Introduction
The first heat packs are believed to date back to 12th century China, during the winter locals would carry around hot rocks to use as portable “kairos”. Then, in the early 1900s, people began to use smoldering dried peat moss. The burning heat source was usually placed in a portable tin to keep in pockets. This report examines the chemical composition of a modern heat pack. Particularly, what is the most desirable configuration to prevent a human hand from getting frostbite. Although most modern heat packs use salt, iron powder, and water, it is still possible to create a heat pack with only a chemical salt and water.
When a salt is dissolved in water, the reaction that occurs is usually endo or exo thermic. This experiment focused …show more content…
Part A
It was important to test each of the four samples given for the largest increase in temperature. To begin, 20ml of purified water was poured into a beaker. Its temperature was then recorded in real time as 1.5g of the chosen salt was added. The computer monitored the process and returned the maximum observed temperature. The total change on temperature was compiled into a table similar to that of table 1. It was essential to repeat the process with each salt, taking care to note the sign of Δ T, half of the salts produced endothermic reactions.
With the table in part a filled out, the salt with the most positive temperature gain was the most exothermic. This salt was used to complete part B.
Part B
In part B, the salt/ water mixture was optimized to produce the correct amount of heat. After adding 70mL of tap water to the calorimeter, the temperature probe was inserted to record. After around 1 minute of recording, 9.8g of CaCl was added. The mixture was then agitated with a stirring rod as the salt dissolved, temperature increased sharply until it reached a plateau. The highest temperature was compared with the initial to find the difference. Table B included this and other relevant information. Heat produced was calculated q=m Δ T, and Δ H was calculated by dividing the heat …show more content…
While table 1 suggests the use of MgS04, the correct salt to use for this application was surprisingly CaCl2. The constant mass of salt, 9.8g for each run, was essential. It was also crucial to only vary volume of water to receive a good temperature change approximation. The mass of the solution was easily calculated by adding the weight of water and salt together. Change in temperature was found by subtracting the final and initial temps during dissolution. Trial two had a more noticeable temperature change, yet still had a lower amount of heat produced as given by the equation q=mcΔT. ΔH was simply our Heat produced per gram of salt. As expected, the Δ H of around -60kJ/mol is about 30% shy of the true value of around -80kJ/mol. For enhanced accuracy, the remaining calculations relied on the exact -83kJ/mol value for Δ
The first heat packs are believed to date back to 12th century China, during the winter locals would carry around hot rocks to use as portable “kairos”. Then, in the early 1900s, people began to use smoldering dried peat moss. The burning heat source was usually placed in a portable tin to keep in pockets. This report examines the chemical composition of a modern heat pack. Particularly, what is the most desirable configuration to prevent a human hand from getting frostbite. Although most modern heat packs use salt, iron powder, and water, it is still possible to create a heat pack with only a chemical salt and water.
When a salt is dissolved in water, the reaction that occurs is usually endo or exo thermic. This experiment focused …show more content…
Part A
It was important to test each of the four samples given for the largest increase in temperature. To begin, 20ml of purified water was poured into a beaker. Its temperature was then recorded in real time as 1.5g of the chosen salt was added. The computer monitored the process and returned the maximum observed temperature. The total change on temperature was compiled into a table similar to that of table 1. It was essential to repeat the process with each salt, taking care to note the sign of Δ T, half of the salts produced endothermic reactions.
With the table in part a filled out, the salt with the most positive temperature gain was the most exothermic. This salt was used to complete part B.
Part B
In part B, the salt/ water mixture was optimized to produce the correct amount of heat. After adding 70mL of tap water to the calorimeter, the temperature probe was inserted to record. After around 1 minute of recording, 9.8g of CaCl was added. The mixture was then agitated with a stirring rod as the salt dissolved, temperature increased sharply until it reached a plateau. The highest temperature was compared with the initial to find the difference. Table B included this and other relevant information. Heat produced was calculated q=m Δ T, and Δ H was calculated by dividing the heat …show more content…
While table 1 suggests the use of MgS04, the correct salt to use for this application was surprisingly CaCl2. The constant mass of salt, 9.8g for each run, was essential. It was also crucial to only vary volume of water to receive a good temperature change approximation. The mass of the solution was easily calculated by adding the weight of water and salt together. Change in temperature was found by subtracting the final and initial temps during dissolution. Trial two had a more noticeable temperature change, yet still had a lower amount of heat produced as given by the equation q=mcΔT. ΔH was simply our Heat produced per gram of salt. As expected, the Δ H of around -60kJ/mol is about 30% shy of the true value of around -80kJ/mol. For enhanced accuracy, the remaining calculations relied on the exact -83kJ/mol value for Δ