Prentice Hall © 2003Chapter 13 Chapter 13 Properties of Solutions CHEMISTRY The Central Science 9th Edition David P. White.

Prentice Hall © 2003 Chapter 13

Chapter 13Chapter 13Properties of SolutionsProperties of Solutions

CHEMISTRY The Central Science

9th Edition

David P. White


• A solution is a homogeneous mixture of solute (present in smallest amount) and solvent (present in largest amount).

• Solutes and solvent are components of the solution.• In the process of making solutions with condensed

phases, intermolecular forces become rearranged.

The Solution ProcessThe Solution Process




• Consider NaCl (solute) dissolving in water (solvent):– the water H-bonds have to be interrupted,

– NaCl dissociates into Na+ and Cl-,

– ion-dipole forces form: Na+ … -OH2 and Cl- … +H2O.

– We say the ions are solvated by water.

– If water is the solvent, we say the ions are hydrated.




Energy Changes and Solution Formation

• There are three energy steps in forming a solution:– separation of solute molecules (H1),

– separation of solvent molecules (H2), andformation of solute-solvent interactions (H3).

• We define the enthalpy change in the solution process as

Hsoln = H1 + H2 + H3.

Hsoln can either be positive or negative depending on the intermolecular forces.





• Breaking attractive intermolecular forces is always endothermic.

• Forming attractive intermolecular forces is always exothermic.




• To determine whether Hsoln is positive or negative, we consider the strengths of all solute-solute and solute-solvent interactions: H1 and H2 are both positive.

H3 is always negative.

– It is possible to have either H3 > (H1 + H2) or H3 < (H1 + H2).




• Examples: – NaOH added to water has Hsoln = -44.48 kJ/mol.

– NH4NO3 added to water has Hsoln = + 26.4 kJ/mol.

• “Rule”: polar solvents dissolve polar solutes. Non-polar solvents dissolve non-polar solutes. Why?– If Hsoln is too endothermic a solution will not form.

– NaCl in gasoline: the ion-dipole forces are weak because gasoline is non-polar. Therefore, the ion-dipole forces do not compensate for the separation of ions.




– Water in octane: water has strong H-bonds. There are no attractive forces between water and octane to compensate for the H-bonds.



Solution Formation, Spontaneity, and Disorder

• A spontaneous process occurs without outside intervention.

• When energy of the system decreases (e.g. dropping a book and allowing it to fall to a lower potential energy), the process is spontaneous.

• Some spontaneous processes do not involve the system moving to a lower energy state (e.g. an endothermic reaction).




• If the process leads to a greater state of disorder, then the process is spontaneous.

• Example: a mixture of CCl4 and C6H14 is less ordered than the two separate liquids. Therefore, they spontaneously mix even though Hsoln is very close to zero.

• There are solutions that form by physical processes and those by chemical processes.






Solution Formation and Chemical Reactions

• Example: a mixture of CCl4 and C6H14 is less ordered

• Consider:

Ni(s) + 2HCl(aq) NiCl2(aq) + H2(g).

• Note the chemical form of the substance being dissolved has changed (Ni NiCl2).

• When all the water is removed from the solution, no Ni is found only NiCl2·6H2O. Therefore, Ni dissolution in HCl is a chemical process.



Solution Formation and Chemical Reactions

• Example:

NaCl(s) + H2O (l) Na+(aq) + Cl-(aq).

• When the water is removed from the solution, NaCl is found. Therefore, NaCl dissolution is a physical process.



• Dissolve: solute + solvent solution.• Crystallization: solution solute + solvent.• Saturation: crystallization and dissolution are in

equilibrium.• Solubility: amount of solute required to form a saturated

solution.• Supersaturated: a solution formed when more solute is

dissolved than in a saturated solution.

Saturated Solutions and Saturated Solutions and SolubilitySolubility



Solute-Solvent Interaction• Polar liquids tend to dissolve in polar solvents.• Miscible liquids: mix in any proportions.• Immiscible liquids: do not mix.• Intermolecular forces are important: water and ethanol

are miscible because the broken hydrogen bonds in both pure liquids are re-established in the mixture.

• The number of carbon atoms in a chain affect solubility: the more C atoms the less soluble in water.

Factors Affecting Factors Affecting SolubilitySolubility


Solute-Solvent Interaction• The number of -OH groups within a molecule increases

solubility in water.• Generalization: “like dissolves like”.• The more polar bonds in the molecule, the better it

dissolves in a polar solvent.• The less polar the molecule the less it dissolves in a polar

solvent and the better is dissolves in a non-polar solvent.



Solute-Solvent Interaction



Solute-Solvent Interaction



Solute-Solvent Interaction• Network solids do not dissolve because the strong

intermolecular forces in the solid are not re-established in any solution.

