Thermochemistry Questions

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.)

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 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 (qreaction) 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 to be supplied to the system by the surroundings) then that process is described as endothermic.

word image 16

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:

 

 

 

 

So therefore:

and

Where:

 

 

 

qsystem =

=

qreaction =

=

qcontents = qcalorimeter =

m = c =

ΔT =

C =

qreaction +qcontents +qcalorimeter Equation 1a

0

–(qcontents +qcalorimeter) Equation 1b

–(m × c × ΔT + C × ΔT) Equation 1c

m × c × ΔT

C × ΔT mass of contents

Specific Heat of contents (assumes solution is pure water) temperature change

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 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.

word image 17

 

Where:

 

 

 

ΔGreaction

ΔGreaction

ΔHreaction

T ΔSreaction

=

=

=

=

=

ΔHreaction – T × ΔSreaction Equation 2

change in Gibb’s Free Energy change in Enthalpy temperature at which reaction takes place change in Entropy

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.

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 TS 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 (<3%) NH3

non hazardous

2 M hydrochloric acid (<10%) HCl

corrosive

ammonium chloride NH4Cl

harmful

 

Experiment 7F equipment located in your TECH BIN.

2 x 250mL glass beaker

1 * 100mL measuring cylinder 1 * 100mL beaker

1 x 250mL or 500mL measuring cylinder

Additional equipment and location

On the bench: Calorimeter with thermometer and stirrer, timer or stopwatch, ammonium chloride with microbalance.

On trolley: Ammonia solution 0.4M and Hydrochloric acid 2M

Please ensure this is correct and your area is clean and tidy before leaving the laboratory!

 

 

PART ONE DETERMINATION OF ΔHneutralisation OF AQUEOUS AMMONIA SOLUTION WITH DILUTE HYDROCHLORIC ACID

PROCEDURE

  1. Clean and dry the calorimeter.
  2. Using a measuring cylinder, add 250 mL aqueous ammonia solution ([NH3(aq)] = 0.40 mol L-1) to the calorimeter.
  3. Stir the solution and record its temperature in your report book until you are sure that it is constant.
  4. Using a measuring cylinder, add 50 mL hydrochloric acid solution ([HCl(aq)] = 2.00 mol L-1) to the calorimeter. (For our purposes, you may assume that the temperature of the acid is the same as that of the ammonia solution.) 5 Place the lid on the calorimeter.
  5. Stir the contents continuously.
  6. Record the temperature at 60 second intervals in Table 8.1 in your report book until you are convinced that it is stable.

PART TWO DETERMINATION OF ΔHsolution OF AMMONIUM CHLORIDE

PROCEDURE

1 Place 300 mL of de-ionised water into your clean, dry calorimeter.

2

Weigh precisely, approximately 5.4 g ammonium chloride into a small beaker.

3

Record the temperature of the contents of the calorimeter in Table 8.2. Do not continue until the temperature is stable.

4

Empty the contents of the small beaker into the calorimeter and replace the lid. If significant traces of ammonium chloride remain in the beaker you will need to reweigh it and adjust your recorded weight.

5

Stir the contents and record the temperature until it is stable in Table 8.2.

 

 

 

 

 

PART THREE APPLICATION OF HESS’S LAW TO DETERMINE ΔHformation OF AMMONIUM CHLORIDE

Hess’s Law can be stated as follows: when reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.

Apply Hess’s Law to your experimental data and the data contained in the appendix to this chapter to determine the ΔHformation of ammonium chloride, NH4Cl(s). Diagram 8.2 may help you.

 

Diagram 8.2

ΔHformation = ???

½ N2(g) + 2 H2(g) + ½ Cl2(g) NH4Cl(s)

ΔHformation(NH3 (aq)) + -ΔHsolution(NH4Cl) ΔHformation(HCl(aq))

 

 

NH3(aq) + HCl(aq) NH+4(aq) + Cl–(aq)

 

PART FOUR DETERMINATION OF SPONTANEITY

Answer the questions in your report book.

APPENDIX

Density of Water as a Function of Temperature

Density of Water

15

17

19

21

23

25

27

29

31

33

35

0.994

0.995

0.996

0.997

0.998

0.999

1

temperature (°C)

d

e

n

s

i

t

y

 

a

s

 

a

 

f

u

n

c

t

i

o

n

 

o

f

 

t

e

m

p

e

r

a

t

u

r

e

 

a

t

 

1

 

A

t

m

o

s

p

h

e

r

e

 

Heat Capacity of Calorimeter Ccalorimeter

100 J K-1

Specific Heat of Pure Water cwater =

Standard Enthalpy Values

4.18 ± 0.01 J g-1 K-1

½ N2(g) + 1½ H2(g) + aq →

NH3(aq) ΔH° = –80.7 kJ mol-1

½ Cl2(g) + ½ H2(g) + aq →

HCl(aq) ΔH° = –165.5 kJ mol-1

C(s) + O2(g) →

CO2(g) ΔH° = –393 kJ mol-1

CO(g) + ½ O2(g) →

CO2(g) ΔH° = –282 kJ mol-1

Sample calculation

A neutralisation reaction performed in aqueous solution with a total volume of 400 ml carried out at room temperature (25°C) results in a temperature rise of 2.3°C (= 2.3 K). What is qreaction?

