Calorimetry. Enthalpy of Neutralization

Transcription

1 Calorimetry Enthalpy of Neutralization

2 Introduction A calorimeter is a device that can measure the heat absorbed or released by a reaction (Petrucci, 2011). A calorimeter is thermally insulated from its surroundings, so that no heat from the reaction is lost to them; this means it is an isolated system (Petrucci, 2011). There are two main types of calorimeters: a bomb calorimeter, and a coffee cup calorimeter. In this experiment a coffee cup calorimeter was used. This is a simple calorimeter in which the reactants are mixed in a double Styrofoam cup. Since Styrofoam is such a good insulator, very little heat escapes from the system into the surroundings. So we treat the coffee cup calorimeter as an isolated system (book). The flow of heat is calculated using this equation: q=mc T (Petrucci, 2011). In this equation q is the heat transferred, c is the specific heat capacity, m is the mass, and T is temperature (Petrucci, 2011). Since the coffee cup calorimeter is not sealed, the external pressure will equal the pressure inside the calorimeter. This means that the initial q value, or the heat of reaction, will be equal to the change in enthalpy, ΔH. A neutralization reaction is when an acid and a base react to form a salt and water (Negi & Anand, 2004). The heat of reaction is a result of the formation of water from the hydronium ions and the hydroxide ions (Negi & Anand, 2004). As long as the products and reactants of the neutralization reaction are all ionized, they will not directly affect the heat of neutralization. Neutralization reactions are generally exothermic. An exothermic reaction is one that produces heat, and as a result would cause an increase in temperature in an isolated system, such as a coffee cup calorimeter (Petrucci, 2011). The heat, or enthalpy, given off by a neutralization reaction is the combination of heat absorbed by the solution and the calorimeter (book). All neutralization reactions

3 3 involving strong acids and strong bases will produce approximately the same amount of heat: kj (Petrucci, 2011). Weak electrolytes only partially dissociate into ions (Negi & Anand, 2004). Since this is true, these reactions will have other enthalpy terms (heat of dissociation) as well as the heat per mol of H + (which is kj) (Petrucci, 2011). The different types of weak acids or bases will determine whether or not the reaction is endothermic or exothermic, which will determine whether or not the heat of reaction is smaller or larger than that of a strong acid-base reaction (Negi & Anand, 2004). Thus, when finding heat of reaction, one will be able to note whether or not the reaction is with strong or weak acids. The purpose of this lab is to learn the methods of using a calorimeter, and learning to calculate the heat and enthalpy of neutralization using the laws of Thermodynamics. Experimental Procedure The experimental procedure used for this experiment was outlined in the CHEM 120L lab manual, Experiment 4. All the steps were followed without deviation. Experimental Observations Table 1: Molarities Compound NaOH HCl HNO3 Phenol Molarity (mol/l) M M M M Unknown #2

4 Table 1: This table shows the molarities of the different compounds used in this experiment. Table 2: Time and Temperature of Neutralization Reactions Part A Trial 1 Part A Trial 2 Part B Trial 1 Part B Trial 2 Part C Trial 1 Part C Trial 2 Time Temp. Time Temp. Tim e Temp Time Temp Time Tem p Time Tem p

5 Table 2: This table shows the temperature (Temp), in C, and the time, in seconds, to which it corresponds for the neutralization reaction in each parts A, B, and C.

6 Table 3: Time and Temperature of the Unknown Reaction Part D: Trial One Part D: Trial Two Time (s) Temperature ( C) Time (s) Temperature ( C)

7 Table 3: This table shows the temperature, and the time to which it corresponds for the neutralization reaction in Part D.

8 Results and Conclusions Table 4: Averages Part A-D Part Trial Base (M) (NaOH) Acid (M) ΔT C Moles H 2O q (J) q/mol A mol/ HCl L 1.885mol/L Average part A: B mol/ HNO L 1.984mol/L Average part B: C mol/ Phenol L mol/ L Average part C: D mol/ # L mol/l Average part D: Table 4: This table shows all findings from all 8 reaction trials, with the averages for each part. Sample Calculations: Numbers used from Part A, Trial One Average Temperature

