The University of Adelide Thermochemistry Questions

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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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Laboratory Manual THERMOCHEMISTRY EXPERIMENT 7F THERMOCHEMISTRY This experiment is done in pairs. Useful background reading (this is not compulsory but may be helpful): “Chemistry Core Concepts”, 1st Editions, John Wiley & Sons Australia Ltd. Section 8.1, 8.3 (page 329-340), 8.4, 8.5 Worked examples 8.3, 8.7; Practice Exercises 8.4, 8.6 “Chemistry Core Concepts”, 2nd Editions, John Wiley & Sons Australia Ltd. Section 8.1, 8.3 (page 466-481), 8.4, 8.5 Worked examples 8.3, 8.7; Practice Exercises 8.4, 8.6 What is the relevance of this prac? This experiment demonstrates Hess’s Law, which is covered in the Chemical Energy section of lectures. You will get the chance to collect temperature data and then work through the calculations required to determine the enthalpies of neutralisation (ΔHneut) and solution (ΔHsoln) and then use these in conjunction with Hess’s Law to calculate the enthalpy of formation (ΔHf). Learning Objectives (remember these are different to the scientific objectives): On completion of this practical, you will have:  Gained an understanding of the principles of thermochemistry.  Learned how to use a constant pressure calorimeter.  Become familiar with the equations and calculations required to determine enthalpies of neutralisation and solution.  Learned how to apply Hess’s Law in the use of thermodynamic data to calculate the enthalpy of formation. Introduction (see pages 13 and 14 of this prac script for further background information) It is recommended that you read the chapters relevant to thermochemistry in your text book before continuing (see chapter references under “Useful background reading” above). Thermodynamics is the study of energy changes. Thermochemistry is the study of energy changes of chemical reactions. Thermodynamics allows us to calculate the maximum energy that may be released (and therefore work done) when a certain process occurs. This process might be the controlled burning of one litre of petrol in a car engine or the illumination of one square metre of solar water heater for one day etc. A calorimeter (Diagram 8.1) is a device that allows the heat change for a process, (q Joules), to be determined. Under normal circumstances the Heat Capacity (C) should be determined EXPERIMENT 7F—1 Laboratory Manual THERMOCHEMISTRY for the calorimeter before it is used.1 For the purposes of this experiment, the Heat Capacity of your calorimeter is included in the appendix to this chapter. Diagram 8.1 thermometer stirrer calorimeter HEAT Although ‘heat’ is a commonly used term in thermodynamics and elsewhere, it is difficult to define clearly. Its use as a verb is not problematic and has no special meaning in thermodynamics and as such has been used freely in this script. However, as a noun the following definition is required: If energy flows between two bodies, or between a system and its surroundings, due solely to their temperature difference then the energy transferred is heat. In most cases ‘energy’ or ‘thermal energy’ can replace ‘heat’ with no loss of meaning. If the process occurs at constant pressure it can be shown that the heat change (q reaction) is equivalent to the enthalpy change (Hreaction). Heat changes are not measured directly but by the effect the heat has on the temperature of the substances involved. Therefore we need to determine the heat capacity, (C, Joules K-1), of the calorimeter and the mass and specific heat, (c, Joules K-1 g-1), of any substances involved. This step represents the calibration of the calorimeter. Once this is done the heat required to raise the temperature of the inside of the calorimeter can be calculated. Any process that gives off heat is called an exothermic process. If chemical energy is converted to thermal energy during a chemical reaction then that process is described as exothermic. If thermal energy is converted to chemical energy during a chemical reaction (that is, heat has 1The Heat Capacity is the amount Kelvin (degree Celsius). EXPERIMENT 7F—2 of heat required to raise the temperature of an object (in this case a calorimeter) by one Laboratory Manual THERMOCHEMISTRY to be supplied to the system by the surroundings) then that process is described as endothermic. Reaction progress HEAT AND CALORIMETRY The calorimeters that you will use are well insulated and you can assume that there will be neglible heat conduction from the system during the experiments that you perform. The calorimeter cannot be sealed, however, to prevent vapour escaping from the system. Under these conditions the following is true: qsystem qreaction So therefore: and Where: = qreaction +qcontents +qcalorimeter = 0 = –(qcontents +qcalorimeter) Equation 1b = –(m × c × ΔT + C × ΔT) Equation 1c qcontents = m × c × ΔT qcalorimeter = C × ΔT Equation 1a m = mass of contents c = Specific Heat of contents (assumes solution is pure water) ΔT = temperature change C = Heat Capacity of calorimeter GIBBS FREE ENERGY AND SPONTANEITY In thermodynamics, as in many other disciplines, words in common usage take on specific meanings that differ slightly but importantly from their more general definition. The word ‘spontaneous’ is one such word. In the field of thermodynamics a spontaneous process is one that lowers the free energy of the system. Such a process is thermodynamically favourable but will only occur if a suitable pathway exists. A process may be spontaneous and yet proceed almost infinitely slowly. Gibbs Free Energy (G) is a thermodynamic function that can be used EXPERIMENT 7F—3 Laboratory Manual THERMOCHEMISTRY to predict if a process is spontaneous at constant temperature and pressure, i.e. a function is spontaneous if ΔG < 0. A spontaneous process releases energy and can be used to do work. Conversely, a nonspontaneous process requires work to be done to cause it to happen. Reaction progress Chemistry connections… Diamonds are not forever! The formation of diamonds is not a spontaneous process at standard temperature and pressure – it requires very high pressure and temperature deep in the Earth’s mantle to convert carbon-bearing minerals into diamonds. However, the reverse reaction, going from diamond to graphite (another form of carbon), has a negative value for ΔG: it is a spontaneous reaction! However, the process is so slow as to be unnoticeable on a human timescale. So all of you who own diamond jewellery, don’t worry; it won’t be crumbling away as graphite anytime soon! A thorough discussion of Gibbs Free Energy is not possible here but it can be established that, for any reaction: Where: ΔGreaction = ΔHreaction – T × ΔSreaction ΔGreaction = change in Gibb’s Free Energy ΔHreaction = change in Enthalpy T = temperature at which reaction takes place ΔSreaction = change in Entropy Equation 2 ENTROPY AND SPONTANEITY Entropy is a direct measure of the randomness or disorder of a system. The second law of thermodynamics tells us that to be spontaneous a reaction must lead to an increase in the entropy of the universe. EXPERIMENT 7F—4 Laboratory Manual THERMOCHEMISTRY Equation 2 says that for a process carried out at temperature T, if the changes in enthalpy and entropy of the system are such that the right hand side of the equation is less than zero, the process must be spontaneous. In order to predict the sign of G, we need to know both H and S. A negative H (an exothermic reaction) and a positive S (a reaction that results in an increase in disorder of the system) tend to make G negative, although a process with a positive H may still be spontaneous if the TS term is large and positive. Although you are not required to calculate the absolute uncertainty for each answer, you should still consider what would be a reasonable number of significant figures to quote. The appendix at the end of this practical script contains physical data that you will need for your calculations. EXPERIMENTAL This experiment can be conveniently divided into four parts: 1 Determination of ΔHneutralisation of aqueous ammonia solution, NH3(aq), with dilute hydrochloric acid, HCl(aq). 2 Determination of ΔHsolution of ammonium chloride, NH4Cl(s) 3 Application of Hess’s Law to your experimental data to determine ΔHformation of ammonium chloride, NH4Cl(s) 4 Consideration of factors that determine the spontaneity of a process Hazardous substances 0.4 M ammonia solution (
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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�...


Anonymous
Excellent resource! Really helped me get the gist of things.

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