the mass of water will be:

The volume of water = 70.0ml+70.0ml=140.0ml

\text{Mass of water}=\text{Density of water}\times \text{Volume of water}=1g/ml\times 140.0ml=140.0g

Now we have to calculate the heat absorbed during the reaction.

q=m\times c\times (T_{final}-T_{initial})

where,

q = heat absorbed = ?

c = specific heat of water = 4.184J/g^oC

m = mass of water = 140.0 g

T_{final} = final temperature of water = 26.00^oC

T_{initial} = initial temperature of metal = 23.14^oC

Now put all the given values in the above formula, we get:

q=140.0g\times 4.184J/g^oC\times (26.00-23.14)^oC

q=1675.27J=1.675kJ

Thus, the heat released during the neutralization = -1.675 KJ

Now we have to calculate the enthalpy of neutralization.

\Delta H=\frac{q}{n}

where,

\Delta H = enthalpy of neutralization = ?

q = heat released = -1.675 KJ

n = number of moles used in neutralization = 0.0147 mole

\Delta H=\frac{-1.675KJ}{0.0147mole}=-113.9KJ/mole

The negative sign indicate the heat released during the reaction.

Therefore, the enthalpy of neutralization is, -113.9 KJ/mole

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<h3>Further explanation</h3>

Given

Reaction

4 NH₃ (g) + 5 O2 (g) → 4 NO (g) + 6 H2O (l)

Required

a. volume of NO

b. mass of H2O

c. volume of O2

Solution

Assume reactants at STP(0 C, 1 atm)

Products at 1000 C (1273 K)and 1 atm

a. mol ratio NO : O2 from equation : 4 : 5, so mo NO :

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mol ratio of NO : O2 = 4 : 5, so mol O2 :

\tt \dfrac{5}{4}\times 0.34=0.425

Volume O2 at STP :

\tt 0.425\times 22.4=9.52~L

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