2510 Thermochemistry

THERMOCHEMISTRY: HEATS OF REACTIONS AND HESS'S LAW

1.0 INTRODUCTION

Thermochemistry is a study of energy and heat changes which accompany chemical reactions. Heat given off or absorbed by a chemical reaction at constant pressure is called the enthalpy change, ΔH. A reaction which loses heat to the surrounding is said to be exothermic. Reactions which absorb heat from the surrounding are said to be endothermic.

The sign of ΔH is relative to the surroundings. Exothermic reactions have negative enthalpy changes (ΔH < 0). Conversely, endothermic reactions have positive enthalpies (ΔH > 0). A way to remember this is that if heat is subtracted from the reaction final products (ΔH < 0), it must go to the surroundings and is therefore exothermic.

Thermochemical reactions follow Hess's law, named after Germain Hess in the year 1840. It can be summarized as:

"If a process proceeds through one or several steps, the enthalpy change for the overall process can be calculated as the sum of the enthalpy changes for the individual steps".

This is a useful law. It allows the measurement of enthalpy changes for chemical reactions that are difficult or even impossible to determine directly. It can be written as an algebraic formula:

ΔH (Reaction 1 + Reaction 2) = ΔH (Reaction 1) + ΔH (Reaction 2)

This experiment measures the following exothermic reactions:

• Enthalpy of the reaction of dissolving solid sodium hydroxide in water
• Enthalpy of neutralization reaction of aqueous hydrochloric acid by aqueous sodium hydroxide
• Enthalpy of neutralization reaction of solid sodium hydroxide by an aqueous solution of hydrochloric acid

The overall analysis of results should permit verification of Hess's Law.

Heat and temperature measurements are carried out in equipment called a calorimeter. A calorimeter is designed to insulate a container to minimize loss or gain of heat from the surroundings. A simple one consists of two Styrofoam™ cups nested together. The cups are covered by a corrugated cardboard square with a center hole for a thermometer. No calorimeter is perfectly insulated from its surrounding. Measurements must be corrected due to losses (or gains) from the calorimeter.

The immediate surroundings for the reaction are the reaction solution plus the calorimeter. The measured heat of the reaction, ΔHrxn is equal to the heat gained by the surroundings:

ΔHrxn = - qsurr = - (qsoln + qcal)

Remember that the three reactions that are being studied in this experiment are all exothermic, hence ΔHrxn will be negative.

The heat of the solution, qsoln is given by:

qsoln = msoln x Csoln x ΔTsoln

where msoln is the mass of the solution and ΔTsoln is the temperature change of the solution. The specific heat, Csoln, is the quantity of heat required to raise the temperature of one gram of a material by one degree Celsius. The total heat capacity of the double nested foam cups and corrugated board, qcal, is calculated by the formula:

qcal = Ccal x ΔTcal

The heat capacity of calorimeter, Ccal, is the heat required to raise the temperature of an empty calorimeter by one degree Celsius, and therefore is measured in Joules per degree C. Unless you measure Ccal separately, assume that the heat capacity of the double-nested foam cup shown in the picture is 15 J/°C.

Double-Nested Foam Cup Calorimeter for Measuring Heats of Reactions

At least three factors, make it difficult to measure temperature changes:

• Heat is lost from the calorimeter to the surroundings. This causes the temperature to drop by about 0.1 °C/min
• It takes time for the calorimeter to reach thermal equilibrium. Usually this is at least 1 minute.
• It takes for sodium hydroxide to dissolve, sometimes as much as 4 min.

After correction is made for the heat absorbed by the calorimeter, the standard enthalpy of the reaction, ΔH, is calculated by dividing the negative of the corrected heat by the number of moles of a reactant and/or product.

Standard enthalpy of neutralization is calculated per one mole of water formed in the course of neutralization reaction. To find the number of moles of water formed, use the stoichiometry of the neutralization reaction:

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

A table may help organize the reactions and mixtures present in the parts of this experiment.

