Solutions and The Dissolving Process
Solutions and Their Characteristics
Introduction:
In 2010, Hamilton, Ontario, warned residents of potential lead contamination in drinking water, especially homes built before 1955 due to lead pipes connecting homes to the main water supply.
Chronic lead exposure is dangerous, impairing brain and nervous system development, particularly in children.
Lead Contamination:
Lead, being just above hydrogen in the activity series, can react with water, especially if it’s slightly acidic, releasing toxic lead ions.
The longer water is in contact with lead pipes, the higher the concentration of lead ions.
Officials advised residents to run taps for five minutes before using water or replace lead pipes.
Tap Water Composition:
Tap water is a mixture of dissolved substances, including natural minerals and gases, as well as intentionally added substances like chlorine for disinfection and fluoride to prevent dental decay.
Municipalities conduct frequent tests to ensure tap water is a clear, colorless, and safe solution.
What Is a Solution?
Definition:
A solution is a homogeneous mixture of two or more substances with only one phase and uniform mixing, resulting in a uniform appearance.
Samples from different locations in the solution have the same composition.
States of Solutions:
Solutions can be solids, liquids, or gases.
Liquid and gaseous solutions are transparent because the entities are too small to block light.
Solutions may be colored or colorless depending on the substances they contain.
Example: Glucose Solution:
Glucose (C6H12O6) is dissolved in water for intravenous (IV) drips.
After mixing, glucose grains are no longer visible as molecules separate and disperse evenly.
Each molecule is in direct contact with water molecules and will not settle due to their small mass.
The diameter of these molecules is extremely small (in the order of 10^{-9} m), allowing light to pass through, creating a clear, homogeneous mixture.
Homogeneous and Heterogeneous Mixtures
Homogeneous Mixtures (Solutions):
Uniform and have only one phase.
Example: Solutions.
Heterogeneous Mixtures:
Have two or more phases.
Example: Oil and vinegar separate into distinct layers.
Appearance vs. Reality:
Some mixtures like blood and milk appear homogeneous but are heterogeneous upon closer inspection.
Blood contains different cells suspended in a liquid, making its composition non-uniform.
All translucent or opaque liquid and gaseous mixtures are heterogeneous.
Air as an Example:
Dry, clean air is a homogeneous mixture (solution) of nitrogen, oxygen, argon, and carbon dioxide.
It is transparent because gas entities are too small to interfere with light.
Classroom air appears homogeneous, but dust particles deflect projector light, revealing it as a heterogeneous mixture.
The same effect can be seen in liquids.
Components of a Solution
Solute and Solvent:
A solution is a homogeneous mixture of a solute (the substance in the lesser quantity) in a solvent (the substance in the greater quantity).
Example: In a glucose solution, glucose is the solute, and water is the solvent.
Example: White Rum:
White rum is a 40% by volume solution of ethanol in water.
Ethanol is the solute, and water is the solvent.
In countries where rum contains up to 75% ethanol, water becomes the solute, and ethanol is the solvent.
Multiple Solutes:
A solution may contain more than one solute.
Types of Solutions
Variety of Solutions:
Solutions exist in different states all around us and within us.
Air is a gaseous solution of gases dissolved in nitrogen.
Urine is an aqueous solution of molecular and ionic compounds.
Liquid Solutions with Non-Water Solvents:
Many liquid solutions do not use water as the solvent.
Gasoline is a solution of liquid hydrocarbons and other solutes.
Different grades of gasoline have different compositions of solutes, affecting how they burn.
Concentration:
Concentration describes the ratio of solute to solution or solvent.
Concentration = \frac{Quantity : of : Solute}{Quantity : of : Solution : or : Solvent}
Concentrated vs. Dilute Solutions:
A concentrated solution has a relatively high quantity of solute per volume of solution.
A dilute solution has a relatively low quantity of solute per volume of solution.
Alloys
Definition:
A solution of two or more metals.
Gold used for jewelry is alloyed with silver and copper to stiffen it and change its color.
Purity is measured in karats: 24k gold is pure; 14k gold contains 58% gold.
