Solubility Of Kno3 In Water
Abstract
Water is a polar solvent hence it is the most normally used fluid. Apart from the polarity of a solvent in relation to the substance being dissolved, temperature also plays an important role in the dissolution of substances. This experiment expected to establish the relationship between temperature and the solubility of KNOthree. The solubility of KNO3 in water and in other salt solutions was adamant at different temperatures. Information technology was realized that the solubility of KNO3 increased with a raise in temperature and that the chemical composition of dissimilar solvents affected the solubility of KNO3.
H2o is the well-nigh suitable solvent for dissolving ionic compounds due to its polarity. The polar nature of water allows negative ions from ionic compounds to be attracted to the partially positive hydrogen ion. Ions with positive charges, conversely, are pulled towards oxygen ions, which possess negative charges. Strong forces of attraction hold the oppositely charged ions within molecules. Therefore, energy is required to overcome these forces of attraction. This energy is known as the enthalpy of solution and is denoted by -∆Ho soln. Factors such equally temperature and the nature of the solvent too impact the amount of energy required to dissolve a chemical compound.
This experiment was aimed at finding the solubility of KNOthree in water equally well as in solutions of other salts. It intended to observe out the temperature dependence of the solubility of KNO3 in water. The experiment also sought to approximate the heat of solution for potassium nitrate. The final objective was to estimate the molal solubility of potassium nitrate in relation to temperature in solutions containing any of the two ions that fabricated upwards KNO3.
Procedure
Solubility of KNOiii in H2o
About 20g of water was weighed equally accurately as possible into a test-tube whose length was eight inches. 9g of potassium nitrate was added to the water, and the test tube was shaken until all the salt was submerged in the water. A coated wire stirrer was so placed in the tube later which a thermometer was inserted in such manner that the bulb of the thermometer did not arrive contact with the sides or the bottom of the test tube. It was ensured that the wire stirrer could move up and downwards (in the examination tube) without upsetting the thermometer. The solution was stirred and heated lightly. The heating stopped when all the crystals liquefied. The flame was removed, after which a beaker containing common cold water was placed under the test tube.
The chalice was raised so that the test tube was partially immersed in the water. Thereafter, the solution was stirred vigorously until it became cloudy. The cloudiness was an indication that crystals were starting to develop. The precise temperature at which crystals began to form was noted and recorded as the solubility temperature. The more accurate estimation of the same temperature was determined by reheating the tube and stirring the contents vigorously until the temperature was between ii and 3 oC higher than the previously recorded temperature. The test tube was removed from the heat, after which the salt was stirred vigorously without cooling in a water bath. The temperature at which the beginning cloudiness appeared was noted and recorded every bit the truthful solubility temperature.
The above procedures were repeated until consistent temperatures were obtained. The purpose of repetition was to check the accuracy of the results. 3g of KNO3 was added to the previously used mixture to make a concentration of 12g. The solubility of the new concentration was determined as previously described. The procedure was repeated to obtain solubility temperatures for concentrations of 15g and 18g. The obtained values were recorded in a table aslope corresponding molal solubility in terms of moles of KNO3 per kilogram of solvent mass (Grossie and Underwood 259).
Handling of Data
Assuming that the KNO3 in aqueous solution was completely dissociated, then
Ksp= [Thousand+]•[NO3 –] = S2 (8) where 's' was the molal solubility. The application of thermodynamic equations to ionic solutions gave the equation
Ln Msp= -∆Ho soln/R• (1/T)+ C (9) where Ln Ksp was the solubility product abiding, -∆Ho soln denoted the standard heat of solution for the solute, T was the absolute temperature, R was the gas constant and C was a abiding of integration (Grossie and Underwood 259). The molal solubility was determined and tabulated for each of the four temperatures. The values of the parameters in equation 9 were likewise calculated and tabulated. Thereafter, a plot of Ln One thousandsp versus one/T was fabricated. The plot was used to determine the value of ∆Ho soln.
A second plot of ∆G o soln (J/Mol) against T (Yard) was made. The plot was used to obtain ∆Ho soln from the slope. The experimental values were and so compared to the literature values.
Solubility in Solutions with a Mutual Ion
It was anticipated that the solubility of KNO3 would reduce in solutions that contained ions that were nowadays in both salts because of the mutual ion effect. The solubility of potassium nitrate in salts that independent common ions was established using i or ii molal solutions of KCl, NaNOiii and NaCl. The procedure that was used was the same as that in department A. The solubility temperatures of KNO3 in different salts were tabulated. The KNOiii solubility data were plotted aslope the additional solubility data on the same nautical chart.
Results
Tabular array 1: The solubility of KNO3 in water.
∆Ho soln=-R*Gradient =-eight.3143*-6830 =56,786.669 J/mol =56.787 KJ/mol
-∆So soln= Gradient
-∆Due southo soln= 210.five
∆So soln= 210.five J/Kmol
Table ii: The solubility of KNO3 in 1M NaCl.
∆Southo soln= 174.i J/Thousand*mol; ∆Ho soln=44.926 KJ/mol
Table 3: The solubility of KNOthree in 2M NaCl.
∆So soln= 220.2 J/G*mol; ∆Ho soln=58.778 KJ/mol
Table four: The solubility of KNO3 in 1M NaNOthree.
∆So soln=175.iv J/K*mol; ∆Ho soln=45.284 KJ/mol
Table five: The solubility of KNOiii in 2M NaNO3.
∆Due southo soln= 240.2 J/K*mol; ∆Ho soln=66.858 KJ/mol
Table 6: The solubility of KNO3 in 1M KCl.
∆So soln= 190.viii J/Thou*mol; ∆Ho soln=l.429 KJ/mol
Tabular array 7: The solubility of KNO3 in 2M KCl.
∆So soln= 212.7 J/Chiliad*mol; ∆Ho soln=57.988 KJ/mol
Discussion
Slight differences were seen in the experimental values and literature values of ∆So soln and ∆Ho soln. The experimental ∆Ho soln for KNO3 in water was institute to exist +56.787 KJ/Mol, whereas the literature value co-ordinate to Grossie and Underwood was +fifty.nine KJ/Mol (263). The experimental ∆So soln, in contrast, was +210.5 J/K*mol while the literature value was +193 J/K*mol. There were slight differences in the molal solubility values from the experiment and the literature values. For example, the experimental molal solubility for KNO3 in h2o at 20 oC was four.914, whereas the literature value at the same temperature was 3.13. From figure nine, information technology was realized that the solubility of KNO3 varied in different solvents. The lowest solubility was in 2M NaCl, whereas the highest solubility was achieved in 2M KCl. It was realized that the solubility was low in the 1M solutions and higher in the 2M solutions. It was likewise seen that the solubility in the 1M solutions was lower than in water. All the same, the highest solubility was seen in a 2M solution of KCl. The differences could be due to experimental errors, such as erroneous thermometer readings as well as taking erroneous weights of solvents and salts. Information technology was also possible that some of the temperature readings were taken too early or too late when determining the solubility of KNO3 in water.
The observed overall trend was that the solubility of KNOiii rose as the temperature was elevated. The observed values were within the experimental uncertainties, such as those reported in other similar experiments. Therefore, it was ended that the solubility of KNO3 increased with a raise in the temperature of the solution regardless of the nature of the solvent.
Works Cited
Grossie, A David and Kirby Underwood. Laboratory Guide for Chemistry. 6th ed. 2013. Plymouth, MI: Hayden-McNeil. Print.
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