Solution:

Using Equation 1c: qreaction = –(m × c × ΔT + C × ΔT)

All of the terms in Equation 1c are known or provided except for m, which can be calculated from the volume of water given using the equation density = mass/volume.

Using the density vs temperature graph above, the density of water at 25°C is 0.9970 g ml-1 (g ml-1 =

g cm-3). Therefore, mass = density x volume = 0.9970 g ml-1 x 400 ml = 398.8 g. Now, qreaction = –(398.8 g × 4.18 J g-1 K-1 × 2.3 K + 100 J K-1 × 2.3 K) = – 4064 J

REPORT BOOK QUESTIONS

These questions from the Report Book are included in the manual to help you prepare for the Experiment and the Computer Preliminary Exercises.

  1. a) Write an equation describing the reaction of HCl with water.
    1. 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.)

  1. Calculate the energy transfer for the calorimeter alone (qcalorimeter).
  2. Did energy flow from the calorimeter to its contents or vice versa?
  3. 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.

  1. a) State Hess’s Law.
    1. How does Hess’s Law apply in this experiment?
  2. 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?
  3. 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.)

Information Sheet

CORROSIVE

(causes severe burns)

CONCENTRATED HYDROCHLORIC ACID

IDENTIFICATION

Name Hydrochloric acid

Structure HCl

PHYSICAL DESCRIPTION AND PROPERTIES

Description

is water saturated with hydrogen chloride gas to give a colourless to yellow fuming liquid (~37% HCl)

 

is a strong acid, pKa -2.2

Boiling Point

not relevant

Melting Point

not relevant

Vapour Pressure

not relevant

Flammability

non combustible

Density

1.194 g/ml at -26° C

Solubility

very soluble: water

Reactivity

Reacts with many metals, generating highly flammable hydrogen gas, which can explode. Acid must always be added to water; water should not be added to concentrated acid.

HEALTH HAZARD INFORMATION

Major hazards

Highly corrosive; causes severe burns upon eye and skin contact and upon inhalation of fumes.

Toxicity

Inhalation: Inhalation of hydrochloric acid gas can cause severe irritation and injury to the upper respiratory tract and lungs and exposure to high concentrations may cause death. The effects of inhalation however are severe enough to usually lead to prompt withdrawal before further effects can take place.

Eye contact: The liquid and the hydrogen chloride gas that evolves are strong eye irritants and lachrymators (tear producing substances). Contact of the acid or fumes with the eyes may cause severe injury, resulting in permanent impairment of vision and possible blindness.

Skin contact: Contact with the skin may cause severe burns. Frequent contact of the dilute acid with the skin can cause dermatitis.

Swallowing: Ingestion of hydrochloric acid can severe burns of the mouth, throat, and gastrointestinal system and can be fatal.

FIRST AID INFORMATION

Eyes: Hold the eyelid wide open, wash the eye for at least 10 minutes with flowing water.

Lungs: Remove patient to fresh air. If unconscious, do not give anything to drink. If no pulse, apply artificial respiration or cardiopulmonary respiration. If there is a pulse, place in coma position. If conscious, make the casualty lie or sit quietly, give oxygen if available. Lung congestion may occur-a conscious person with breathing difficulties should be placed in a sitting position.

Mouth: DO NOT INDUCE VOMITING as the acid will cause burns on the gastrointestinal tract again as it comes up. If unconscious, treat as described above. If conscious, drink lots of water. Do not attempt to neutralise the acid.

Skin: Remove contaminated clothing immediately. Wash acid off skin with running water for at least 10 minutes.

DISPOSAL OF SMALL AMOUNTS/SPILLAGES BY DEMONSTRATORS

Cover the spill with solid sodium bicarbonate and slowly and carefully mix to a slurry. Carefully scoop up and wash down the drain with plenty of running water. Allow 2-3 minutes between each stage as heat is produced by neutralisation.

Other Information

The full Material Safety Data Sheet for this chemical is available from Chemwatch Gold, on-line at:

http://www.adelaide.edu.au/hr/ohs/legislation/chemwatch

 

Information Sheet

HARMFUL

(can affect health if exposed to large doses or to low doses over a long period of time)

AMMONIUM CHLORIDE

IDENTIFICATION

Name Ammonium chloride

Structure NH4Cl

PHYSICAL DESCRIPTION AND PROPERTIES

Description

White odourless solid

Boiling Point

Not available

Melting Point

340° C

Vapour Pressure

not available

Flammability

Not combustible

Density

1.52 g/ml at 25° C

Solubility

Highly soluble: water

Reactivity

 

HEALTH HAZARD INFORMATION

Major hazards

Can irritate the skin and eyes on contact. Inhalation of dust can irritate the respiratory tract and ingestion can irritate the digestive tract.

Toxicity

Inhalation: Inhalation of dust can cause respiratory tract irritation.

Eye contact: Contact of solid with eyes can cause irritation.

Skin contact: Contact of the solid with the skin can cause irritation.