Amalgams:
Alloys of mercury, the only liquid metal at room temperature.
Silver-colored dental fillings are amalgams of mercury and silver.
Mercury can be absorbed by gold, discoloring it.
Solute/Solvent Combinations
States of Solutes and Solvents:
Solutes and solvents can be solid, liquid, or gas.
Metals only form alloys when heated until molten.
Examples of Solute/Solvent Combinations:
Air: gas in gas (oxygen, argon, carbon dioxide in nitrogen)
Carbonated beverages: gas in liquid (carbon dioxide in water)
Humidity: liquid or solid in gas (water molecules in air)
Alcoholic beverages: liquid in liquid (ethanol in water)
Dental fillings: solid in liquid (mercury amalgams)
Air fresheners: solid in gas (vapors from scented solids in air)
Clear apple juice: solid in liquid (flavor compounds in water)
Brass: liquid in liquid (copper and zinc)
Aqueous Solutions
Definition:
Solutions in which water is the solvent (
Solutions in which water is the solvent
Electrical Conductivity of Aqueous Solutions
Molecular Compounds:
Form solutions containing neutral molecules and do not conduct electricity.
Example: Sugar dissolved in water.
Ionic Compounds:
Form solutions containing ions and conduct electricity.
Example: Salt dissolved in water.
Electrolytes:
Substances that form conducting solutions are called electrolytes.
Strong Electrolytes:
Ionic compounds that dissociate completely into ions when dissolved in water.
Form solutions that are very good conductors of electricity.
Example: Sodium chloride (NaCl).
Weak Electrolytes:
Molecular compounds that form a small number of ions when dissolved in water.
Form solutions that are weak conductors of electricity.
Example: Acetic acid (CH3COOH).
Non-Electrolytes:
Substances that dissolve in water but do not produce ions.
Form solutions that do not conduct electricity.
Example: Sugar (C12H22O11).
Solubility
Definition:
The maximum quantity of solute that can dissolve in a certain quantity of solvent at a specified temperature.
Saturated Solution:
A solution containing the maximum quantity of solute.
Additional solute will not dissolve and will remain undissolved.
Unsaturated Solution:
A solution containing less than the maximum quantity of solute.
More solute can be dissolved.
Supersaturated Solution:
A solution containing more than the maximum quantity of solute.
This is a temporary, unstable condition. If disturbed the excess solute will precipitate out of the solution as crystals
Factors Affecting Solubility
Solute-Solvent Interactions:
"Like dissolves like": Polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes.
Example: Water (polar) dissolves sugar (polar), while gasoline (nonpolar) dissolves oil (nonpolar).
Temperature:
Generally, the solubility of solids in liquids increases with temperature.
The solubility of gases in liquids decreases with temperature.
Example: Carbonated beverages lose carbon dioxide (CO2) more readily when warm.
Pressure:
Has little effect on the solubility of solids and liquids.
Significantly affects the solubility of gases in liquids.
Henry’s Law: The solubility of a gas is directly proportional to its partial pressure above the solution.
Example: Carbonated beverages are bottled under high pressure to increase the solubility of carbon dioxide.
Solubility and Saturation
Dynamic Equilibrium:
In a saturated solution, a dynamic equilibrium exists between dissolved and undissolved solute.
The rate of dissolution equals the rate of precipitation.
Crystallization:
The process by which solute molecules leave the solution and form crystals.
Can be induced by cooling the solution, adding a seed crystal, or evaporating some of the solvent.
Precipitation:
The process by which solute molecules leave the solution and form a solid phase.
Occurs when the concentration of solute exceeds its solubility.
Applications of Solubility
Environmental Chemistry:
The solubility of pollutants in water affects their transport and distribution in the environment.
Solubility determines whether contaminants can dissolve in groundwater or remain in soil.
Pharmaceuticals:
The solubility of drugs affects their absorption and bioavailability in the body.
Drug formulations are designed to optimize solubility and dissolution rates.
Food Chemistry:
The solubility of food components affects their taste, texture, and stability.
Solubility plays a role in the preparation and preservation of foods.
