Table kPA and 298 K), should boil at

2Background p. 2Method p. 4Raw Data p. 5Processed Data p. 7Conclusion p. 10Evaluation p. 11Risk Assessment p.

12Bibliography p.12Introduction: Distilled water, under standard conditions (101 kPA and 298 K), should boil at exactly 373 K (or 100 °C). However, when one adds different substances (solutes) to that amount of water, then the boiling point increases. This change in temperature (?T) is given by the equation: ?T =Kb *m*i (where Kb is molal boiling point constant, m is molal of the solute, and i is the number of ions that the solute breaks into when mixed with water). This is called the boiling point elevation equation. This investigation will be centered around proving that the number of ions released by a compound (solute) in solution (water is the solvent) is directly proportional to the increase in boiling point of that solution.

Background: I love to ski, and when I go skiing the road is often iced over. To ensure that cars do not slide and slip on these ice sheets, people pour ethylene glycol. This lowers the freezing point depression and essentially unices the roads. Now, bringing solutions to under freezing in a high school laboratory would be extremely difficult. So I decided instead to investigate the elevation boiling point. During my research I learnt that in order to boil water, enough energy must be put in to overcome the intermolecular (H-bonding and London Dispersion Forces (LDFs)) forces in the water so that they are able to overcome the vapor pressure on the surface of the liquid.

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However, since salt (NaCl) breaks down into ions when dissolved in water these ions affect the forces between the water molecules themselves. The water molecules are dipoles and in fact the dipole-ion interactions are stronger than the hydrogen bonds holding the water molecule together. So more energy, and therefore temperature, is needed to break these new bonds and make the solution come to a boil and begin to transform into the gas state. In fact, even non-charged particles increase the boiling point of a solution because the pressure the water feels is intensified due to the fact that the particles exert their own, independent of water, pressure on the air around the liquid. The change in boiling point of a solution is a very common topic; however, most of what I found online were labs proving the Kb constant. I thought that it would be more interesting to investigate the relationship between the number of ions and the change in boiling point. The equation suggests that this relationship is linear.

I believed using compounds that would break into three or ions/compounds to be the most direct way of proving this relationship as when there are more ions the change in boiling point is less linearly proportional to the number of ions/compounds formed.Objective:Is the relationship linear between the number of ions formed from a solute and the increase in boiling point (change in temperature) when the solute is mixed with (distilled) water?MethodProcedure:• Put on lab coat, goggles, and gloves• Clean 15 (125 mL) erlenmeyer flasks carefully, make sure there are no remaining substances in the beakersfill three beakers with each of these substance:To measure 100g of water, use a 100mL volumetric flaskTo measure quantity of substance use a scale, a weighing boat (different one for each compound) (for ethylene glycol use a beaker), and a metal spatuladistilled/pure water — 100gsucrose (C12H22O11)—  0.10 mole and distilled water — 100g ethylene glycol (C2H6O2) — 0.20 mole and distilled water — 100gNaCl — 0.20 mole and distilled water — 100gCaCl2 — 0.

20 mole and distilled water — 100gPlace magnetic stirrer in erlenmeyer flask and stir until solute is completely dissolved in solvent, remove magnetic stirrer with magnetic wandSet up a probe to hang around the middle of the erlenmeyer flask using a stand and a clamp.connect the probe to a LabQuest and then to Logger Pro on a computerPut the erlenmeyer flask on heating plate and put temperature to around 350°Cheat the solution in the erlenmeyer flask until the graph on LoggerPro seems to plateau—temperature stays almost constant—and then let it continue heating until there is a clear prolonged plateau in the temperature. record temperature at which the boiling point plateaus rinse the metal spatula between the different compoundsEmpty the contents of the erlenmeyer flasks and the distilled water used to clean the instruments in a waste bucket (labeled with its contents)Repeat this process for all the erlenmeyer flasksEquipment list:12 beakersthe next values are rounded up to account for 4 trials instead of 3:2 kg of water140 grams of sugar (C12H22O11)50 grams of NaCl50 grams ethylene glycol (C2H6O2)90 grams CaCl2LabQuesttemperature probescale heating/hot platestandclampweighing boatsmetal spatulamagnetic stirrermagnetic stirrer removerIndependent Variable:solute -> Type of substanceDependent Variable:Boiling Point (Temperature — ?T)Controlled Variable:amount of solvent (100g of water)ExperimentObservationsFor some of the reactions when I was mixing the solute into the solvent I had to heat the solution to completely dissolve the solute. The solute started granuarly but over time they dissolved and were no longer visible to the naked eye.However, in the sucrose trials some of the sucrose was still visible in the solution, it was not completely dissolved.When the solution became close to reaching its boiling point some water vapor that looked similar to smoke rose out of the erlenmeyer flask.Even though some of the trials had the same final boiling point they were bubbling much less. So although on Logger Pro two reactions plateaued at the same value, one of the solutions was almost bubbling over while the other barely showed signs of boiling.

