The crystallization process is happening in a hand warmer.
The properties of a solution are different from those of pure solutes or solvent.
The chemical identity of the solute affects many solution properties.
Compared to pure water, a solution of hydrogen chloride is more acidic, a solution of ammonia is more basic, a solution of sodium chloride is more dense, and a solution of sucrose is more viscous.
This small set of properties is important to many natural phenomena and technological applications.
Several units commonly used to express the concentrations of solution components were introduced in an earlier chapter of this text, each providing certain benefits for use in different applications.
molar concentrations will vary because of solution volumes.
The concentration of a solution with identical numbers of solute and solvent species will be different at different temperatures due to the contraction/expansion of the solution.
mole-based concentration units have values that are not dependent on temperature.
Since these units are computed using only mass and molar amounts, they do not vary with temperature and are better suited for applications requiring temperature-independent concentrations, as will be described in this chapter module.
A mixture of equal volumes of water and ethylene glycol is used in most automobile radiators.
The mole fraction is a ratio of properties with the same units.
To convert from one concentration unit to another, you need to compare the two unit definitions.
Both units have the same numerator but different denominators.
3.0 mol NaCl is the numerator for this solution's mole fraction.
The definition for mole fraction is 55 mol H2O.
The mole fraction of I2 is 0.115.
A lowering of the liquid's vapor pressure can be achieved by dissolving a nonvolatile substance.
The effect of solute molecule on the liquid's vaporization and condensation processes can be rationalized.
The solvent must be present on the solution's surface.
The rate of solvent vaporization is reduced by the presence of solute.
Since the rate of condensation is unaffected by the presence of solute, the net result is that the vaporizationcondensation equilibrium is achieved with fewer solvent molecules in the vapor phase.
To understand the lowering of a liquid's vapor pressure, it is necessary to note that the greater entropy of a solution in comparison to its separate solvent and solute serves to effectively stabilizing the solvent molecule and hinder their vaporization.
The OpenStax book is available for free at http://cnx.org/content/col11760/1.9 a higher boiling point as described in the next section of this module.
The vapor pressure of a solution is lowered by the presence of nonvolatile solutes.
The partial pressure exerted by any component of an ideal solution is the same as the vaporpressure of the pure component.
The mole fraction of A is the solution.
The provided mass data can be used to calculate the molar amounts of each component.
Next, calculate the mole fraction of the solvent and use Raoult's law to compute the solution.
A solution contains urea, CO(NH2)2 and water.
Due to the presence of nonvolatile solutes, the solution's boiling point will be increased.
Compared to pure solvent, a solution will require a higher temperature to achieve any given vapor pressure, including one equivalent to that of the surrounding atmosphere.
Depending on the identity of the solvent, boiling point elevation constants are characteristic properties.
The total number of solute particles present in a given amount of solvent, not the mass or size of the particles, determines the extent to which the vapor pressure of a solvent is lowered and the boiling point is elevated.
The equation relating boiling point elevation to solute molality can be used to solve this problem.
The pure solvent's boiling point should be added to the boiling point elevation.
If the solution is ideal and the iodine is nonvolatile, you can find the boiling point of it.
Four steps can be used to solve this problem.
The molality of the solution is determined by the number of moles of solute and the mass ofsolvent.
To determine how much the boiling point changes, use the direct proportionality between the change in boiling point and molalconcentration.
Determine the new boiling point from the boiling point of the pure solvent.
An ideal solution is what you should assume.
Distillation is used in both the laboratory and in industrial settings to separate the components of a mixture.
It is used to purify water.
The separation technique involves heating a sample mixture to get it to evaporate, condense, and collect interest.
A schematic diagram of the components is shown in a photograph.
When components reach adequately cool zones, they condense and are collected.
Crude oil is a complex mixture that is separated by large-scale fractional distillation.
Solutions freeze at lower temperatures than pure liquids.
This phenomenon is used in de-icing schemes that use salt, calcium chloride, or urea to melt ice on roads and sidewalks, and in the use of ethylene glycol as an antifreeze in automobile radiators.
The body fluids of fish and other cold-blooded sea animals that live in the ocean remain unfrozen even at a lower temperature than fresh water.
Rock salt, calcium chloride, or a mixture of the two are used to melt ice.
The values of these properties depend on the chemical identity of the solvent.
The equation relating freezing point depression to solute molality can be used to solve the problem.
The pure solvent's freezing point is observed.
Due to the fact that a solution of any one of these salts will have a freezing point lower than 0 degrees Centigrade, they are often used to de-ice roads and sidewalks.
The group 2 metal salts are often mixed with the cheaper and more readily available rock salt, since they are less corrosive and provide a larger depression of the freezing point.
If you want to de-ice a plane prior to takeoff, you shouldn't use these ionic compounds.
For these applications, covalent compounds are often used.
The freezing point of the liquid is lowered by the use of glycols, which elevate the boiling point, making the fluid useful in both winter and summer.