Expressing the Concentration of Solutions
Different Units of Concentration:
Percent by mass (% m/m)
Percent by volume (% v/v)
Mass/volume percent (% m/v)
Molarity (M)
Molality (m)
Parts per million (ppm)
Parts per billion (ppb)
Percent by Mass (% m/m):
The mass of the solute divided by the mass of the solution, multiplied by 100%.
Percent : by : Mass = (Mass : of : Solute / Mass : of : Solution) * 100\%
Percent by Volume (% v/v):
The volume of the solute divided by the volume of the solution, multiplied by 100%.
Percent : by : Volume = (Volume : of : Solute / Volume : of : Solution) * 100\%
Mass/Volume Percent (% m/v):
The mass of the solute divided by the volume of the solution, multiplied by 100%.
Mass/Volume : Percent = (Mass : of : Solute / Volume : of : Solution) * 100\%
Molarity (M):
The number of moles of solute per liter of solution.
Molarity = Moles : of : Solute / Liters : of : Solution
Commonly used in chemistry because it relates directly to the number of molecules of solute in a given volume of solution.
Molality (m):
The number of moles of solute per kilogram of solvent.
Molality = Moles : of : Solute / Kilograms : of : Solvent
Useful for experiments where temperature changes are significant, as molality is independent of volume changes caused by temperature.
Parts per Million (ppm) and Parts per Billion (ppb):
Used for very dilute solutions, such as pollutants in water.
ppm = (Mass : of : Solute / Mass : of : Solution) * 10^6
ppb = (Mass : of : Solute / Mass : of : Solution) * 10^9
Colligative Properties
Definition:
Properties of solutions that depend on the concentration of solute particles, but not on the nature of the solute.
Types of Colligative Properties:
Vapor pressure lowering
Boiling point elevation
Freezing point depression
Osmotic pressure
Vapor Pressure Lowering:
The vapor pressure of a solution is lower than that of the pure solvent.
Caused by the presence of solute particles, which reduce the number of solvent molecules that can escape into the gas phase.
Boiling Point Elevation:
The boiling point of a solution is higher than that of the pure solvent.
The elevation is proportional to the concentration of solute particles.
\Delta Tb = Kb * m
Where \Delta Tb is the boiling point elevation, Kb is the ebullioscopic constant, and m is the molality of the solution.
Freezing Point Depression:
The freezing point of a solution is lower than that of the pure solvent.
The depression is proportional to the concentration of solute particles.
\Delta Tf = Kf * m
Where \Delta Tf is the freezing point depression, Kf is the cryoscopic constant, and m is the molality of the solution.
Osmotic Pressure:
The pressure required to prevent the flow of solvent across a semipermeable membrane.
Proportional to the concentration of solute particles.
\Pi = iMRT
Where \Pi is the
Molecular Compounds:
Form solutions containing neutral molecules and do not conduct electricity.
Example: Sugar dissolved in water.
Ionic Compounds:
Form solutions containing ions and conduct electricity.
Example: Salt dissolved in water.
Electrolytes:
Substances that form conducting solutions are called electrolytes.
Strong Electrolytes:
Ionic compounds that dissociate completely into ions when dissolved in water.
Form solutions that are very good conductors of electricity.
Example: Sodium chloride (NaCl).
Weak Electrolytes:
Molecular compounds that form a small number of ions when dissolved in water.
Form solutions that are weak conductors of electricity.
Example: Acetic acid (CH3COOH).
Non-Electrolytes:
Substances that dissolve in water but do not produce ions.
Form solutions that do not conduct electricity.
Example: Sugar (C12H22O11).
Solubility
Definition:
The maximum quantity of solute that can dissolve in a certain quantity of solvent at a specified temperature.
Saturated Solution:
A solution containing the maximum quantity of solute.
Additional solute will not dissolve and will remain undissolved.
Unsaturated Solution:
A solution containing less than the maximum quantity of solute.
More solute can be dissolved.
Supersaturated Solution:
A solution containing more than the maximum quantity of solute.
