Lab 4 Heat Reactions

Lab 4: Heat Reactions Abstract: The purpose of this lab is to use calorimetry to measure the heats of reaction, while also attempting to use the standard formation heats. It was also to become familiar with the heat transfer involved in a solution and solvent, while also attempting to understand the heat capacity of water and how everything plays in together. When heat is absorbed in the final temperature, it means that the change is endothermic. When heat is released in the final temperature, it means that the change is exothermic. For the heat of solution with the water and potassium nitrate, it was found that heat was released, hence making it an exothermic reaction. For the heat of neutralization with the hydrochloric acid and sodium hydroxide solution, it was found that heat was absorbed, hence making it an endothermic reaction. For the magnesium strip and hydrochloric acid, it was found that heat was also absorbed, hence making it an endothermic reaction. There was a large percent error for part 1, which was 82.66%. The percent error for part 2 was around 1.37%, and the percent error for part 3 was 16.07%. The difference in temperature for part 1 was -1.7˚C. The difference in temperature for part 2 was 5.4˚C, and the difference in temperature for part 3 was 15.4˚C. Introduction: When the symbol q is given, the physical and chemical changes that are involved include either the absorption or release of heat. When heat is absorbed, it means that the change is endothermic and is given a positive sign. When the heat is released, it means that the change is exothermic and is given a negative sign. So for endothermic, it’s (q > 0), and for exothermic, it’s (q < 0). When there is constant pressure, the heat flow is equal to the total heat. Under certain conditions, the heat of reactions can be measured by a calorimeter. The most useful and widely used heats of reactions to measure is the standard heat flow of formation of a substance, ΔHf˚. This is what’s involved in the forming of one mole of a substance under certain conditions. For example, for H2O the ΔHf˚ is given by the thermochemical reaction: H2 (g) + ½O2 (g)  H2O (l) ΔHf˚ = -285.83 kJ/mol This reaction explains that when one mole of liquid water is formed, it releases 285.83 kJ exothermically considering it’s a negative number. The reason for using the heat flow of formations is so that all data can be used to calculate the heat flow of any reaction under certain conditions, given that we know the ΔHf˚ values for both the reactants and products. For any reaction under the certain conditions, the enthalpy, ΔHf˚, of the reaction is the sum of the enthalpies of the products multiplied by their stoichiometric coefficients in the reaction minus the sum of the formation of the reactants multiplied by their stochiometric coefficients in the reaction. In this experiment, a “coffee-cup calorimeter” was used, which is two polystyrene cups together (one inside of the other), which are then balanced with a ring on a ring stand and covered with a 10cm piece of square fiber board with a small hole on top in which a thermometer is inserted. A stir bar that’s placed inside of the cup helps to continuously stir the solution. Regardless of the design and which calorimeter is used, it’s purpose is universal—to help determine the heats of reaction by calculating the heat change that’s experiences by the calorimeter itself and not the reaction. This is usually found by the equation: qcal = Ccal ΔT qcal is the heat change for the calorimeter, and the Ccal is the heat capacity of the calorimeter. The ΔT represents the temperature change, which is the final temperature minus the initial temperature. In this experiment, dilute solutions were created, where the heat capacities aren’t significantly different from pure water, which has a specific heat of 4.184 J/˚C*g. Due to this, the heat capacity of the calorimeter and the specific heat of water is calculated using the equation: Ccal = (grams of solution) * 4.184 J/(˚C*g) If the temperature of the liquid in the coffee-cup calorimeter rose as a result of the chemical reaction or physical change, the calorimeter absorbed heat which results in an endothermic heat change. If the temperature of the liquid in the coffee-cup calorimeter decreased as a result of the chemical reaction or physical change, the calorimeter lost heat which results in an exothermic heat change. The purpose of the experiment was to find out whether each of trials entailed an endothermic or exothermic reaction, while subsequently using temperature change in order to determine what each experiment entailed. Method: In part one, the mass of both polystyrene cups with the stir bar placed inside it were measured on an analytic balance and was recorded to the nearest ±0.01 g. 100 mL of deionized water was measured with a graduated cylinder and poured into the measured cups. The mass of everything was then measured again on an analytic balance and was recorded to the nearest ±0.01 g. A single aluminum weighing dish was then weighed on an analytic balance, and then potassium nitrate was added, which had a mass between 2.0 to 2.5 g. The mass of the weighing dish and potassium nitrate was recorded to the nearest ±0.01 