The freezing point depression can be used to remove ice from the control surfaces of aircraft.
The phase diagrams for a pure liquid and a solution derived from that liquid are used to show the colligative effects on vapor pressure, boiling point, and freezing point.
The phase diagrams show water and a solution of nonelectrolyte.
The solvent and solution have the same solid-gas curves.
This is the case for many solutions.
transitions between the solid and gaseous phases are not subject to colligative effects because they are composed solvent only.
A more simplistic example of technology is the use of dialysis tubing to remove waste from blood.
The molecule will diffuse in both directions.
A net transfer of solvent molecule from pure solvent to solution is achieved.
The volume of the solution will increase when it is carried out in an apparatus like that shown.
The weight of the column of solution in the tube causes the level of the solution to rise, increasing its hydrostatic pressure, and this results in a faster transfer of solvent molecule back to the pure solvent side.
When the pressure reaches a value that yields a reverse solvent transfer rate, bulk transfer of solvent ceases.
This technique is used for large-scale desalination of seawater and on smaller scales to produce high-purity tap water for drinking.
Applying a pressure greater than the osmotic pressure of the solution will reverse osmosis.
The solvent is pushed into the solution.
In the process of osmosis, water is moved through a semipermeable membrane from a less concentrated solution to a more concentrated solution.
Osmosis is caused by the amount of pressure applied to the solution.
The water will go from being more concentrated to being less concentrated if more pressure is applied.
This is called reverse Osmosis.
Water purification using reverseosmosis is used in many applications, from desalination plants in coastal cities to water-purifying machines in grocery stores and smaller reverseosmosis household units.
Vehicles can be used to transport generator-operated RO units to remote locations.
There are reverse osmosis systems for purifying drinking water shown here.
In many biological systems, cells are surrounded by semipermeable membranes.
If carrots and celery are placed in water, they can be crisp again.
Water moves into the carrot cells.
A cucumber placed in a concentrated salt solution loses water and becomes a pickle.
Animals can also be affected byOsmosis.
When solutions are injected into the body, concentrations of solute are important.
The osmotic pressure is approximately 7.7 atm.
Red blood cell membranes are impermeable to water and will swell and possibly die in a hypotonic solution, as well as maintaining normal volume and shape in an isotonic solution.
Concentration of solute present is directly proportional to Osmotic pressure and changes in freezing point, boiling point, and vapor pressure.
We can use a measurement of one of these properties to determine the mass of the solute.
A solution of 4.00 g of nonelectrolyte dissolved in 58.0 g of benzene is found to freeze at 2.32 degrees.
The following steps can be used to solve this problem.
Determine the mass of solvent used to make the solution and the number of moles of compound in it.
Determine the number of moles in the mass of the solute.
The boiling point of the solution of 35.7 g of nonelectrolyte in 220.0 g of chloroform is 64.5%.
The osmotic pressure of the 0.500 L sample is 5.9 torr.
Determine the molar concentration from the osmotic pressure by converting it to atmospheres.
The number of moles in the solution is determined by the concentration and volume of the solution.
Determine the number of moles in the mass from the mass of hemoglobin.
The colligative properties of a solution are dependent on the number, not the kind of solute species dissolved.
One mole of any nonelectrolyte dissolved in 1 kilogram of solvent lowers the freezing point the same way as one mole of any other nonelectrolyte.
The effect on the freezing point is the same for each individual ion.
In a solution containing 4.2 g of NaCl dissolved in 125 g of water, the concentration of ion is approximately the same.
Determine the freezing temperature of the solution, which is approximately equal to the freezing temperature of seawater, by assuming that each of the ion in the NaCl solution has the same effect on the freezing point of water as a nonelectrolyte molecule.
The following steps can be used to solve this problem.
The unit conversionfactor is used to convert from grams to moles of NaCl.
Determine the number of moles of ion present in the solution using the number of moles of ion in 1 mole of NaCl as the conversion factor.
Determine the molality of the solution from the number of moles and the mass of solvent.
To determine how much the freezing point changes, use the direct proportionality between the change in freezing point and molalconcentration.
Determine the new freezing point from the old one.
The value of 3.4 degC is obtained when this solution is actually prepared.
As solute concentrations increase, the differences between the measured and expected colligative property values become more significant.
The observations suggest that the ion of sodium chloride is not completely separated from the solution.
To account for this and avoid the errors associated with the assumption of total dissociation, an experiment named in honor of a German chemist is used.
The chemists Peter Debye and Erich Huckel came up with a theory to explain the incomplete ionization of strong electrolytes in 1923.
Interionic attraction is reduced by solvation of the ion and the polar solvent, but it is not completely nullified.
In some cases, a positive and negative ion may actually touch, giving a solvated unit called an ion pair.
Ions become more and more widely separated, and the residual interionic attractions become less and less.
In extremely dilute solutions, the effective concentrations of the ion activities are the same as the actual concentrations.