Pressure Effects• Solubility of a gas in a liquid is a function of the pressure

of the gas.



Pressure Effects



Pressure Effects• The higher the pressure, the more molecules of gas are

close to the solvent and the greater the chance of a gas molecule striking the surface and entering the solution.– Therefore, the higher the pressure, the greater the solubility.

– The lower the pressure, the fewer molecules of gas are close to the solvent and the lower the solubility.

• If Sg is the solubility of a gas, k is a constant, and Pg is the partial pressure of a gas, then Henry’s Law gives:


gg kPS


Pressure Effects• Carbonated beverages are bottled with a partial pressure

of CO2 > 1 atm.

• As the bottle is opened, the partial pressure of CO2 decreases and the solubility of CO2 decreases.

• Therefore, bubbles of CO2 escape from solution.



Temperature Effects• Experience tells us that sugar dissolves better in warm

water than cold.• As temperature increases, solubility of solids generally

increases.• Sometimes, solubility decreases as temperature increases

(e.g. Ce2(SO4)3).



Temperature Effects• Experience tells us that carbonated beverages go flat as

they get warm.• Therefore, gases get less soluble as temperature

increases.

• Thermal pollution: if lakes get too warm, CO2 and O2 become less soluble and are not available for plants or animals.



Mass Percentage, ppm, and ppb• All methods involve quantifying amount of solute per

amount of solvent (or solution).• Generally amounts or measures are masses, moles or

liters.• Qualitatively solutions are dilute or concentrated.• Definitions:

Ways of Expressing Ways of Expressing ConcentrationConcentration

100solution of mass total

solutionin component of masscomponent of % mass


Mass Percentage, ppm, and ppb

• Parts per million (ppm) can be expressed as 1 mg of solute per kilogram of solution. – If the density of the solution is 1g/mL, then 1 ppm = 1 mg

solute per liter of solution.

• Parts per billion (ppb) are 1 g of solute per kilogram of solution.



solutionin component of masscomponent of ppm


Mass Percentage, ppm, and ppb

Mole Fraction, Molarity, and Molality• Recall mass can be converted to moles using the molar

mass.



solutionin component of masscomponent of ppb

solution of moles totalsolutionin component of moles

component offraction Mole

solution of literssolute moles

Molarity


Mole Fraction, Molarity, and Molality• We define

• Converting between molarity (M) and molality (m) requires density.


solvent of kgsolute moles

Molality, m


• Colligative properties depend on quantity of solute molecules. (E.g. freezing point depression and melting point elevation.)

Lowering Vapor Pressure• Non-volatile solvents reduce the ability of the surface

solvent molecules to escape the liquid.• Therefore, vapor pressure is lowered.• The amount of vapor pressure lowering depends on the

amount of solute.

Colligative PropertiesColligative Properties


Lowering Vapor Pressure



Lowering Vapor Pressure

• Raoult’s Law: PA is the vapor pressure with solute, PA is the vapor pressure without solvent, and A is the mole fraction of A, then

• Recall Dalton’s Law:


AAA PP

totalPP AA


Lowering Vapor Pressure• Ideal solution: one that obeys Raoult’s law.• Raoult’s law breaks down when the solvent-solvent and

solute-solute intermolecular forces are greater than solute-solvent intermolecular forces.

Boiling-Point Elevation• Goal: interpret the phase diagram for a solution.• Non-volatile solute lowers the vapor pressure.• Therefore the triple point - critical point curve is lowered.



Boiling-Point Elevation• At 1 atm (normal boiling point of pure liquid) there is a

lower vapor pressure of the solution. Therefore, a higher temperature is required to teach a vapor pressure of 1 atm for the solution (Tb).

• Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m:


mKT bb


Freezing Point Depression• At 1 atm (normal boiling point of pure liquid) there is no

depression by definition• When a solution freezes, almost pure solvent is formed

first.– Therefore, the sublimation curve for the pure solvent is the

same as for the solution.

– Therefore, the triple point occurs at a lower temperature because of the lower vapor pressure for the solution.



Freezing Point Depression• The melting-point (freezing-point) curve is a vertical line

from the triple point.

• The solution freezes at a lower temperature (Tf) than the pure solvent.

• Decrease in freezing point (Tf) is directly proportional to molality (Kf is the molal freezing-point-depression constant):


mKT ff


Freezing Point Depression



Osmosis• Semipermeable membrane: permits passage of some

components of a solution. Example: cell membranes and cellophane.

• Osmosis: the movement of a solvent from low solute concentration to high solute concentration.

• There is movement in both directions across a semipermeable membrane.

• As solvent moves across the membrane, the fluid levels in the arms becomes uneven.