g. The coffee-cup calorimeter was made, and the motor for stirring was turned on. The speed was adjusted to make sure there was adequate stirring but no splash-back. The temperature of the water in the cup was recorded just before the potassium nitrate was added. The mixture was then added and covered immediately with the lid and was stirred for approximately one minute. The temperature was then recorded once all the solid potassium nitrate was dissolved. In part two, approximately 50 mL of 1.00 M hydrochloric acid was measured, recorded, and poured into the polystyrene cup. With a separate graduated cylinder, between 30 and 40 mL of 1.00 M sodium hydroxide solution was measured and recorded. The coffee-cup calorimeter was made, and the motor for stirring was turned on—ensuring that the speed was adjusted to make sure there was adequate stirring but no splash-back. The temperature of the hydrochloric acid was measured and recorded before the sodium hydroxide solution was added. All of the sodium hydroxide solution was added to the cup that contained the hydrochloric acid solution, and the lid was immediately placed over the cup. The mixture was stirred for approximately one minute until the maximum or minimum temperature was achieved. The final temperature was then recorded. In part three, the polystyrene cup was measured with the stir bar inside of it on an analytic balance, and the mass was recorded to the nearest ±0.01 g. Using a graduated cylinder, approximately 75 mL of 1.00 M of hydrochloric acid solution was measured and poured into the cup. Everything was then measured on the analytic balance and the mass was recorded to the nearest ±0.01 g. A 15 cm strip of magnesium ribbon that weighed about 0.1 g was sanded down with sandpaper until was shiny. The strip was shaped into a zig-zag pattern, and the mass was determined using an analytic balance and recorded to the nearest ±0.0001 g. The coffee-cup calorimeter was made, and the motor for stirring was turned on—ensuring that the speed was adjusted to make sure there was adequate stirring but no splash-back. The temperature of the hydrochloric acid was recorded before the magnesium strip was added to the cup. The lid was then immediately placed over the top. The mixture was stirred for approximately one minute until the maximum or minimum temperature was achieved. The final temperature was then recorded. Data: Table 1 shows the heat of solution for the water and potassium nitrate. Each of the data was either measured in grams, J/˚C, moles, or kJ/mol. Table 1. This table shows the heat of solution for the water and potassium nitrate with the proper calculations, and measurements. Mass of empty cups with stir bar Mass of cups, stir bar, and 100 mL water Mass of water Mass of weighing dish Mass of weighing dish and KNO3 (s) Mass of KNO3 Mass of solution Heat capacity of calorimeter Initial Temperature (Ti) Final Temperature (Tf) Temperature change (ΔT) Calorimeters heat change (qcal) Solution heat change (qsoln = -qcal) Moles of KNO3 (s) dissolved Heat of solution per mole of KNO3 (s) dissolved Exothermic/Endothermic Percent error, given value +34.89 kJ/mol 7.5235 g 104.9720 g 974485 g 0.9351 g 2.0516 g 1.1165 g 9.565 g 412.3956 J/˚C 20.5 ˚C 18.8 ˚C -1.7 ˚C -701.073 J 701.073 J 0.011 mol 63.733 kJ/mol Exothermic 82.66% Table 2 shows the heat of neutralization data for the hydrochloric acid and sodium hydroxide solution. Each of the data was either measured in mL, g, J/˚C, ˚C, J, moles, or kJ/mol. Table 2. This table shows the heat of neutralization data for the hydrochloric acid and sodium hydroxide solution with the proper calculations, and measurements. Exact volume of 1.00 M HCl (aq) Exact volume of 1.00 M NaOH (aq) Moles of NaOH Mass of solution Heat capacity of calorimeter, Ccal Initial Temperature (Ti) Final Temperature (Tf) Temperature change (ΔT) Calorimeters heat change (qcal) Neutralization heat change, qrxn = -qcal Moles of H3O+ neutralized with OHHeat of neutralization per mole of H3O+ neutralized Exothermic/Endothermic Percent error, given value -57.36 kJ/mol 50.2 mL 31.9 mL 0.0319 mol 82.1 g 343.506 J/˚C 20.7 ˚C 26.1 ˚C 5.4 ˚C 1854.93 J 1854.93 J 0.0319 mol 58.14 kJ/mol Exothermic 1.37% Table 3 shows the heat of reaction of the magnesium and hydrochloric acid. Each of the data was measured in either g, J/˚C, ˚C, J, moles, or kJ/mol. Table 3. This table shows the heat of reaction of the magnesium and hydrochloric acid with the proper calculations and measurements. Mass of empty cups with stir bar Mass of cups, stir bar, and 75 mL of 1.00 M HCl (aq) Mass of HCl (aq) solution Mass of Mg (s) Mass of solution Heat capacity of calorimeter, Ccal Initial Temperature (Ti) Final Temperature (Tf) Temperature change (ΔT) Calorimeters heat change (qcal) Reaction heat change, qrxn = -qcal Moles of Mg (s) reacted Heat of reaction per mole of Mg (s) reacted Heat of reaction using HCl (g) and MgCl2 (aq) Percent error, given value 981.5 kJ/mol 7.5365 g 81.5517 g 74.0152 g 0.5462 g 74.2614 g 310.709 J/˚C 20.5 ˚C 35.9 ˚C 15.4 ˚C 4784.929 J -4784.929 J 0.0042 mol 1139.268 kJ/mol 981.5 kJ/mol 16.07% Calculations: The calculation for the literature value for the heat of reaction using HCl (g) and MgCl2 (aq) reaction is shown below. ΔHrxn = ∑MgCl2 – ∑HCl ΔHrxn = (-796.9) – (-184.6) ΔHrxn = 981.5 kJ/mol The calculation for the calculation for the reaction heat change using HCl (g) and MgCl2 (aq) reaction is shown below. Ccal = mass solution * 4.184 (cup of water) * Ccal = 74.2614 g * 4.184 Ccal = 310.709 J/˚C qcal = Ccal ΔT qcal = (310.709 J/˚C) * (15.4˚C) qcal = 4784.929 qrxn = -qcal qrxn = -4784.929 Discussion: For the heat of solution with the water and potassium nitrate, it was found that there was temperature lost which meant that there was an exothermic reaction. Once the potassium nitrate was mixed with the water, it was proven that once potassium nitrate dissolves, the temperature drops a little bit. There is an exothermic reaction being produced because the heat was released within the cup once the stirring occurred. The heat of solution in this first part means that heat is either absorbed or released when one mole of the solute is dissolved in an amount of water. This amount of water should cause no significant change in the temperature once the substance is being dissolved. Due to this, the result of the experiment is more likely to be exothermic because the initial measured temperature of the solution was going to be less than the actual value. The percent error for this part ended up being 82.66%. For the heat of neutralization with the hydrochloric acid and sodium hydroxide solution, it was found that there was temperature gained (meaning heat was absorbed), hence making it an exothermic reaction. Once the sodium hydroxide was added to the hydrochloric acid, they both dissolved together to form water. This is due to the fact that a strong base and strong acid are added together, and the neutralization reaction occurs. The reaction of both of these substances was the temperature change observed when one mole of the hydroxide ion is added to and reacts with a hydronium ion. Due to this, heat is usually absorbed. The percent error for this part ended up being 1.37%. For the reaction of the magnesium and hydrochloric acid, it was found that there was temperature gained (meaning heat was absorbed), hence making it an endothermic reaction. Since magnesium was mixed with hydrochloric acid, there was hydrogen gas and magnesium chloride that was formed. Due to the hydrogen gas being formed, heat was absorbed in order to make that reaction happen which led to the temperature to increase significantly. The percent error for this part ended up being 16.07%. For the heat of solution with the water and potassium nitrate, the “expected” temperature change would be that it would be exothermic. Once the potassium nitrate is added to the water, it should cause no significant change in the temperature—hence it should be less than what it started with. Since there was temperature lost, the percent error is going to be large because the temperature dropped, and heat was lost. The calculation for delta T is shown below. ΔT =final temperature ( Tf )−initial temperature (Ti) Δ T =18.8 −20.5 ΔT =−1.7 For the heat of neutralization with the hydrochloric acid and sodium hydroxide solution, the expected temperature change would be that heat would be lost, and the temperature would also subsequently drop due to a strong base and a strong acid mixing together. However, the temperature went up, meaning heat was absorbed. Since there was heat that was absorbed, the percent error is going to be small because the temperature went up. The calculation for delta T is shown below. ΔT =final temperature ( Tf )−initial temperature (Ti) ΔT =26.1−20.7 ΔT =5.4 For the reaction of the magnesium and hydrochloric acid, the expected temperature change would be that it would increase causing heat to be absorbed. This would be due to the fact that a strong acid is mixed with a weak base. In turn, the percent error would also be small because the temperature went up and heat was absorbed. The calculation for delta T is shown below. ΔT =final temperature ( Tf )−initial temperature (Ti) ΔT =35.9−20.5 ΔT =15.4 For the entirety of all of the experiments, there were a few possible sources of error. There could have been too much or too little potassium nitrate that was added to the water, too much or little of the hydrochloric acid or sodium hydroxide was used, or that the magnesium strip wasn’t sanded down enough. It was also possible that the crucible lid wasn’t placed properly or quickly enough in order to not let any outside oxygen in. Conclusion: In this experiment, there was an attempt to use a calorimeter to try and determine if the heat of the neutralization was endothermic or exothermic using three different trials—water and potassium nitrate, hydrochloric acid and sodium hydroxide, and a magnesium strip and hydrochloric acid. The calorimeter was used to measure the heats of reaction by measuring the different substances and then calculating how much of the substance was neutralized. The calculations for all of the parts were the same, with just a couple of different formulas. The neutralization of KNO3 turned out to be exothermic. The neutralization of NaOH and the magnesium strip turned out the both be endothermic. There was a large percent error for part 1 (82.66%), while it was a little less (1.37%) and (16.07%) for parts 2 and 3 respectively.