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Heat of Neutralization in a Calorimeter with Weak and Strong Electrolytes

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Heat of Neutralization in a Calorimeter with Weak and Strong Electrolytes

Erin

Lab Performed: November 11th, 2008

Section 006

T.A:  Teresa


Introduction

The purpose of this experiment is to find the concentration of the unknown HCl as a strong electrolyte to observe the enthalpy changes between strong and weak electrolytes. This is conducted by using a “coffee-cup” calorimeter and a series of solutions: HCl, HNO3, NaOH (all strong electrolytes) and Phenol (weak electrolyte).

Both strong electrolytes and weak electrolytes are used in this lab. The difference between strong electrolytes and weak electrolytes is that strong electrolytes completely ionize in aqueous solution making the solution a good conductor. In contrast, weak electrolytes do not behave the same way and will only partially ionize, but is only a fairly good conductor of electricity (Petrucci et al, 2007).  The examples of strong electrolytes in this experiment are HCl, HNO3 and NaOH. The phenol solution that it is use is a weak electrolyte.

The reaction is happening within a “coffee-cup” or Styrofoam cup calorimeter where the neutralization reaction between the NaOH and the acid will act as the system. The heat that is absorbed or released by the system during the reaction within the calorimeter, is called the heat of the reaction, q. The formula for q is: q = mcΔt

The heat of the system is comes from the combination of hydrogen ions and hydroxide ions to make water. When the H+ (hydrogen ions) and OH- (hydroxide ions) ions come from two strong electrolytes, they will have no effect on the heat of neutralization, but it is found that reactions with strong electrolytes produce the same amount of heat (-55.90KJ/mol). Likewise, within a reaction with a weak electrolyte, such as phenol, the nature of the solution will determine the heat of the reaction – either smaller or larger than -55.90KJ/mol of H+ (Strathopulous, 2008).

Materials and Methods

The experimental procedure used for this experiment was outlined in the CHEM 120L lab manual, Experiment 4. All steps were followed without deviation.

Results

Table 1: Observation of Acids and Bases

Trial

Concentration of Solution Used

 Acid            Base

ΔT from graph

(°C)

Moles of H2O Formed (nwater)

Number of kJ/mole of H2O formed

A

1

2.0372M

2.0343M

11.5

0.081mol

53KJ/mol

2

2.0372M

2.0343M

11.5

0.081mol

53KJ/mol

B

3

1.7377M

2.0343M

10

0.070mol

54KJ/mol

4

1.7377M

2.0343M

9.9

0.070mol

54KJ/mol

C

5

0.5000M

2.0343M

2.8

0.025mol

47KJ/mol

6

0.5000M

2.0343M

2.4

0.025mol

40KJ/mol

D

7

2.0343M

12.6

8

2.0343M

12.7

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Discussion

Since the number of KJ/mol of water formed will be the same as the KJ/mol as the electrolyte. Seeing as these acids or electrolytes should all be strong, except for phenol, these values are fairly accurate. This is known because the correct because the amount of heat that is given off, theoretically by strong electrolytes is 55.90KJ/mol. The values that were found are all technically negative values because the process is exothermic, meaning that the heat moves from the system, neutralization, to the surroundings.

 The calculations show that the neutralization reaction with NaOH and HNO3 produces an average molar enthalpy of 54 kJ/mol. Both HNO3 and NaOH are strong electrolytes and therefore both completely ionize in water. Of the four parts conducted in this experiment, this was the closest value to the number of kilojoules per mole of water formed.

To solve for the molarity of the unknown hydrochloric acid, the average of the heat of reactions from both trials of part A were used. By finding the moles of water (which is equal to the amount of moles of HCl) the molar enthalpy was determined using q = mcΔt. By taking the average molar enthalpy, -53KJ/mol, the moles of the unknown HCl can be determined. The average morality of HCl unknown #5 that was found was 2.24M. The calculations are shown in the results section as part D.

With an average molar enthalpy of 54 kJ/mol, the neutralization reaction between HNO3 and NaOH produced the most heat per mole of H2O formed. This reemphasizes on the point that strong electrolytes dissociate completely as well as releases more heat for the reaction. If compared to part c, neutralization reaction between phenol and sodium hydroxide (a weak electrolyte), it is apparent that the reaction with the strong electrolytes produces more heat; average heat of reaction for part C equalled 40 kJ/mol versus the heat of reaction for part B which was 54 kJ/mol. The solutions in part C do not completely ionize which is a result of the lower molar enthalpy. These results should have shown more of a difference in heat, however there are many sources of error so these results are slightly off.

Conclusion

In this experiment, the strong acids that were used were HCl and HNO3 – strong electrolytes. The weak acid that was used was called phenol – weak electrolytes. The base that was used for all 4 neutralization reactions was a strong base called sodium hydroxide which also happens to be a strong electrolyte.

        Because there were so many trials and required readings, there was plenty of room for error in this lab. Some sources of error include reading errors such as reading the thermometer incorrectly and not keeping the readings constant (i.e. every 10 seconds, every 5 seconds), keeping time inconsistently. When placing the thermometer in the calorimeter, it was suggested not to allow the thermometer to touch the bottom of the coffee cup; a few times, because this was so tedious, the thermometer did touch the bottom and that may have affected the results as well. Stirring the solution aids the reaction, but some parts of the experiment may have been stirred more rapidly than others which may have affected the results. Finally, at the beginning of each trial it was recommended to take a reading of the temperature every second, however it was difficult to do this and the temperatures may be off. This lab displayed the functions of the calorimeter accurately due to the fact that the calorimeters are not perfect and do let some heat escape.


References

Department of Chemistry 2008 First Year Chemistry Chem 120L Laboratory Manual. University

of Waterloo, Waterloo. pp 16-20.

Petrucci, Ralph et.al. General Chemistry Principles & Modern Applications, Ninth Edition.

        Upper Saddle River: Pearson Education, 2007.

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