Raw DataTable 1: Raw Data values of mass, recorded boiling point, volume of solvent, and number of ions/compounds formed from soluteThe values for “Boiling Point” were recorded in Celsius (°C) but were converted to Kelvin (K). To convert from Celsius to Kelvin I added 273 to the value recorded in Celsius. The average boiling point of the solvent (water) is exactly 273.0 °C as measured in three trialsWater (H2O)Trial 1/ KTrial 2/ KTrial 3/ KAverage/ KBoiling Point273.

0272.9273.2273.0SoluteMass of substance/ ±0.001gBoiling Point/ ±0.2 KVolume of solvent (H20)/ ±0.16 mLNumber of ions formed from soluteSucrose (C12H22O11)Target: 34.230100.

001Trial 134. 238374.6100.001Trial 234.250374.4100.001Trial 334.

239374.4100.001Ethylene Glycol (C2H6O2)Target: 12.414100.

001Trial 112.601374.1100.001Trial 212.617373.7100.

001Trial 312.565374.4100.

001Sodium Chloride(NaCl)Target: 11.689100.002Trial 111.

702374.7100.002Trial 211.700375.2100.002Trial 311.703375100.

002Calcium Chloride(CaCl2)Target: 22.1968100.003Trial 122.185375.9100.003Trial 222.190376100.003Trial 322.

144376100.003An example of the graph recorded in LoggerPro to determine the boiling pointProcessed DataValues used for the Molar Mass of the substances comes from the Chemistry IB bookletAll the numbers displayed are rounded. The actual values used in the calculations are the values with the full number of significant figures available.Table 2: Processed data to find Moles, Molal, experimental ?b.p. and theoretical ?b.

p. SoluteMoles/ molMolal/ MDifference in boiling point in relation to water/ ± 0.4KTheoretical difference in boiling point/ KSucrose (C12H22O11)Target: 0.09999± 0.0029%Target: 0.9999 ± 0.16%Average: 1.

5Target: 0.512± 0.16%Trial 10.10001 ± 0.0029%1.0001± 0.16%1.60.

512± 0.16%Trial 20.10005 ± 0.0029%1.0005± 0.

16%1.40.512± 0.16%Trial 30.

10001± 0.0029%1.0001± 0.16%1.40.

512± 0.16%Ethylene Glycol (C2H6O2)Target: 0.19997± 0.0081%Target: 1.9997± 0.

17%Average: 1.1Target: 1.024± 0.17%Trial 10.

20298 ± 0.0079%2.0298± 0.17%1.11.039± 0.17%Trial 20.

20324± 0.0079%2.0324± 0.17%0.71.

041± 0.17%Trial 30.20240± 0.0080%2.0240± 0.

17%1.41.036± 0.17%Sodium Chloride(NaCl)Target: 0.20002± 0.

0086%Target: 2.0002± 0.17%Average: 2.0Target: 2.

048± 0.17%Trial 10.20024± 0.0085%2.0024± 0.

17%1.72.050± 0.

17%Trial 20.20021± 0.0085%2.0021± 0.17%2.22.050± 0.17%Trial 30.

20026± 0.0085%2.0026± 0.17%22.

051± 0.17%Calcium Chloride(CaCl2)Target: 0.20001±  0.

0045%Target: 2.0001± 0.16%Average: 3.0Target: 3.

072± 0.16%Trial 10.19990±  0.0045%1.9990± 0.16%2.93.

070± 0.16%Trial 20.19995±  0.0045%1.9995± 0.16%33.071± 0.16%Trial 30.

19953±  0.0045%1.9953± 0.16%33.065± 0.16%All calculations done for sucrose (C12H22O11), trial 1To calculate the number of moles we use the equation n=m/MM n = 34. 238/342.34 = 0.

10001 molUnc: the uncertainty in n is 0.001g or 0.0029% (absolute uncertainty/value*100%) and the molar mass has no unc.

The final unc is ±0.0029% molTo calculate the number of molal we use the equation m = n/masssolvent       mass is measured in kg M = (0.10001)/(0.10000) = 1.

0001 mol K-1Unc: the unc in n is 0.0029% mol. The unc in mass of the solvent is 0.00016 or 0.16%. Final unc = 0.0029% + 0.16% ? 0.

16%To calculate the difference/change in boiling point we use the equation ?b.p.= b.p.solution – b.p.water ?b.

p. = 374.6 – 373.

0 = 1.6 KUnc: Both of these values have an uncertainty of ± 0.2 K, so the final uncertainty is ± 0.4KTo calculated the expected change in temperature of the boiling point is given by the equation ?T = Kb*M*i  Kb: boiling point elevation constant 0.512M: molali: number of ions formed when the solute is dissolved into the solvent ?T = 0.512*1.0001*1 = 0.512 KUnc: There is no unc in Kb.

There is an uncertainty of 0.16% in M. There is no uncertainty in i. Final unc = 0.16%Summary Table: Boiling PointSucrose (C12H22O11)Ethylene Glycol (C2H6O2)Sodium Chloride(NaCl)Calcium Chloride(CaCl2)Average experimental change in boiling point/ ±0.

4K1.51.12.03.0Theoretical change in boiling point0.

512 ± 0.16%1.024 ± 0.17%2.048 ± 0.17%3.

