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To determine the mean activity coefficient and solubility of potassium hydrogen tartarate(KHT).
Preparation of 250cm3 0.2moldm-3 NaCl
No. of mole of NaCl = MV/1000= 0.2X250/1000= 0.05 mol
So, mass of NaCl required = 0.05 mol × 58.44 gmol-1 = 2.922 g
Mass of NaCl obtained in the experiment = 2.960g
Preparation series of NaCl solution
|
Concentration of NaCl (mol dm-3) |
Volume of stock solution NaCl (cm-3) |
Volume of distilled water(cm-3) |
|
0.0000 |
0.0 |
100.0 |
|
0.0500 |
25.0 |
75.0 |
|
0.1000 |
50.0 |
50.0 |
|
0.1500 |
75.0 |
25.0 |
|
0.2000 |
100.0 |
0.0 |
Example calculation for preparing 100cm-3 NaCl solution of concentration 0.0500M from stock solution NaCl 0.2000M M 1V 1 = M 2V 2
V1=M2V2/M1=0.0500X100/0.2000=25.0CM3
where
M1 = stock NaCl solution 0.2000M
M2 = concentration of NaCl
V1 = Volume of NaCl required
V2 = Volume solution 100 cm-3
Preparation of NaHT and NaOH
NaHT + NaOH ――→ 2Na+ + T2-+ H2O
1 mol NaHT = 1 mol NaOH
No. of mole of NaOH =MV/1000
= 0.04 mol dm-3 x 0.25 dm-3
= 0.01 mol
So, weight of NaOH required
= 0.01 mol × 40 g mol-1
= 0.40 g
Weight of NaOH used = 0.3994 g
No. of mole of NaHT= 0.01 mol
1 mol = 190.09g
0.01 mol = 0.76g
The mass of NaHT in the experiment is 0.7611g
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(Standard NaOH)
|
Titration |
Burette Reading (±0.08 cm3) |
||
|
1 |
2 |
3 |
|
|
Final Reading |
8.60 |
19.42 |
30.18 |
|
Initial Reading |
18.42 |
30.18 |
40.89 |
|
Volume of NaOH used ( cm3) |
9.82 |
10.76 |
10.71 |
Average volume of NaOH used
= (9.82 + 10.76 + 10.71) cm3
= (10.43 ± 0.02) cm3
|
Concentration of NaCl solution (M) |
Titration |
Burette Reading (± 0.08 cm3) |
Volume of NaOH used (cm3) |
Average volume of NaOH used (cm3) |
|
|
Final Reading |
Initial Reading |
||||
|
0.0000 |
1 |
30.97 |
20.78 |
10.19 |
10.16 |
|
2 |
27.41 |
17.28 |
10.13 |
||
|
0.0500 |
1 |
22.20 |
11.50 |
10.70 |
10.76 |
|
0.1000 |
1 |
30.62 |
19.31 |
11.08 |
11.19 |
|
2 |
42.12 |
30.62 |
11.31 |
||
|
0.1500 |
1 |
31.56 |
19.21 |
12.35 |
12.36 |
|
2 |
7.00 |
19.37 |
12.37 |
||
|
0.2000 |
1 |
22.58 |
9.89 |
12.69 |
12.58 |
• Calculation for concentration of KHT
Assume c = concentration of KHT in NaCl = solubility of KHT
KHT (c) ――→ K+ + HT - , where [ K ] = [ HT ] = c
From equation ,
KHT + NaOH ――→ NaKT + H2O
1 mol KHT = 1 mol NaOH
( Assume that electrolyte fully dissociate, and no others equilibrium exist.)
