Solution Properties by Owen Borville Lesson 13 October 26, 2025
A solution is a homogeneous mixture of two or more substances. A solution consists of a solvent and one or more solutes. Examples of solutions include air, carbonated water, vinegar, H2 gas in palladium, ethanol in water, mercury in silver, saltwater, and bronze (tin in copper), steel (carbon in iron), and brass (Cu/Zn).
Solutions can be classified by the amount of solute dissolved. An unsaturated solution is one that contains less solute than the solvent has the capacity to dissolve at a specific temperature. A saturated solution is one that contains the maximum amount of solute that will dissolve in a solvent at a specific temperature.
Supersaturated solutions contain more dissolved solute than is present in a saturated solution and are generally unstable.
Intermolecular forces:
Solvation occurs when solute molecules are separated from one another and surrounded by solvent molecules. Solvation depends on three types of interactions: (1) solute-solute interactions (2) solvent-solvent interactions (3) solute-solvent interactions.
Components of a mixture can have different properties, and there is a greater variety of intermolecular forces to consider:
Ion-dipole molecular forces: The charge of an ion is attracted to the partial charge on a polar molecule, such as NaCl or KI in H2O.
Dipole-induced dipole molecular forces: The partial charge on a polar molecule induces a temporary partial charge on a neighboring nonpolar molecule or atom. Ex. He or CO2 in H2O.
Ion-induced dipole molecular forces: The charge of an ion induces a temporary partial charge on a neighboring nonpolar molecule or atom. Ex. Fe2+ and O2.
Entropy of a system is a measure of how dispersed or spread out its energy is. There is a natural tendency for the energy of a system of a system to become dispersed (entropy increases). When separated gases are mixed together and the barrier is removed, entropy increases as disorder increases.
Like dissolves like: Two substances with similar type and magnitude of intermolecular forces are likely to be soluble in each other. Ex. CCl4 and C6H6 (nonpolar/nonpolar). CH3OH and CH3CH2OH (polar/polar).
Two liquids are miscible if they are completely soluble in each other in all proportions.
Concentration is the amount of solute relative to the volume of a solution or to the amount of solvent in a solution. Molarity (M) = moles of solute/liters of solution.
Mole fraction of a component A = xa = moles of A/sum of moles of all components.
Molality (m) is the number of moles of solute dissolved in 1 kg (1000g) solvent: molality = m moles of solute / mass of solvent (kg)
Percent by mass is percent mean parts per hundred and the multiplier is 100 (10^2)
percent by mass = mass of solute/(mass of solute + mass of solvent) x 100 %
parts per thousand and the multiplier is 1000 (10^3)
ppm is parts per million and the multiplier is 1,000,000 (10^6)
others follow with a change in the multiplier only
Comparison of Units: Concentration choice is based on the experiment:
Molarity: Titration, Gravimetric analysis (easiest in a lab situation) Based on volume of solution and temperature dependent.
Mole fraction: gases and vapor pressure
Molality and Percent by Mass: independent of temperature, dependent on mass of solvent.
Factors that Affect Solubility: Temperature affects the solubility of most substances. Pressure greatly influences the solubility of a gas.
Henry's Law states that the solubility of a gas in a liquid is proportional to the pressure of the gas over the solution: c = kP
c = molar concentration (mol/L) P = pressure (atm) k = proportionality constant called Henry's law constant
Colligative properties are properties that depend on the number of solute particles in solution. Colligative properties do not depend on the nature of the solute particles. The colligative properties are: vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.
Raoult's Law states that the partial pressure of a solvent over a solution is given by the vapor pressure of the pure solvent times the mole fraction of the solvent in the solution.
Psolution = Xsolvent * Psolvent
P1 = x1P1°
ΔP = x2P1°
P1 = partial pressure of solvent over solution
P1° = vapor pressure of pure solvent
X1 = mole fraction of solvent
ΔP = P1° - P1
x2 = mole fraction of solute
If both components of a solution are volatile, the vapor pressure of the solution is the sum of the individual partial pressures:
Pa = xaPa°
Pb = xbPb°
Pt = xaPa° + xbPb°
An ideal solution obeys Raoult's law.
Solutions boil at a higher temperature than the pure solvent.
ΔTb = Tb - Tb°
ΔTb = Kbm
ΔTb = boiling point elevation
Kb = boiling point elevation constant (°C/m)
m = molality
Freezing Point Depression: Solutions freeze at a lower temperature than the pure solvent.
ΔTf = Tf ° -Tf
ΔTf = Kfm
ΔTf = freezing point depression
Kf = freezing point depression constant (°C/m)
m = molality
Osmosis is the selective passage of solvent molecules through a porous membrane from a more dilute solution to a more concentrated one.
