The dissolution of ammonium nitrate, NH4NO3, in water is an endothermic process. Since the calorimeter is not a perfect insulato
r, will the enthalpy of solution for ammonium nitrate be reported as too high or too low if this heat change is ignored? Explain
1 answer:
Answer:
Explanation: As already mentioned that the dissolution of ammonium nitrate in water is an endothermic process, which explains that more energy would be needed to break the bond between the reactants rather than the formation of the products.
Hence, the enthalpy of the solution for ammonium nitrate would be positive in nature as energy is being given to the reactants molecules to get converted into the products.
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Answer:
The correct answer is: Ka= 5.0 x 10⁻⁶
Explanation:
The ionization of a weak monoprotic acid HA is given by the following equilibrium: HA ⇄ H⁺ + A⁻. At the beginning (t= 0) we have 0.200 M of HA. Then, a certain amount (x) is dissociated into H⁺ and A⁻, as is detailed in the following table:
HA ⇄ H⁺ + A⁻
t= 0 0.200 M 0 0
t -x x x
t= eq 0.200M -x x x
At equilibrium, we have the following ionization constant expression (Ka):
Ka=
Ka=
Ka=
From the definition of pH, we know that:
pH= - log [H⁺]
In this case, [H⁺]= x, so:
pH= -log x
3.0= -log x
⇒x = 10⁻³
We introduce the value of x (10⁻³) in the previous expression and then we can calculate the ionization constant Ka as follows:
Ka= =
= 5.025 x 10⁻⁶= 5.0 x 10⁻⁶