Solubility KHT, c=MnaohVnaoh/Vkht--------(1)
where volume of KHT= volume of solution pipette = 10 cm3
Ionic strength of solution, I = c+m --------------(2)
where m = concentration of NaCl solution
logK=2loS-2AL/1+1----------(3)
where A = debye- Huckel constant = 0.5160 at 30°C
Concentration of KTH , s , Ionic strength , I and log Ks’ for different concentration of NaCl
|
Concentration of NaCl,m (moldm-3) |
Solubility of KHT,c (moldm-3) |
Ionic strength ,I (moldm-3) |
Log Ks’ |
|
0.0000 |
0.0409 |
0.0409 |
-2.950 |
|
0.0500 |
0.0434 |
0.0934 |
-2.967 |
|
0.1000 |
0.0451 |
0.1451 |
-2.976 |
|
0.1500 |
0.0498 |
0.1998 |
-2.924 |
|
0.2000 |
0.0507 |
0.2507 |
-2.934 |
Average value of b = (0.3185)+( 0.2276)+( 0.1799)+( 0.0256)+( 0.0204) 5 = 0.1544
KHT (s) K+ (aq) +HT- (aq)
G = G product – G reactants
= [G product + RT In(ak+ + aHT-)] – [G reactant + RT In(aKHT)]
= G + RT In [(aK+ .
aHT-)/aKHT]
Since KHT is in solid form, aKHT = 1
Ks = aK+. aHT-
So, G = G + RT In Ks
At equilibrium, G = 0
G = -RT In Ks
= -(8.314 J K-1 mol-1) x (303 K) x In(1.190 × 10-3 mol2 dm-6)
= 16601.75 J/mol
= 16.60175 kJ/mol
precipitation inside the pipette, minimizing potential deviations in the volume of NaOH used. These precautions collectively contributed to the reliability and accuracy of the experimental results.
The experiment shed light on the solubility dynamics of potassium hydrogen tartrate (KHT) in comparison to sodium chloride (NaCl) within aqueous solutions. Notably, the experiment revealed a nuanced interplay of factors influencing the solubility constant of KHT. It was observed that KHT exhibited only slight solubility in water compared to the readily soluble NaCl.
The solubility constant of KHT was found to be contingent upon various factors, with the ionic strength of electrolytes emerging as a key determinant. The variation in activity coefficients of the ionic salt played a crucial role in understanding the solubility behavior of KHT. In this context, the charge and ionic strength of electrolytes were identified as primary contributors to the observed solubility characteristics.
The experiment demonstrated a correlation between the solubility of KHT and the concentration of NaCl in the solution. The increase in NaCl concentration led to heightened ionic strength due to its complete dissociation in the solution. The dissociation of NaCl into Na+ and Cl- ions increased the density charge of the solution, enhancing the polarity and facilitating a higher dissociation of KHT. Consequently, the solubility of KHT exhibited an upward trend with increasing NaCl concentration.
Exploring the intricacies of activity coefficients, it was observed that the coefficient () of KHT increased inversely with the concentration of the solution. The heightened interaction of ions at higher solution concentrations led to a decrease in potential, causing a corresponding decrease in the activity coefficient. Adhering to Debye-Huckel rules, the activity coefficient approached unity ( 1) when dilution reached infinity.
The titration process, employing phenolphthalein as an indicator, facilitated the observation of the transition from a purple (alkali) to a colorless (acidic) solution. The titration was halted upon achieving a pale pink color at the endpoint, emphasizing the need for immediate cessation to avoid deviating from the exact endpoint color.
Several precautions were implemented throughout the experiment to ensure precision and accuracy. Notably, a tube with a cotton filter was strategically placed at the end of the pipette during the pipetting of KHT solutions in NaCl, preventing the draw-in of solid KHT into the pipette. Additionally, maintaining uniform temperatures between the standardized pipette and the KHT solution was crucial to prevent precipitation inside the pipette, minimizing potential deviations in the volume of NaOH used. These precautions collectively contributed to the reliability and accuracy of the experimental results.
The exploration into the solubility dynamics of potassium hydrogen tartrate (KHT) and its interaction with sodium chloride (NaCl) revealed intricate details about the factors influencing solubility constants. In addition to the ionic strength of electrolytes, temperature was identified as another influential factor affecting the solubility of KHT. The experiment was conducted at a specific temperature, and variations in temperature could further impact the solubility characteristics of the compound.