Osmotic pressure (π) = MRT
(π) = osmotic pressure
M = molarity (moles/L)
R = gas constant (0.08206 L*atm/mol*K)
T =absolute temperature (Kelvin)
Electrolyte solutions: electrolytes undergo dissociation when dissolved in water. The van't Hoff factor (i) accounts for this effect.
i = (actual number of particles in solution after dissociation)/(number of formulas units initially dissolved in solution
ΔTf = iKfm
ΔTb = iKbm
π = iMRT
The van't Hoff factor (i) is 1 for all nonelectrolytes:
C12H22O11(s) =H2O=> C12H22O11(aq) 1 particle dissolved, i = 1
For strong electrolytes, i should be equal to the number of ions:
NaCl(s) =H2O=>Na+(aq) + Cl-(aq) 2 particles dissolved, i = 2
Na2SO4(s) =H2O=>2Na+(aq)+(SO4)2-(aq) 3 particles dissolved, i = 3
The van't Hoff factor (i) is usually smaller than predicted due to the formation of ion pairs. An ion pair is made up of one or more cations and one or more anions held together by electrostatic forces.
Concentration has an effect on experimentally measured van't Hoff factors (i).
Colligative properties (nonelectrolyte) may be used to find solution molarity, molality, and/or a solute's molar mass.
Percent dissociation is the percentage of dissolved molecules (or formula units, in the case of an ionic compound) that separate into ions in a solution and this can be found using colligative properties.
Strong electrolytes should have complete, or 100% dissociation, however, experimentally determined van't Hoff factors indicate that this is not the case. Percent dissociation of a strong electrolyte is more complete at lower concentration.
Percent ionization of weak electrolytes is also dependent on concentration.
Colloid is a dispersion of particles of one substance throughout another substance. Colloid particles are much larger than the normal solute molecules. Categories of colloids are: aerosols, foams, emulsions, sols, gels. Examples of colloids are fog, mist, smoke, whipped cream, mayonnaise, milk of magnesia, styrofoam, jelly, butter, metal alloys like steel and gemstones.
Colloids with water as the dispersing medium can be categorized as hydrophilic (water loving) or hydrophobic (water fearing). Hydrophilic groups on the surface of a large molecule stabilize the molecule in water.
Negative ions are adsorbed (stick to the surface, not absorbed) onto the surface of hydrophobic colloids. The repulsion between like charges prevents aggregation of the articles.
Hydrophobic colloids can be stabilized by the presence of hydrophilic groups on their surface.
Emulsification is the process of stabilizing a colloid that would otherwise not stay dispersed. An emulsion is a stable mixture of two liquids that don't normally mix, like oil and water. Vigorous mixing is required to emulsify along with an emulsifier, a substance that helps stabilize the mixture by reducing surface tension between the two liquids. Without an emulsifier, the two liquids will eventually separate.
A solution is a homogeneous mixture of two or more substances. A solution consists of a solvent and one or more solutes. Examples of solutions include air, carbonated water, vinegar, H2 gas in palladium, ethanol in water, mercury in silver, saltwater, and bronze (tin in copper), steel (carbon in iron), and brass (Cu/Zn).
Solutions can be classified by the amount of solute dissolved. An unsaturated solution is one that contains less solute than the solvent has the capacity to dissolve at a specific temperature. A saturated solution is one that contains the maximum amount of solute that will dissolve in a solvent at a specific temperature.
Supersaturated solutions contain more dissolved solute than is present in a saturated solution and are generally unstable.
Intermolecular forces:
Solvation occurs when solute molecules are separated from one another and surrounded by solvent molecules. Solvation depends on three types of interactions: (1) solute-solute interactions (2) solvent-solvent interactions (3) solute-solvent interactions.
Components of a mixture can have different properties, and there is a greater variety of intermolecular forces to consider:
Ion-dipole molecular forces: The charge of an ion is attracted to the partial charge on a polar molecule, such as NaCl or KI in H2O.
Dipole-induced dipole molecular forces: The partial charge on a polar molecule induces a temporary partial charge on a neighboring nonpolar molecule or atom. Ex. He or CO2 in H2O.
Ion-induced dipole molecular forces: The charge of an ion induces a temporary partial charge on a neighboring nonpolar molecule or atom. Ex. Fe2+ and O2.
Entropy of a system is a measure of how dispersed or spread out its energy is. There is a natural tendency for the energy of a system of a system to become dispersed (entropy increases). When separated gases are mixed together and the barrier is removed, entropy increases as disorder increases.
Like dissolves like: Two substances with similar type and magnitude of intermolecular forces are likely to be soluble in each other. Ex. CCl4 and C6H6 (nonpolar/nonpolar). CH3OH and CH3CH2OH (polar/polar).
Two liquids are miscible if they are completely soluble in each other in all proportions.
Concentration is the amount of solute relative to the volume of a solution or to the amount of solvent in a solution. Molarity (M) = moles of solute/liters of solution.
Mole fraction of a component A = xa = moles of A/sum of moles of all components.