9 9 Moles H 2 O q (kj) q /mol Part D

10 Part A Graph 1: Trial One Graph 2: Trial 2 Part B Graph 3: Trial One Graph 4: Trial Two Part C Graph 5: Trial One Graph 6: Trial Two Part D Graph 7: Trial One Graph 8: Trial Two Graphs 1-8 represent the temperature change over time of the various neutralization reactions. Discussion The values obtained matched fairly well when compared with those in literature. The accepted value for any neutralization reaction with strong acids and strong bases is kj (Petrucci, 2011). For Part A, the average enthalpy per mole H 2 O was kj (since q=56.4 kj, and ΔH=-q). Since HCl is a strong acid, and NaOH is a strong base, when mixed these substances will yield water and NaCl (Negi & Anand, 2004). So for Part A, the difference between the accepted value and the experimental value is only 0.83, which is only a 1.49% error. In Part B, NaOH is a strong base, and HNO3 a strong acid (Petrucci, 2011). Therefore the experimental value should be close to the accepted

11 11 value of kj. Since the experimental value is kj per mole H 2 O, there is a 3.96% error, which is still quite close. In Part C, NaOH is still a strong base, but since phenol is not on the table of strong acids, it is assumed to be a weak acid (book). Therefore, depending on whether it is exo- or endothermic it will have a value that is more or less than kj (Petrucci, 2011). The experimental value calculated is kj per mole H 2 O. This value is the furthest off from the accepted value of 25.3 kj/mol of phenol (internet book). This gives a 17.40% error, which is quite significant. The error in these values could be due to several different things. Human error in reading the temperature could be a large part of why the values are too big or too small. Another human error would be not stirring the reaction consistently, which would lead to different temperature changes. Also, the coffee cup calorimeter was not actually sealed, so some heat probably escaped through the top of the calorimeter. Since we made the assumption that no energy escaped the system, this will throw off our values slightly. We also made the assumption that the density and heat capacity of the acids and bases will be the same as water. This however, will not be true since NaOH and the acids used will not have the exact density of water. However, the densities are very similar since the concentrations of the acids and bases are so low, so these assumptions will only very slightly alter the experimental results, if at all. We also made the assumption that the initial temperature would be the same as the average of the initial base and acid temperature. This is a valid assumption, since they were only 0.1 degrees apart in the first place, so this number would be considerably accurate. The assumption that the rate of heat lost to the surroundings being linear was made. This was used to extrapolate the peak temperature. This is a fairly accurate assumption, as heat lost over time would

12 represent a linear graph. The assumption that no heat would be absorbed by the calorimeter was also made. The experimental result would have been more accurate if the specific heat capacity of the calorimeter was accounted for. However, it would have been a very small amount of heat lost; therefore this was a fairly accurate assumption to make. The unknown #2, which is an HCl acid of unknown concentration, was found to have a molarity of mol/l. The difference in temperature was the same for both trials in Part D, so that part of the experiment was probably done very accurately. The average q value used from Part A to find molarity for part D carried its error with it, though, since it was not completely accurate. Therefore, this result would have at least the error in Part A (1.49%) plus any human error that could have occurred during Part D. However, this is most likely a fairly accurate result, as the temperature difference was exactly the same, and the average for Part A only had a percent error of 1.49%.

13 13 Conclusion The purpose to learn how to use a calorimeter was accomplished, as well as how to calculate heat and enthalpy of neutralization, using q = mcδt. The ability to read a thermometer accurately was acquired as well. The laws of Thermodynamics were observed in this lab, and calculated to find experimental values and compare them to accepted values. The assumptions made and the human error led the experimental data to be slightly off, thus not proving the laws of Thermodynamics completely. However, error was, for the most part, quite small. Therefore the purpose of this lab was accomplished. The unknown concentration of HCl acid #2 was found to be mol/l, using the enthalpy of neutralization from Part A and the assumption that the specific heat capacity and density is the same as water. This result is fairly accurate. The error in this lab was due to somewhat faulty assumptions as well as human error. The base and acids used in this lab did not have density the exact same as that of water, the coffee cup calorimeter is not completely sealed (and thus lets some heat escape the system), and it absorbs some of the heat of the reaction. The thermometer could have been read slightly wrong, and the reaction stirred inconsistently. Therefore, there was error in the lab, and this was significant enough to produce experimental values slightly off from the accepted, but only to a small degree. The errors in most parts of the lab were quite small. Therefore, with more acquired skills, and with fewer assumptions, this lab would have produced more accurate results. However, this lab accomplished its purpose.

14 References Negi, A. S., Anand S.C. (2004). A Textbook of Physical Chemistry. New Delhi: New Age International Limited Publishing. Petrucci, et. al. (2001). In Bennet G. (Ed.), General chemistry principles and modern applications (10th ed.). Toronto, ON: MacMillan Publishing Company. University of Waterloo. Department of Science. CHEM 120 Courseware Fall Waterloo, Fall 2012.U University of Waterloo. Department of Science. CHEM 120L Laboratory Manual. Waterloo, Fall 2012.