Calorimetry Lab

The boiling point elevation constant for water that was experimentally determined in this analysis was 0.4396 °C/m, which was derived from the slope of the trend line in Figure 2. This is slightly lower than the constant provided in lecture of 0.51 °C/m. This could be due to further evaporation of water from the solutions tested via refractive index after the boiling temperature was recorded.

The next step in this lab is to rinse the Erlenmeyer flask with distilled water down the drain and then repeat the experiment, this time adding 10 ml of 0.10M KI and 10 ml of distilled water to the flask instead. The flask should again be swirling to allow the solution to succumb to the same temperature as the water bath and once it has reached the same temperature, 10 ml of 3% H2O2 must then be added and a stopper must be immediately placed on the flask and recording should then begin for experiment two. After recording the times, the Erlenmeyer flask must then be rinsed again with distilled water down the drain. After rinsing the flask, the last part of the lab can now be performed. Experiment three is performed the same way, but instead, 20 ml of 0.10 ml M KI and 5 ml of distilled water will be added and after the swirling of the flask, 5 ml of 3% H2O2 will be added. After the times have been recorded, data collection should now be complete.

The main purpose of the lab “Determination of the Formula of a Copper Oxide” was to determine the formula of a copper oxide. Specifically, this is a compound of copper combined with oxygen. This was to be done by heating the copper oxide thoroughly until all of the oxygen had been driven off. To accomplish this experiment, we first had to take and measure the mass of a specified color of copper oxide, ours being red. Then, we used a fischer burner to provide the heat needed for the split of copper oxide, in which our amount resided in a test tube. But, in order for the copper to not recombine with oxygen that could be found in the surrounding atmosphere of our lab, we also had to have a flow of methane gas into the test tube that fed into

The initial circumference of the balloons were 41.5cm. The temperature of the hot water was 58℃, the room temperature water was 25℃ and the cold water was -4℃. After the balloons were placed in the water for 3 minutes, we got our results. The final circumference of the balloon in the hot water was 40.5. It shrank 1cm less which is a surprise since it should expand.

Our task was to identify the powder in each bag containing different amounts of moles. The mole provides a standard unit of measure that can be used to compare a wide variety of substances. A mole of atoms gives you a physical representation of what a single atom. The molar mass is the total mass of an element divided by one mole in that element. The molar mass can be used as a physical property to identify unknowns by converting it to moles and seeing which one is the closest to the moles in the bag. Our guiding question to answer was: “What are the identities of the unknown compounds?”

The experiment was started preparing 300 mL of a 2 M HCl solution. A graduated cylinder was used to collect 200 ml of deionized H2O, then it was added to an empty 600 mL beaker which was designated to be the 2 M HCl solution. 100 mL of 6M HCl was then added to a sanitized graduated cylinder and poured into the 600 mL beaker with H2O. The solution was then stirred with the glass rod. 150 mL of a 2 M NaOH solution was then prepared. 50 mL of deionized H2O was added to a 400 mL beaker. Then, a graduated cylinder of 100 mL 3 M NaOH was added to the beaker. Repeat stirring. LabQuest was then configured and setup for data collection of Part A; the interval box should be set to 15 seconds.

Procedure: In preparation 600mL of water was added to a beaker then an Erlenmeyer flask was placed inside to determine the water level when the Erlenmeyer flask with unknown was added. This was done to prevent water from spilling over the edge while also

Next, one of the 5 mL syringes was used to collect 1.5 mL of tap water. Before filling, the cap was removed from all the syringes. The water was checked for air bubbles. Throughout the experiment, if the fluid in the syringe had air bubbles it was emptied and filled until air bubbles were no longer present. The 1.5 mL of tap water was emptied into well #2. The same syringe was then filled with 2.0 mL of tap water. This water was then placed in well #3. Well #1 was left without H2O. The syringe was finally placed in the H2O beaker in order to keep the syringes

The red product was filtered off by using a vacuum filtration flask and funnel. Air was sucked through the solid to help to dry it. The crystals were transferred to a preweighed vial and were weighed on a top loading balance. The weight was recorded the vial was labelled with the name, the experiment number and the formula of the compound. About 1 g of the product was dissolved in 5 mL of conc. HCl in a small beaker. The solution was diluted to about 25 mL by using distilled water. The solution was warmed for about 5 minutes. This solution was called solution A. 5 mL of solution A were diluted to 50 mL with distilled water. 6 M NaOH solution was added slowly until it became basic and its solubility and colour were noted. A fresh precipitate of Fe(OH)3 was prepared from a separate 5 mL of solution A. an excess of conc. HCl was added to the precipitate and its solubility was noted. 2-3 drops of solution A were taken by using a dropping pipette, and were diluted to about 150 mL with distilled water in a clean 250 mL beaker. 5 ml of 1 M potassium thiocyanate (KSCN) were added to the diluted solution, and the colour produced was

Figure 1 shows that the steepest line is the 10 minutes’ line. Since the steeper the line , the faster the rate of change, one can conclude that in the first ten minute interval, the rate of change was the fastest , which means that in those 10 minutes Osmosis was the fastest. The reason for this could be that during osmosis water is diffused through a selectively permeable membrane, from a region od higher water potential to a region of lower water potential. So once the dialysis bags were immersed in the water, water entered the dialysis bag faster than any other time interval.

The first sample was placed in a water bath of 38⁰ C. An ice bath was used to create a sample of water at 3.5⁰ C and the third sample was maintained at a constant 18.5⁰ C which was room temperature. Dialysis bags were then attached to glass tubing using string in order to create an osmometer. All three osmometers were then filled with the same 20% sucrose solution. A thermometer was placed into all three water samples to ensure the temperatures remained constant throughout the experiment. A stopwatch was used to time each osmometer in each water sample for a total of twenty minutes and all the data was recorded. The increase in amount of water in each osmometer was recorded each minute using a ruler (millimeters). At the end of the twenty minutes all the data had been

The purpose of this lab is to find most efficient way to capture energy from a combusted chip using a calorimeter. We made a basic calorimeter by using a steel can. Throughout our lab we made modifications to the calorimeter to increase the percent of energy captured by the water.

The procedure was repeated with the different amounts of water and the different chemicals indicated

The reaction occurred under three different temperatures: 40ْ C, 10ْ C, and 0ْ C. The experiment in this part was carried out just like part one. In flask 1, 10 mL of 0.010 M KI, 10 mL of 0.0010 M Na2S2O3, and 10 mL of H2O were mixed. In flask 2, 10 mL of 0.040 M KBrO3 and 10 mL of 0.10 M HCl were mixed, and 3 drops of starch was added. Then, the two flasks were put in an ice bath and cooled to about 10ْ C. Afterwards, the two solutions were mixed together while in the ice bath with swirling until turning blue. The time was recorded. The reaction was done the same way for the other two degrees (0ْ C, 10ْ C). If the water needed to be cold, ice was added, and if hot water was needed, the water was heated.

In this experiment, hot water was mixed with room temperature water in a calorimeter to calculate the specific heat capacity of the calorimeter. The specific heat of the calorimeter was then used to find the enthalpy of sodium hydroxide and hydrochloric acid, ammonia and hydrochloric acid, and ammonium chloride and sodium hydroxide by mixing each, respectively, into the beaker after each other. The beaker was washed with distilled water after each reaction. The enthalpy of sodium hydroxide and hydrochloric acid and ammonia and hydrochloric acid were subtracted together to calculate the enthalpy of ammonium chloride and sodium hydroxide to prove Hess’ Law.