Osmosis• Eventually the pressure difference between the arms

stops osmosis.



Osmosis• Osmotic pressure, , is the pressure required to stop

osmosis:

• Isotonic solutions: two solutions with the same separated by a semipermeable membrane.


MRT

RTVn

nRTV


Osmosis• Hypotonic solutions: a solution of lower than a

hypertonic solution.• Osmosis is spontaneous.• Red blood cells are surrounded by semipermeable

membranes.



Osmosis• Crenation:

– red blood cells placed in hypertonic solution (relative to intracellular solution);

– there is a lower solute concentration in the cell than the surrounding tissue;

– osmosis occurs and water passes through the membrane out of the cell.

– The cell shrivels up.



Osmosis



Osmosis• Hemolysis:

– red blood cells placed in a hypotonic solution;

– there is a higher solute concentration in the cell;

– osmosis occurs and water moves into the cell.

– The cell bursts.

• To prevent crenation or hemolysis, IV (intravenous) solutions must be isotonic.



Osmosis– Cucumber placed in NaCl solution loses water to shrivel up and

become a pickle.

– Limp carrot placed in water becomes firm because water enters via osmosis.

– Salty food causes retention of water and swelling of tissues (edema).

– Water moves into plants through osmosis.

– Salt added to meat or sugar to fruit prevents bacterial infection (a bacterium placed on the salt will lose water through osmosis and die).



Osmosis• Active transport is the movement of nutrients and waste

material through a biological system.• Active transport is not spontaneous.



• Colloids are suspensions in which the suspended particles are larger than molecules but too small to drop out of the suspension due to gravity.

• Particle size: 10 to 2000 Å.• There are several types of colloid:

– aerosol (gas + liquid or solid, e.g. fog and smoke),

– foam (liquid + gas, e.g. whipped cream),

– emulsion (liquid + liquid, e.g. milk),

– sol (liquid + solid, e.g. paint),

ColloidsColloids


– solid foam (solid + gas, e.g. marshmallow),

– solid emulsion (solid + liquid, e.g. butter),

– solid sol (solid + solid, e.g. ruby glass).

• Tyndall effect: ability of a Colloid to scatter light. The beam of light can be seen through the colloid.

ColloidsColloids


ColloidsColloids


Hydrophilic and Hydrophobic Colloids• Focus on colloids in water.• “Water loving” colloids: hydrophilic.• “Water hating” colloids: hydrophobic.• Molecules arrange themselves so that hydrophobic

portions are oriented towards each other.• If a large hydrophobic macromolecule (giant molecule)

needs to exist in water (e.g. in a cell), hydrophobic molecules embed themselves into the macromolecule leaving the hydrophilic ends to interact with water.

ColloidsColloids


Hydrophilic and Hydrophobic Colloids• Typical hydrophilic groups are polar (containing C-O, O-

H, N-H bonds) or charged.• Hydrophobic colloids need to be stabilized in water.• Adsorption: when something sticks to a surface we say

that it is adsorbed.• If ions are adsorbed onto the surface of a colloid, the

colloids appears hydrophilic and is stabilized in water.• Consider a small drop of oil in water.• Add to the water sodium stearate.

ColloidsColloids


Hydrophilic and Hydrophobic Colloids

ColloidsColloids


Hydrophilic and Hydrophobic Colloids• Sodium stearate has a long hydrophobic tail

(CH3(CH2)16-) and a small hydrophobic head (-CO2-Na+).

• The hydrophobic tail can be absorbed into the oil drop, leaving the hydrophilic head on the surface.

• The hydrophilic heads then interact with the water and the oil drop is stabilized in water.

ColloidsColloids

ColloidsColloids


Hydrophilic and Hydrophobic Colloids• Most dirt stains on people and clothing are oil-based.

Soaps are molecules with long hydrophobic tails and hydrophilic heads that remove dirt by stabilizing the colloid in water.

• Bile excretes substances like sodium stereate that forms an emulsion with fats in our small intestine.

• Emulsifying agents help form an emulsion.

ColloidsColloids


Removal of Colloidal Particles• Colloid particles are too small to be separated by physical

means (e.g. filtration).• Colloid particles are coagulated (enlarged) until they can

be removed by filtration.• Methods of coagulation:

– heating (colloid particles move and are attracted to each other when they collide);

– adding an electrolyte (neutralize the surface charges on the colloid particles).

ColloidsColloids


Removal of Colloidal Particles• Dialysis: using a semipermeable membranes separate

ions from colloidal particles.

ColloidsColloids


End of Chapter 13End of Chapter 13Properties of SolutionsProperties of Solutions

Prentice Hall © 2003Chapter 13 Chapter 13 Properties of Solutions CHEMISTRY The Central Science 9th Edition David P. White.

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