This is a temporary, unstable condition. If disturbed the excess solute will precipitate out of the solution as crystals
Factors Affecting Solubility
Solute-Solvent Interactions:
"Like dissolves like": Polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes.
Example: Water (polar) dissolves sugar (polar), while gasoline (nonpolar) dissolves oil (nonpolar).
Temperature:
Generally, the solubility of solids in liquids increases with temperature.
The solubility of gases in liquids decreases with temperature.
Example: Carbonated beverages lose carbon dioxide (CO2) more readily when warm.
Pressure:
Has little effect on the solubility of solids and liquids.
Significantly affects the solubility of gases in liquids.
Henry’s Law: The solubility of a gas is directly proportional to its partial pressure above the solution.
Example: Carbonated beverages are bottled under high pressure to increase the solubility of carbon dioxide.
Solubility and Saturation
Dynamic Equilibrium:
In a saturated solution, a dynamic equilibrium exists between dissolved and undissolved solute.
The rate of dissolution equals the rate of precipitation.
Crystallization:
The process by which solute molecules leave the solution and form crystals.
Can be induced by cooling the solution, adding a seed crystal, or evaporating some of the solvent.
Precipitation:
The process by which solute molecules leave the solution and form a solid phase.
Occurs when the concentration of solute exceeds its solubility.
Applications of Solubility
Environmental Chemistry:
The solubility of pollutants in water affects their transport and distribution in the environment.
Solubility determines whether contaminants can dissolve in groundwater or remain in soil.
Pharmaceuticals:
The solubility of drugs affects their absorption and bioavailability in the body.
Drug formulations are designed to optimize solubility and dissolution rates.
Food Chemistry:
The solubility of food components affects their taste, texture, and stability.
Solubility plays a role in the preparation and preservation of foods.
Expressing the Concentration of Solutions
Different Units of Concentration:
Percent by mass (% m/m)
Percent by volume (% v/v)
Mass/volume percent (% m/v)
Molarity (M)
Molality (m)
Parts per million (ppm)
Parts per billion (ppb)
Percent by Mass (% m/m):
The mass of the solute divided by the mass of the solution, multiplied by 100%.
Percent : by : Mass = (Mass : of : Solute / Mass : of : Solution) * 100\%
Percent by Volume (% v/v):
The volume of the solute divided by the volume of the solution, multiplied by 100%.
Percent : by : Volume = (Volume : of : Solute / Volume : of : Solution) * 100\%
Mass/Volume Percent (% m/v):
The mass of the solute divided by the volume of the solution, multiplied by 100%.
Mass/Volume : Percent = (Mass : of : Solute / Volume : of : Solution) * 100\%
Molarity (M):
The number of moles of solute per liter of solution.
Molarity = Moles : of : Solute / Liters : of : Solution
Commonly used in chemistry because it relates directly to the number of molecules of solute in a given volume of solution.
Molality (m):
The number of moles of solute per kilogram of solvent.
Molality = Moles : of : Solute / Kilograms : of : Solvent
Useful for experiments where temperature changes are significant, as molality is independent of volume changes caused by temperature.
Parts per Million (ppm) and Parts per Billion (ppb):
Used for very dilute solutions, such as pollutants in water.
ppm = (Mass : of : Solute / Mass : of : Solution) * 10^6
ppb = (Mass : of : Solute / Mass : of : Solution) * 10^9
Colligative Properties
Definition:
Properties of solutions that depend on the concentration of solute particles, but not on the nature of the solute.
Types of Colligative Properties:
Vapor pressure lowering
Boiling point elevation
Freezing point depression
Osmotic pressure
Vapor Pressure Lowering:
The vapor pressure of a solution is lower than that of the pure solvent.
Caused by the presence of solute particles, which reduce the number of solvent molecules that can escape into the gas phase.
Boiling Point Elevation:
The boiling point of a solution is higher than that of the pure solvent.
The elevation is proportional to the concentration of solute particles.
\Delta Tb = Kb * m
Where \Delta Tb is the boiling point elevation, Kb is the ebullioscopic constant, and m is the molality of the solution.