Molality (m) is the number of moles of solute dissolved in 1 kg (1000g) solvent: molality = m moles of solute / mass of solvent (kg)
Percent by mass is percent mean parts per hundred and the multiplier is 100 (10^2)
percent by mass = mass of solute/(mass of solute + mass of solvent) x 100 %
parts per thousand and the multiplier is 1000 (10^3)
ppm is parts per million and the multiplier is 1,000,000 (10^6)
others follow with a change in the multiplier only
Comparison of Units: Concentration choice is based on the experiment:
Molarity: Titration, Gravimetric analysis (easiest in a lab situation) Based on volume of solution and temperature dependent.
Mole fraction: gases and vapor pressure
Molality and Percent by Mass: independent of temperature, dependent on mass of solvent.
Factors that Affect Solubility: Temperature affects the solubility of most substances. Pressure greatly influences the solubility of a gas.
Henry's Law states that the solubility of a gas in a liquid is proportional to the pressure of the gas over the solution: c = kP
c = molar concentration (mol/L) P = pressure (atm) k = proportionality constant called Henry's law constant
Colligative properties are properties that depend on the number of solute particles in solution. Colligative properties do not depend on the nature of the solute particles. The colligative properties are: vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.
Raoult's Law states that the partial pressure of a solvent over a solution is given by the vapor pressure of the pure solvent times the mole fraction of the solvent in the solution.
Psolution = Xsolvent * Psolvent
P1 = x1P1°
ΔP = x2P1°
P1 = partial pressure of solvent over solution
P1° = vapor pressure of pure solvent
X1 = mole fraction of solvent
ΔP = P1° - P1
x2 = mole fraction of solute
If both components of a solution are volatile, the vapor pressure of the solution is the sum of the individual partial pressures:
Pa = xaPa°
Pb = xbPb°
Pt = xaPa° + xbPb°
An ideal solution obeys Raoult's law.
Solutions boil at a higher temperature than the pure solvent.
ΔTb = Tb - Tb°
ΔTb = Kbm
ΔTb = boiling point elevation
Kb = boiling point elevation constant (°C/m)
m = molality
Freezing Point Depression: Solutions freeze at a lower temperature than the pure solvent.
ΔTf = Tf ° -Tf
ΔTf = Kfm
ΔTf = freezing point depression
Kf = freezing point depression constant (°C/m)
m = molality
Osmosis is the selective passage of solvent molecules through a porous membrane from a more dilute solution to a more concentrated one.
Osmotic pressure (π) = MRT
(π) = osmotic pressure
M = molarity (moles/L)
R = gas constant (0.08206 L*atm/mol*K)
T =absolute temperature (Kelvin)
Electrolyte solutions: electrolytes undergo dissociation when dissolved in water. The van't Hoff factor (i) accounts for this effect.
i = (actual number of particles in solution after dissociation)/(number of formulas units initially dissolved in solution
ΔTf = iKfm
ΔTb = iKbm
π = iMRT
The van't Hoff factor (i) is 1 for all nonelectrolytes:
C12H22O11(s) =H2O=> C12H22O11(aq) 1 particle dissolved, i = 1
For strong electrolytes, i should be equal to the number of ions:
NaCl(s) =H2O=>Na+(aq) + Cl-(aq) 2 particles dissolved, i = 2
Na2SO4(s) =H2O=>2Na+(aq)+(SO4)2-(aq) 3 particles dissolved, i = 3
The van't Hoff factor (i) is usually smaller than predicted due to the formation of ion pairs. An ion pair is made up of one or more cations and one or more anions held together by electrostatic forces.
Concentration has an effect on experimentally measured van't Hoff factors (i).
Colligative properties (nonelectrolyte) may be used to find solution molarity, molality, and/or a solute's molar mass.
Percent dissociation is the percentage of dissolved molecules (or formula units, in the case of an ionic compound) that separate into ions in a solution and this can be found using colligative properties.
Strong electrolytes should have complete, or 100% dissociation, however, experimentally determined van't Hoff factors indicate that this is not the case. Percent dissociation of a strong electrolyte is more complete at lower concentration.
Percent ionization of weak electrolytes is also dependent on concentration.
Colloid is a dispersion of particles of one substance throughout another substance. Colloid particles are much larger than the normal solute molecules. Categories of colloids are: aerosols, foams, emulsions, sols, gels. Examples of colloids are fog, mist, smoke, whipped cream, mayonnaise, milk of magnesia, styrofoam, jelly, butter, metal alloys like steel and gemstones.
Colloids with water as the dispersing medium can be categorized as hydrophilic (water loving) or hydrophobic (water fearing). Hydrophilic groups on the surface of a large molecule stabilize the molecule in water.
Negative ions are adsorbed (stick to the surface, not absorbed) onto the surface of hydrophobic colloids. The repulsion between like charges prevents aggregation of the articles.
Hydrophobic colloids can be stabilized by the presence of hydrophilic groups on their surface.
Emulsification is the process of stabilizing a colloid that would otherwise not stay dispersed. An emulsion is a stable mixture of two liquids that don't normally mix, like oil and water. Vigorous mixing is required to emulsify along with an emulsifier, a substance that helps stabilize the mixture by reducing surface tension between the two liquids. Without an emulsifier, the two liquids will eventually separate.