Description
1. a) Write an equation describing the reaction of HCl with water.
b) Hence, write an equation describing the reaction between an aqueous solution of ammonia and hydrochloric acid. Indicate the acid/base conjugate pairs.
2. Calculate the energy absorbed by the calorimeter alone, qcalorimeter. Refer to Laboratory Manual, Page 2–3.
3. Calculate the energy absorbed by the contents of the calorimeter, qcontents. You may assume that the specific heat and density of the solution equals that of pure water. Refer to Laboratory Manual, Page 2–3.
4. Calculate the total thermal energy, qreaction, of the reaction.
5. Calculate the thermal energy per mole of reactants. Is this energy released or absorbed?
6. Was the reaction exothermic or endothermic?
7. Report the value for the molar enthalpy of neutralisation (ΔHneutralisation) of NH3(aq) with HCl(aq). Be sure to quote the correct units and comply with sign conventions. This value will vary with the concentration of the reactants, so when quoting your ΔHneutralisation you should state the concentration at which it was determined.
8. Calculate the energy required to decrease the temperature of the contents of the calorimeter by the observed ΔT (qcontents). You may assume that the specific heat and density of the ammonium chloride solution are the same as pure water. (Remember, mcontents = the total mass of everything you have added to the calorimeter.)
9. Calculate the energy transfer for the calorimeter alone (qcalorimeter).
10. Did energy flow from the calorimeter to its contents or vice versa?
11. Calculate the enthalpy change when one mole of ammonium chloride, NH4Cl(s), is dissolved in water. Record the concentration of the resulting solution as this will affect the value obtained.
12. a) State Hess’s Law. b) How does Hess’s Law apply in this experiment?
13. In Part One you examined a spontaneous process that was exothermic. In Part Two you examined a spontaneous process that was endothermic. What does this tell you about a spontaneous process?
14. You have considered spontaneous reactions with positive and negative enthalpy changes (ΔHreaction). Is it possible to predict the spontaneity of a process given the enthalpy change of that process alone? If not, what does determine spontaneity and thus what other factors must also be considered? (Use an equation to support your answer.)
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Explanation & Answer
I made a version where the calculations are in function of the temperatures, I just need those values to finish.
1.
a)
𝐻𝐶𝑙(𝑎𝑞) + 𝐻2 𝑂(𝑙) → 𝐻3 𝑂 +(𝑎𝑞) + 𝐶𝑙 − (𝑎𝑞)
b)
𝐻𝐶𝑙(𝑎𝑞) + 𝑁𝐻3 (𝑎𝑞) → 𝑁𝐻4+(𝑎𝑞) + 𝐶𝑙 −(𝑎𝑞)
Acid and conjugate base: 𝐻𝐶𝑙/𝐶𝑙 −, by the Brönsted-Lowry definition HCl is a proton
donor.
Base and conjugate acid: 𝑁𝐻3 /𝑁𝐻4+, by the Brönsted-Lowry definition HCl is a proton
acceptor.
2.
𝑞𝑐𝑎𝑙𝑜𝑟𝑖𝑚𝑒𝑡𝑒𝑟 = 𝐶 ∙ ∆𝑇
For both experimental measurements, 𝐶 = 100 𝐽/𝐾 and ∆𝑇 is the temperature difference
between the end and the beginning of the reactions.
3.
𝑞𝑐𝑜𝑛𝑡𝑒𝑛𝑡𝑠 = 𝑚 ∙ 𝑐 ∙ ∆𝑇
In the neutralization reaction, 250 mL of ammonia solution and 50 mL of acid solution are
added to the calorimeter, assuming that the mixture has the properties of water at 25°C:
𝑚 = 300 𝑚𝐿 ∙ 0.997
𝑐 = 4.18
𝑔
= 299.1 𝑔
𝑚𝐿
𝐽
𝑔. 𝐾
𝑞𝑐𝑜𝑛𝑡𝑒𝑛𝑡𝑠 = 1250.2 ∙ ∆𝑇
For the solution process, 5.4 g of the salt are added to 300 mL of water; again, considering
that the mixture has the same properties of water at 25°C, we would get the same result as in the
neutralization reaction.
Therefore for both experimental measurements 𝑞𝑐𝑜𝑛𝑡𝑒𝑛𝑡𝑠 = 1250.2 ∙ ∆𝑇, where ∆𝑇 is the
temperature difference between the end and the beginning of the reactions.
4.
𝑞𝑠𝑦𝑠𝑡𝑒𝑚 = 𝑞𝑟𝑒𝑎𝑐𝑡𝑖𝑜𝑛 + 𝑞𝑐𝑎𝑙𝑜𝑟𝑖𝑚𝑒𝑡𝑒𝑟 + 𝑞𝑐𝑜𝑛𝑡𝑒𝑛𝑡𝑠
Since the system is well insulated, 𝑞𝑠𝑦𝑠𝑡𝑒𝑚 = 0:
𝑞𝑟𝑒𝑎𝑐𝑡𝑖𝑜𝑛 = −(𝑞𝑐𝑎𝑙𝑜𝑟�...