Freezing Point Depression:
The freezing point of a solution is lower than that of the pure solvent.
The depression is proportional to the concentration of solute particles.
\Delta Tf = Kf * m
Where \Delta Tf is the freezing point depression, Kf is the cryoscopic constant, and m is the molality of the solution.
Osmotic Pressure:
The pressure required to prevent the flow of solvent across a semipermeable membrane.
Proportional to the concentration of solute particles.
\Pi = iMRT
Where \Pi is the
Solutions in which water is the solvent
Electrical Conductivity of Aqueous Solutions
Molecular Compounds:
Form solutions containing neutral molecules and do not conduct electricity.
Example: Sugar dissolved in water.
Ionic Compounds:
Form solutions containing ions and conduct electricity.
Example: Salt dissolved in water.
Electrolytes:
Substances that form conducting solutions are called electrolytes.
Strong Electrolytes:
Ionic compounds that dissociate completely into ions when dissolved in water.
Form solutions that are very good conductors of electricity.
Example: Sodium chloride (NaCl).
Weak Electrolytes:
Molecular compounds that form a small number of ions when dissolved in water.
Form solutions that are weak conductors of electricity.
Example: Acetic acid (CH3COOH).
Non-Electrolytes:
Substances that dissolve in water but do not produce ions.
Form solutions that do not conduct electricity.
Example: Sugar (C12H22O11).
Solubility
Definition:
The maximum quantity of solute that can dissolve in a certain quantity of solvent at a specified temperature.
Saturated Solution:
A solution containing the maximum quantity of solute.
Additional solute will not dissolve and will remain undissolved.
Unsaturated Solution:
A solution containing less than the maximum quantity of solute.
More solute can be dissolved.
Supersaturated Solution:
A solution containing more than the maximum quantity of solute.
This is a temporary, unstable condition. If disturbed the excess solute will precipitate out of the solution as crystals
Factors Affecting Solubility
Solute-Solvent Interactions:
"Like dissolves like": Polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes.
Example: Water (polar) dissolves sugar (polar), while gasoline (nonpolar) dissolves oil (nonpolar).
Temperature:
Generally, the solubility of solids in liquids increases with temperature.
The solubility of gases in liquids decreases with temperature.
Example: Carbonated beverages lose carbon dioxide (CO2) more readily when warm.
Pressure:
Has little effect on the solubility of solids and liquids.
Significantly affects the solubility of gases in liquids.
Henry’s Law: The solubility of a gas is directly proportional to its partial pressure above the solution.
Example: Carbonated beverages are bottled under high pressure to increase the solubility of carbon dioxide.
Solubility and Saturation
Dynamic Equilibrium:
In a saturated solution, a dynamic equilibrium exists between dissolved and undissolved solute.
The rate of dissolution equals the rate of precipitation.
Crystallization:
The process by which solute molecules leave the solution and form crystals.
Can be induced by cooling the solution, adding a seed crystal, or evaporating some of the solvent.
Precipitation:
The process by which solute molecules leave the solution and form a solid phase.
Occurs when the concentration of solute exceeds its solubility.
Applications of Solubility
Environmental Chemistry:
The solubility of pollutants in water affects their transport and distribution in the environment.
Solubility determines whether contaminants can dissolve in groundwater or remain in soil.
Pharmaceuticals:
The solubility of drugs affects their absorption and bioavailability in the body.
Drug formulations are designed to optimize solubility and dissolution rates.
Food Chemistry:
The solubility of food components affects their taste, texture, and stability.
Solubility plays a role in the preparation and preservation of foods.
Expressing the Concentration of Solutions
Different Units of Concentration:
Percent by mass (% m/m)
Percent by volume (% v/v)
Mass/volume percent (% m/v)
Molarity (M)
Molality (m)
Parts per million (ppm)
Parts per billion (ppb)
Percent by Mass (% m/m):
The mass of the solute divided by the mass of the solution, multiplied by 100%.
Percent : by : Mass = (Mass : of : Solute / Mass : of : Solution) * 100\%
Percent by Volume (% v/v):
The volume of the solute divided by the volume of the solution, multiplied by 100%.