072 ± 0.16%Graph:I did not use the value (?bp) of sucrose as it was an outlier that did not correlate with the rest of the dataConclusion As we look at the final graph in the processed data, excluding the experimental data for sucrose, the values experimentally determined are very similar to that of the theoretical change in boiling point. As we can see, the relationship between the difference in boiling point and the number of ions formed by the solute is clearly completely linear. The line of best fit (excluding the value of sucrose) has a slope of 0.9500, which is extremely close to 1, which would indicate perfect proportionality.

Also the correlation between the data is 0.9995, which further indicates the validity and correlation of the data.Both sucrose and ethylene glycol stay as one compound when being dissolved as water, sodium chloride breaks down into two ions (Na+ and Cl-), and calcium chloride breaks down into three separate ions (Ca+ and 2 Cl-) However, because the mass of one mole of sucrose is so large, around 340 grams, I elected to only use one-tenth of that amount (1 molal) I used one-fifth of a mole of all the other compounds were used in their respective reactions (2 molal), which explains why their relationship is linear and why the theoretical change in boiling point of ethylene glycol is twice that of sucrose. Normally, the more mass of the compound used the more pronounced the effect on the change in boiling point would be; however, one needs to take into account the fact that only a certain amount of mass can be dissolved into 100 mL of water. I decided that 2 molal for most of the substances (1 for sugar) would be enough mass so that there was an observable difference in the change in boiling point, but not too much mass so that it would not dissolve completely.The equation for change in boiling point when adding a solvent is ?T = M*Kb*i, so we would indeed expect the boiling point to increase linearly, and the experimental values determined are indeed extremely close to the theoretical values (see summary table), so it is clear that, at least when the compounds break down into maximum three ions, the change in boiling point increases linearly to the number of ions formed by the solute when dissolved into the solvent. The absolute uncertainty for all the experimental values is ± 0.

4K and when the percent error and percent uncertainty are calculated for every value the results are:Table to determine percent uncertainty and percent error for each soluteUncertainty and Percent ErrorSucrose (C12H22O11)Ethylene Glycol (C2H6O2)Sodium Chloride(NaCl)Calcium Chloride(CaCl2)Percent Unc26.736.420.013.3Percent error193.07.

42.32.4% uncertainty = (0.4/1.5)*100% = 26.7Percent error = |experimental – theoretical || theoretical |* 100%= (1.5-0.512)/(0.

512)*100% = 193%As we can observe in the table above, every single reaction (except for the sucrose one) contained mostly random errors because the percent uncertainty is so much smaller than the percent error for those data points. However, because the percent error for the value obtained for the change in boiling point of sucrose is so much larger, nearly eight times larger, this part of the experiment must have contained mostly (large) systematic errors, which will be discussed in the evaluation. It would be interesting in further experiments to evaluate if this linear trend (between the change in boiling point and number of ions formed by the solute) continues when the compounds used for solutes form even more ions when broken down. This trend should be evaluated for compounds that break into at least five, six, and maybe even seven ions. Substances that could be used are AlCl3 (4 ions), Al2(SO4)3 (5 ions), or Mn3N4 (7 ions).Evaluation The most evident problem in this lab are the results obtained from the reaction with sucrose. All of the values are nearly triple what they are supposed to be.

This most plausible reason for this is that not all of the nearly 35 grams of sucrose were completely dissolved and therefore more temperature was required to bring the solution to a boil. This is a systematic error that is specific to the sucrose trials. In order to fix this, we could use a larger amount of solvent, say 150 or 200 mL of water instead of 100 mL, this would allow for the sugar to completely dissolve.

Given this would decrease the molality of the solution and there might not be as pronounced an effect as the theoretical model predicts, but I hypothesize that it will be much closer than the values obtained in my trials. In order to decrease random error throughout the entirety of the experiment, we could use 500 mL of water and 10 molal of the substance. Since throughout the experiment I was measuring very small amounts of mass (of the different substances) it was easy to accidentally add a little more mass and get a much greater mass than desired, and one that was quite a high percentage of the total mass I wanted.

If we were to increase the amount of solvent and solute it would inevitably reduce random error as any human mistake would be less significant, proportionally, to the quantities needed for the adjusted experiment. The mass of water (H2O) should have been weighed not taken from volume as at different temperatures density changes (colder water is more dense than hotter water). This would change our mass and it would therefore also change the concentration of the solute in the solvent, meaning that the solute effect on the boiling point of the solution would be lower (if water is cool) or higher (if water is hot) as proportionally it is less or more significant to the total mass of the solution.Risk AssessmentOne should always wear gloves, lab coat, and goggles to protect from any accidents (glass shattering or water bubbling over). Also a waste bucket should be used to dispense of all the solutions used/made in the experiment.The only real risk in this experiment is the potential danger of the solution bubbling over the edges of the flask used to contain it when brought to a boil.

First of all, to prevent this one should simply keep a close eye on the solution being brought to a boil, and even if the graph does not show that the temperature is beginning to plateau, and turn off the heat or remove the solution if it appears that it may boil. However, if for some reason the liquid does boil over, make sure to have paper towels around the heating plate and a “heat glove” ready to remove the flask from the heating place as soon as possible.