Swallowing: May cause irritation of the digestive tract. Is harmful when swallowed. May cause acidosis (reduced alkalinity of the blood and body tissues).

FIRST AID INFORMATION

Eyes: Hold the eyelid wide open, wash the eye for at least 10 minutes with flowing water. Summon medical assistance immediately.

Lungs: Remove patient to fresh air. If unconscious, do not give anything to drink. If no pulse, apply artificial respiration or cardiopulmonary respiration. If there is a pulse, place in coma position. If conscious, make the casualty lie or sit quietly, give oxygen if available.

Mouth: If unconscious, treat as described above. If conscious, induce vomiting, rinse out mouth and give lots of water to drink. Seek medical aid immediately.

Skin: Remove contaminated clothing immediately. Wash off skin with running water for at least 10 minutes. Summon medical aid immediately.

DISPOSAL OF SMALL AMOUNTS/SPILLAGES BY DEMONSTRATORS

Any solid spill should be swept up and placed in a container for further disposal. Avoid raising dust. Small quantities of dilute ammonium chloride solutions (<20% w/w) may be washed to waste with running water.

Other Information

The full Material Safety Data Sheet for this chemical is available from Chemwatch Gold, on-line at:

http://www.adelaide.edu.au/hr/ohs/legislation/chemwatch

 

EXPERIMENT 2

THERMOCHEMISTRY

Extra Background

INFORMATION

Heat is defined as the energy that flows into or out of a system because of a difference in temperature between the system and its surroundings. Heat will flow from a region of high temperature to a region of low temperature until a thermal equilibrium is established. Heat is denoted by the symbol q. When heat is absorbed by a system then q will be positive and when heat is removed from a system then q will be negative. The heat of reaction is the value of q required to return a system to a given temperature when the reaction is complete. An exothermic reaction is one where heat is evolved (q is negative). An endothermic reaction is one in which heat is absorbed (q is positive). Enthalpy is a property of a substance that can be used to obtain the heat absorbed or evolved in a chemical reaction. The change in enthalpy for a reaction at a given temperature and pressure is obtained by subtracting the enthalpy of the reactants from the enthalpy of the products. The symbol Δ is used to denote change in something.

ΔH = Hfinal – Hinitial

ΔH = qp the enthalpy of reaction equals the heat of reaction at constant pressure

When manipulating thermochemical equations there are two important things to remember:

When multiplying a thermochemical equation by any factor (doubling the number of moles used) then the value of ΔH is also multiplied by the same factor (doubling the ΔH when doubling the number of moles).

When a thermochemical equation is reversed the value of ΔH is reversed.

The heat capacity (C) of a substance is the quantity of heat (q) needed to raise the temperature of the sample one degree Celsius. Therefore q = CΔt where Δt is the change in temperature. The molar heat capacity is the heat capacity of one mole of substance. The specific heat capacity (c, also called specific heat) is the quantity of heat required to raise the temperature of one gram of a substance by one degree Celsius at constant pressure. q = c × m × Δt where the mass m is in grams.

How do we measure the heat of a reaction? We use a calorimeter which is a device use to measure the heat absorbed or evolved during a chemical reaction.

What can we do with ΔH values? We can use ΔH from one reaction to help calculate the ΔH for another reaction. We use Hess’s Law to assist with this calculation. Hess’s Law states that if a chemical reaction can be written as the sum of other equations, the ΔH of the overall equation equals the sum of the individual ΔH values. An example is shown below:

2CO (g) + O2(g)

2CO2(g)

ΔH = -566.0 kJ (kJ is kilojoules)

C(graphite) + O2(g)

CO2(g)

ΔH = -393.5 kJ

We can use these equations and ΔH values to calculate the ΔH for the following reaction:

2C(graphite) + O2(g)

Now:

2CO(g)

 

2CO2(g)

2CO(g) + O2(g)

ΔH = +566.0 kJ

2C(graphite) + 2O2(g)

2C(graphite) + O2(g)

2CO2(g)

2CO(g)

ΔH = 2 x -393.5 kJ = -787.0 kJ

ΔH = -221.0 kJ

 

While an exothermic reaction will have a negative ΔH and so heat is evolved, does this always mean that the reaction will occur readily? How do we know whether a reaction will take place? A spontaneous process is a chemical reaction that will occur by itself. We need information about enthalpy (ΔH) and entropy (ΔS) to decide whether a reaction will be spontaneous.

What is entropy? Entropy(S) is a thermodynamic quantity that is a measure of the randomness or disorder in a system. Entropy increases with temperature and when a substance changes from solid to liquid and then to gas.

For a spontaneous process to occur the total entropy of the system must increase.

The Gibbs Free Energy (G) is the quantity that takes into account the enthalpy and entropy of a system and can be used to decide whether a reaction will occur spontaneously.

ΔG = ΔH – TΔS

If ΔG is negative then the reaction will be spontaneous and if ΔG is positive the reaction is non spontaneous. If ΔG is zero then the reaction is in equilibrium.

 

 

 

CALORIMETERS

word image 18

 

  1. The Heat Capacity is the amount of heat required to raise the temperature of an object (in this case a calorimeter) by one Kelvin (degree Celsius).

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