Percent : by : Volume = (Volume : of : Solute / Volume : of : Solution) * 100\%
Mass/Volume Percent (% m/v):
The mass of the solute divided by the volume of the solution, multiplied by 100%.
Mass/Volume : Percent = (Mass : of : Solute / Volume : of : Solution) * 100\%
Molarity (M):
The number of moles of solute per liter of solution.
Molarity = Moles : of : Solute / Liters : of : Solution
Commonly used in chemistry because it relates directly to the number of molecules of solute in a given volume of solution.
Molality (m):
The number of moles of solute per kilogram of solvent.
Molality = Moles : of : Solute / Kilograms : of : Solvent
Useful for experiments where temperature changes are significant, as molality is independent of volume changes caused by temperature.
Parts per Million (ppm) and Parts per Billion (ppb):
Used for very dilute solutions, such as pollutants in water.
ppm = (Mass : of : Solute / Mass : of : Solution) * 10^6
ppb = (Mass : of : Solute / Mass : of : Solution) * 10^9
Colligative Properties
Definition:
Properties of solutions that depend on the concentration of solute particles, but not on the nature of the solute.
Types of Colligative Properties:
Vapor pressure lowering
Boiling point elevation
Freezing point depression
Osmotic pressure
Vapor Pressure Lowering:
The vapor pressure of a solution is lower than that of the pure solvent.
Caused by the presence of solute particles, which reduce the number of solvent molecules that can escape into the gas phase.
Boiling Point Elevation:
The boiling point of a solution is higher than that of the pure solvent.
The elevation is proportional to the concentration of solute particles.
\Delta Tb = Kb * m
Where \Delta Tb is the boiling point elevation, Kb is the ebullioscopic constant, and m is the molality of the solution.
Freezing Point Depression:
The freezing point of a solution is lower than that of the pure solvent.
The depression is proportional to the concentration of solute particles.
\Delta Tf = Kf * m
Where \Delta Tf is the freezing point depression, Kf is the cryoscopic constant, and m is the molality of the solution.
Osmotic Pressure:
The pressure required to prevent the flow of solvent across a semipermeable membrane.
Proportional to the concentration of solute particles.
\Pi = iMRT
Where \Pi is the
Solutions in which water is the solvent
Electrical Conductivity of Aqueous Solutions
Molecular Compounds:
Form solutions containing neutral molecules and do not conduct electricity.
Example: Sugar dissolved in water.
Ionic Compounds:
Form solutions containing ions and conduct electricity.
Example: Salt dissolved in water.
Electrolytes:
Substances that form conducting solutions are called electrolytes.
Strong Electrolytes:
Ionic compounds that dissociate completely into ions when dissolved in water.
Form solutions that are very good conductors of electricity.
Example: Sodium chloride (NaCl).
Weak Electrolytes:
Molecular compounds that form a small number of ions when dissolved in water.
Form solutions that are weak conductors of electricity.
Example: Acetic acid (CH3COOH).
Non-Electrolytes:
Substances that dissolve in water but do not produce ions.
Form solutions that do not conduct electricity.
Example: Sugar (C12H22O11).
Solubility
Definition:
The maximum quantity of solute that can dissolve in a certain quantity of solvent at a specified temperature.
Saturated Solution:
A solution containing the maximum quantity of solute.
Additional solute will not dissolve and will remain undissolved.
Unsaturated Solution:
A solution containing less than the maximum quantity of solute.
More solute can be dissolved.
Supersaturated Solution:
A solution containing more than the maximum quantity of solute.
This is a temporary, unstable condition. If disturbed the excess solute will precipitate out of the solution as crystals
Factors Affecting Solubility
Solute-Solvent Interactions:
"Like dissolves like": Polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes.
Example: Water (polar) dissolves sugar (polar), while gasoline (nonpolar) dissolves oil (nonpolar).
Temperature:
Generally, the solubility of solids in liquids increases with temperature.
The solubility of gases in liquids decreases with temperature.