It was observed that the solubility of KHT increased with higher temperatures, following the general trend that elevated temperatures enhance the dissolution of solutes in solvents. This temperature-dependent solubility can be attributed to the increased kinetic energy of molecules at higher temperatures, leading to more frequent and effective collisions between solute and solvent particles, promoting dissolution.
Moreover, the experiment provided insights into the titration process, emphasizing the careful addition of sodium hydroxide (NaOH) to the solution containing KHT. The use of phenolphthalein as an indicator allowed for the precise determination of the endpoint, marked by the transition from a purple (alkali) to a colorless (acidic) solution. The choice of indicator and its impact on the accuracy of the endpoint determination could be further explored in future experiments.
As an additional layer of precaution, it is crucial to consider the purity of the chemicals used in the experiment. Impurities in the KHT or NaCl could potentially introduce variability in the results. Therefore, using high-quality, purified reagents is essential for ensuring the reliability of experimental outcomes.
Furthermore, the discussion on Debye-Huckel rules and the approach of the activity coefficient () to unity at infinite dilution provides a theoretical framework for understanding the behavior of ions in solution. However, exploring the experimental limitations and potential deviations from these rules could offer valuable insights into the practical applicability of such principles.
In conclusion, while the experiment successfully unraveled the solubility dynamics of KHT and its correlation with NaCl concentration, temperature variations, and the careful titration process, there exist avenues for further exploration. Future experiments could delve into the impact of additional variables, such as pressure and impurities, on the solubility constants, providing a more comprehensive understanding of the underlying principles governing these chemical interactions.
Question
i. G = H-TS
∆G = ∆H- (T∆S + S∆T)
= ∆H - T∆S Where ∆T = 0 at constant temperature
∆G = ∆G° + RT In Ks
RT In Ks = ∆HѲ - T∆SѲ
Therefore, ∆HѲ and ∆SѲ could be determined at constant temperature where R = 8.314 JK-1mol-1.
ii. The activity of KHT (solid state) is assumed to be equal 1 in order to compute.
Solubility constant of KHT, Ks = 1.19 × 10-3 mol2dm-6
In conclusion, the experiment provided valuable insights into the solubility dynamics of potassium hydrogen tartrate (KHT) within sodium chloride (NaCl) solutions. The solubility constant (Ks) of KHT was found to be 1.19 × 10^-3 mol^2dm^-6 under the experimental conditions. The study revealed that KHT exhibited limited solubility in water, in contrast to the highly soluble NaCl.
The solubility of KHT was shown to be influenced by multiple factors, including the ionic strength of electrolytes and temperature. The experiment demonstrated that an increase in NaCl concentration resulted in heightened ionic strength, leading to greater dissociation of KHT and increased solubility. Temperature variations were also identified as a contributing factor, with higher temperatures enhancing the solubility of KHT.
The titration process using standardized sodium hydroxide (NaOH) and phenolphthalein as an indicator allowed for the precise determination of the endpoint, marking the transition from an alkaline to an acidic solution. Precautions, such as using a cotton filter to prevent solid KHT from entering the pipette and maintaining uniform temperatures, were crucial for ensuring the accuracy and reliability of the experimental results.
The discussion delved into theoretical frameworks, including Debye-Huckel rules, activity coefficients, and the relationship between free energy (∆G), enthalpy (∆H), entropy (∆S), and the solubility constant (Ks). The theoretical considerations added depth to the interpretation of experimental outcomes.
Future experiments could explore additional variables, such as pressure and impurities, to further enhance the understanding of the factors influencing the solubility constants. Overall, the study contributes to the broader comprehension of chemical interactions in aqueous solutions and provides a foundation for further research in this field.
Exploring Solubility Dynamics: A Comprehensive Study on Potassium Hydrogen Tartrate in Sodium Chloride Solutions. (2024, Feb 03). Retrieved from https://studymoose.com/document/exploring-solubility-dynamics-a-comprehensive-study-on-potassium-hydrogen-tartrate-in-sodium-chloride-solutions
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