Explain the following observations using chem principles:
When table salt and sugar are dissolved in water, it is observed that:
1. Both solutions have a higher boiling point than pure water.
2. the boiling point of .1 M NaCl solution is higher than that of .1 M C12H22O11 solution.
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1. I was thinking the salt and the sucrose have a higher vapor pressure which causes the vapor pressure of the solution to increase and decrease the volatility.
2. not sure about number 2
Consider on number 2 that NaCl breaks into two particles when dissoved, which doubles the effect on colligative properties.
1 is in the right ball park but the wrong direction.
For #1, both salt and sugar lower the vapor pressure of the solution compared to that of pure water. Lowering the vapor pressure of water means that it must be raised to a higher temperature before it reaches 1 atmosphere vapor presure (which is the definition of boiling point for a liquid).
2. The increase in boiling point, the decrease in freezing point, and the change in vapor pressure when a non-volatile solute is dissolved in a volatile solvent, depends solely upon the number of dissolved particles. Mole for mole, sucrose has 1 particle per molecule while NaCl has two particles per molecule. Therefore, a 0.1 M NaCl will increase the boiling point of water almost twice as much as 0.1 M sucrose.
Ok, for number 2: NaCl breaks into 2 particles when dissolved whereas sucrose remains as one. This doubles the van't Hoff factor which means a higher temperature is required for it to vaporize.
Does that make sense?
Re: number one. Is that answer ok?
Your revised number 2 answer is great. No, your original answer to number 1 is not correct. Take another look at my resonse to your number 1.
I hadn't even seen your response DrBob. Thanks! Making more sense to me now.
To explain the two observations using chemical principles:
1. Both solutions have a higher boiling point than pure water:
When a solute like table salt (NaCl) or sugar (C12H22O11) is dissolved in water, it disrupts the normal intermolecular forces between water molecules. This disruption causes an increase in the boiling point of the solution compared to pure water. The phenomenon is known as boiling point elevation.
In more detail, boiling occurs when the vapor pressure of the liquid equals the atmospheric pressure. The presence of solute particles decreases the vapor pressure of the solution. This is because the solute particles occupy some of the space at the surface of the liquid, reducing the number of solvent molecules available to escape into the gas phase. As a result, the solution requires a higher temperature to reach the same vapor pressure as pure water, thereby elevating its boiling point.
2. The boiling point of a 0.1 M NaCl solution is higher than that of a 0.1 M C12H22O11 solution:
The difference in boiling points between the two solutions can be explained by the nature of the solute particles. NaCl dissociates in water, forming Na+ and Cl- ions. On the other hand, C12H22O11 does not dissociate into ions but remains as individual sugar molecules.
The presence of ions in the solution enhances the boiling point elevation effect. This is because ions interact more strongly with water molecules due to their charges. The attraction between Na+ and Cl- ions and water molecules is stronger than the interactions between sugar molecules and water molecules. As a result, the 0.1 M NaCl solution experiences a higher boiling point elevation than the 0.1 M C12H22O11 solution, leading to a higher boiling point.
In summary, the presence of solute particles disrupts intermolecular forces, reducing the vapor pressure of the solution and elevating its boiling point. The extent of boiling point elevation depends on the nature of the solute particles, with dissociated ions causing a greater effect compared to non-ionic solutes like sugar.