Example: Carbonated beverages lose carbon dioxide (CO2) more readily when warm.
Pressure:
Has little effect on the solubility of solids and liquids.
Significantly affects the solubility of gases in liquids.
Henry’s Law: The solubility of a gas is directly proportional to its partial pressure above the solution.
Example: Carbonated beverages are bottled under high pressure to increase the solubility of carbon dioxide.
Solubility and Saturation
Dynamic Equilibrium:
In a saturated solution, a dynamic equilibrium exists between dissolved and undissolved solute.
The rate of dissolution equals the rate of precipitation.
Crystallization:
The process by which solute molecules leave the solution and form crystals.
Can be induced by cooling the solution, adding a seed crystal, or evaporating some of the solvent.
Precipitation:
The process by which solute molecules leave the solution and form a solid phase.
Occurs when the concentration of solute exceeds its solubility.
Applications of Solubility
Environmental Chemistry:
The solubility of pollutants in water affects their transport and distribution in the environment.
Solubility determines whether contaminants can dissolve in groundwater or remain in soil.
Pharmaceuticals:
The solubility of drugs affects their absorption and bioavailability in the body.
Drug formulations are designed to optimize solubility and dissolution rates.
Food Chemistry:
The solubility of food components affects their taste, texture, and stability.
Solubility plays a role in the preparation and preservation of foods.
Expressing the Concentration of Solutions
Different Units of Concentration:
Percent by mass (% m/m)
Percent by volume (% v/v)
Mass/volume percent (% m/v)
Molarity (M)
Molality (m)
Parts per million (ppm)
Parts per billion (ppb)
Percent by Mass (% m/m):
The mass of the solute divided by the mass of the solution, multiplied by 100%.
Percent : by : Mass = (Mass : of : Solute / Mass : of : Solution) * 100\%
Percent by Volume (% v/v):
The volume of the solute divided by the volume of the solution, multiplied by 100%.
Percent : by : Volume = (Volume : of : Solute / Volume : of : Solution) * 100\%
Mass/Volume Percent (% m/v):
The mass of the solute divided by the volume of the solution, multiplied by 100%.
Mass/Volume : Percent = (Mass : of : Solute / Volume : of : Solution) * 100\%
Molarity (M):
The number of moles of solute per liter of solution.
Molarity = Moles : of : Solute / Liters : of : Solution
Commonly used in chemistry because it relates directly to the number of molecules of solute in a given volume of solution.
Molality (m):
The number of moles of solute per kilogram of solvent.
Molality = Moles : of : Solute / Kilograms : of : Solvent
Useful for experiments where temperature changes are significant, as molality is independent of volume changes caused by temperature.
Parts per Million (ppm) and Parts per Billion (ppb):
Used for very dilute solutions, such as pollutants in water.
ppm = (Mass : of : Solute / Mass : of : Solution) * 10^6
ppb = (Mass : of : Solute / Mass : of : Solution) * 10^9
Colligative Properties
Definition:
Properties of solutions that depend on the concentration of solute particles, but not on the nature of the solute.
Types of Colligative Properties:
Vapor pressure lowering
Boiling point elevation
Freezing point depression
Osmotic pressure
Vapor Pressure Lowering:
The vapor pressure of a solution is lower than that of the pure solvent.
Caused by the presence of solute particles, which reduce the number of solvent molecules that can escape into the gas phase.
Boiling Point Elevation:
The boiling point of a solution is higher than that of the pure solvent.
The elevation is proportional to the concentration of solute particles.
\Delta Tb = Kb * m
Where \Delta Tb is the boiling point elevation, Kb is the ebullioscopic constant, and m is the molality of the solution.
Freezing Point Depression:
The freezing point of a solution is lower than that of the pure solvent.
The depression is proportional to the concentration of solute particles.
\Delta Tf = Kf * m
Where \Delta Tf is the freezing point depression, Kf is the cryoscopic constant, and m is the molality of the solution.
Osmotic Pressure:
The pressure required to prevent the flow of solvent across a semipermeable membrane.
Proportional to the concentration of solute particles.