Summary Table
	Target mixture	Total final volume of solution	M of HCl, NaOH in solution after mixing	Moles of product
Part A:	50.0 ml of 1.00 M HCl added to 50.0 ml of 1.00 M NaOH	100 ml	0.500	0.0500
Part B:	2.00 g NaOH(s) are added to 50.0 ml of laboratory H2O	100 ml	0.500	0.0500
Part C:	2.00 g of NaOH(s) are added to 50.0 ml of 1.00 M HCl(aq) and 50.0 ml of laboratory H2O	100 ml	0.500	0.0500

The concentrations and amounts listed in the above table are ideal quantities. You will be determining the mass of your NaOH(s) which is to be close to 2.00 grams but may not be exactly 2.00 grams. You will therefore have to calculate the number of moles of NaOH(s) for your grams and it may therefore be a little more or a little less than 0.0500 mol. Similar adjustments might be needed for different concentrations of HCl. Remember to record your values.

2.0 SAFETY PRECAUTIONS AND WASTE DISPOSAL

!!Wear your safety goggles!! Sodium hydroxide and hydrochloric acid are both strong eye irritants. Be sure that the pathway to the eye wash station remains always unobstructed. Solid sodium hydroxide is caustic! Do not handle the pellets! If any solution of acid or base used in this experiment is spilled on your hands, rinse thoroughly with plenty of water. Rinse the Styrofoam™ cups well with laboratory water prior to returning them. The cups are reusable. Wash your hands before leaving the laboratory. Do not eat or drink during this experiment. This experiment uses digital thermometers that should not be completely submerged in liquids. Only the measuring probe should be submerged. If glass thermometers are used, they are fragile. Mercury-containing thermometers have mostly been replaced due to problems with cleanup in the event of a break. If a mercury thermometer breaks, notify your instructor and the stockroom for appropriate and safe cleanup procedures. Dispose of liquids down the drain Once the reaction has been completed, the products are non-toxic (salty water). There are no hazardous wastes to be captured at the completion of this experiment.

3.0 CHEMICALS AND SolutionS

Chemical	Concentration	Approximate amount	Notes
HCl, hydrochloric acid	1.0 M	100 ml
NaOH, sodium hydroxide solid	NA	4 g	Hydroscopic! Solid sodium hydroxide absorbs moisture directly from the air. Weigh quickly and seal the container.

4.0 GLASSWARE AND APPARATUS

Item	Use	Notes
400 ml glass beaker	Store and use 1.00 M NaOH(aq)	If different concentration is used, label it properly.
250 ml glass beaker	Store and use 1.00 M HCl(aq)	If different concentration is used, label it properly.
50-ml graduated cylinders	Volume measurement	Keep separate cylinders for NaOH and HCl solutions
Stopwatch or phone	Time measurement
50 or 100 ml beaker	Weighing container for solid NaOH
Spatula		Be sure to clean residues off spatula
Thermometer	Temperature measurement	Do not submerge electronic portion
2 Styrofoam™ cups (8 oz)	Calorimeter
Cardboard square	Insulation	Cut to fit if needed

5.0 PROCEDURE

Part A: The Enthalpy of Neutralization of HCl(aq) and NaOH(aq)

1. Prepare a double nested foam-cup calorimeter as shown in the picture.

2. Place 50.0 ml (measured by a 50-ml graduated cylinder) of 1.0 M NaOH(aq) into calorimeter.

3. Use another 50-ml graduated cylinder to measure 50.0 ml of 1.0 M HCl(aq). Leave the acid in the cylinder.

4. Measure the temperature of sodium hydroxide solution inside the calorimeter (with the accuracy of ±0.1°C) for 3 minutes at 1-minute intervals.

5. At the 4th minute do not read temperature, but instead quickly mix HCl(aq) into the calorimeter. Immediately close the lid to prevent the loss of heat. Gently swirl the calorimeter for 1 minute to mix the solutions.

6. Measure the temperature every minute after mixing. Continue until the temperature starts dropping at a steady rate for at least 5 minutes.

7. For this part, the complete reaction being studied is:

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

It can also be written as:

H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → Na+(aq) + Cl-(aq) + H2O(l)

There are several species on both sides of the above equation. Because those species are not relevant to the reaction (they are called spectator species because they witness the reaction but do not chemically react), they can be canceled from both sides of the chemical equation. The result is a net ionic reaction:

H+(aq) + OH-(aq) → H2O(l)

This is a neutralization reaction. The enthalpy of this reaction is called the enthalpy of neutralization.

Part B: The Enthalpy of dissolving NaOH(s)

1. Because it is difficult to measure exactly 2.00 grams of sodium hydroxide, use between 1.9 and 2.1 grams. Precisely measure the amount of NaOH(s) to at least 3 significant figures. Record the exact mass in your report.

2. Place 50 ml of laboratory water in your rinsed and dried calorimeter.

3. Record the temperature of the water for 3 minutes at one-minute intervals.

4. At the 4th minute, add the solid NaOH. Gently swirl the calorimeter.

5. Measure the temperature every minute after mixing.

6. It will take several minutes for the NaOH to dissolve, so the temperature will continue rising. Continue until the temperature starts dropping at a steady rate for at least 5 minutes.

7. For this part, the complete reaction being studied is:

NaOH(s) → Na+(aq) + OH-(aq)

There are no spectator species present, so the reaction cannot be further simplified. This reaction is a dissolution (or solvation) of a chemical into a solvent. The enthalpy of this reaction is called the enthalpy of dissolving (sometimes called the enthalpy of solvation).

C. The Enthalpy of the Neutralization of HCl(aq) with NaOH(s)

1. Because it is difficult to measure exactly 2.00 grams of sodium hydroxide, use between 1.9 and 2.1 grams. Precisely measure the amount of NaOH(s) to at least 3 significant figures. Record the exact mass in your report.

2. Place 50.0 ml of 1.00 M HCl(aq) in your rinsed and dried calorimeter.

3. Record the temperature of the solution for 3 minutes at one-minute intervals.

4. At the 4th minute, add the solid NaOH. Gently swirl the calorimeter.

5. Measure the temperature every minute after mixing.

6. It will take several minutes for the NaOH to dissolve, so the temperature will continue rising. Continue until the temperature starts dropping at a steady rate for at least 5 minutes.

7. For this part, the complete reaction being studied is:

H+(aq) + Cl-(aq)+ NaOH(s) → Na+(aq) + Cl-(aq) + H2O(l)

The Cl-(aq) is a spectator ion that does not participate in the reaction. So it can be subtracted from both sides of the equation.

H+(aq) + NaOH(s) → Na+(aq) + H2O(l)

Note that if you combine the reaction in C with the reverse reaction in B (and after eliminating common ions) get the reaction in A. The enthalpy of this reaction will be the combination of the enthalpy of dissolving and the enthalpy of neutralization. This enthalpy does not have a specific name.

6.0 DATA RECORDING SHEET

Last Name		First Name		Partner Name(s)	Date
Temperature Readings (oC)
	Time (minutes)	Part A	Part B	Part C
	0.00 (Start)
	1
	2
	3
	4 Mix reagents
	5
	6
	7
	8
	9
	10
	11
	12
	13
	14
	15
	16
	17
	18
	19
	20

Note: Continue to measure the temperature until it has stabilized.

7.0 CALCULATIONS AND ANALYSIS SHEET

	Part A	Part B	Part C
Amount of NaOH added	Conc: Volume:	Grams:	Grams:
Maximum temperature reached, Tmax, °C
Temperature before mixing, To, °C
Temperature change, ΔT = Tmax - To
Moles of NaOH present in the calorimeter (based on ml of NaOH in Part A or grams of solid NaOH in parts B and C)
Heat gained by the solution: qsoln = msoln x csoln x ΔT	*	**	*
Heat gained by the calorimeter, qcal*** = Ccal x ΔT
Heat gained by both calorimeter and NaCl(aq), qsurr, in Joules
Total heat of the reaction, in kJ ΔHrxn = -qsurr
Enthalpy of the reaction in kJ per mole ΔHrxn/ moles of reaction = standard ΔH

* The density of 0.5 M NaCl(aq) is 1.02 g/mL. 100 mL of this solution is therefore 1.02 g/mL X 100 mL = 102 grams. The specific heat of this solution is approximately 4.19 J/g·°C.

** The density of water is 1.00 g/mL. Assume that the specific heat of this solution is approximately 4.184 J/g·°C.

*** Unless you measure Ccal separately, assume that the heat capacity of the double-nested foam cup shown in the picture is 15 J/°C.

8.0 POST-LAB QUESTIONS AND CONCLUSIONS

1. Write out the reactions for parts A, B, and C showing how A and B combine to give C.

2. Use the enthalpies you calculated for parts A, B, and C. Combine them in the same way as you did the reactions to see if your data is consistent with Hess’s Law.

3. Does your data support Hess' Law? Why or why not? What sources of error could be present that might cause differences between the theoretical value and the observed value of the enthalpy of this reaction? How might you change the experiment to minimize sources of these errors?