\Pi = iMRT
Where \Pi is the
Solutions in which water is the solvent
Electrical Conductivity of Aqueous Solutions
Molecular Compounds:
Form solutions containing neutral molecules and do not conduct electricity.
Example: Sugar dissolved in water.
Ionic Compounds:
Form solutions containing ions and conduct electricity.
Example: Salt dissolved in water.
Electrolytes:
Substances that form conducting solutions are called electrolytes.
Strong Electrolytes:
Ionic compounds that dissociate completely into ions when dissolved in water.
Form solutions that are very good conductors of electricity.
Example: Sodium chloride (NaCl).
Weak Electrolytes:
Molecular compounds that form a small number of ions when dissolved in water.
Form solutions that are weak conductors of electricity.
Example: Acetic acid (CH3COOH).
Non-Electrolytes:
Substances that dissolve in water but do not produce ions.
Form solutions that do not conduct electricity.
Example: Sugar (C12H22O11).
Solubility
Definition:
The maximum quantity of solute that can dissolve in a certain quantity of solvent at a specified temperature.
Saturated Solution:
A solution containing the maximum quantity of solute.
Additional solute will not dissolve and will remain undissolved.
Unsaturated Solution:
A solution containing less than the maximum quantity of solute.
More solute can be dissolved.
Supersaturated Solution:
A solution containing more than the maximum quantity of solute.
This is a temporary, unstable condition. If disturbed the excess solute will precipitate out of the solution as crystals
Factors Affecting Solubility
Solute-Solvent Interactions:
"Like dissolves like": Polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes.
Example: Water (polar) dissolves sugar (polar), while gasoline (nonpolar) dissolves oil (nonpolar).
Temperature:
Generally, the solubility of solids in liquids increases with temperature.
The solubility of gases in liquids decreases with temperature.
Example: Carbonated beverages lose carbon dioxide (CO2) more readily when warm.
Pressure:
Has little effect on the solubility of solids and liquids.
Significantly affects the solubility of gases in liquids.
Henry’s Law: The solubility of a gas is directly proportional to its partial pressure above the solution.
Example: Carbonated beverages are bottled under high pressure to increase the solubility of carbon dioxide.
Solubility and Saturation
Dynamic Equilibrium:
In a saturated solution, a dynamic equilibrium exists between dissolved and undissolved solute.
The rate of dissolution equals the rate of precipitation.
Crystallization:
The process by which solute molecules leave the solution and form crystals.
Can be induced by cooling the solution, adding a seed crystal, or evaporating some of the solvent.
Precipitation:
The process by which solute molecules leave the solution and form a solid phase.
Occurs when the concentration of solute exceeds its solubility.
Applications of Solubility
Environmental Chemistry:
The solubility of pollutants in water affects their transport and distribution in the environment.
Solubility determines whether contaminants can dissolve in groundwater or remain in soil.
Pharmaceuticals:
The solubility of drugs affects their absorption and bioavailability in the body.
Drug formulations are designed to optimize solubility and dissolution rates.
Food Chemistry:
The solubility of food components affects their taste, texture, and stability.
Solubility plays a role in the preparation and preservation of foods.
Expressing the Concentration of Solutions
Different Units of Concentration:
Percent by mass (% m/m)
Percent by volume (% v/v)
Mass/volume percent (% m/v)
Molarity (M)
Molality (m)
Parts per million (ppm)
Parts per billion (ppb)
Percent by Mass (% m/m):
The mass of the solute divided by the mass of the solution, multiplied by 100%.
Percent : by : Mass = (Mass : of : Solute / Mass : of : Solution) * 100\%
Percent by Volume (% v/v):
The volume of the solute divided by the volume of the solution, multiplied by 100%.
Percent : by : Volume = (Volume : of : Solute / Volume : of : Solution) * 100\%
Mass/Volume Percent (% m/v):
The mass of the solute divided by the volume of the solution, multiplied by 100%.
Mass/Volume : Percent = (Mass : of : Solute / Volume : of : Solution) * 100\%
Molarity (M):
The number of moles of solute per liter of solution.
Molarity = Moles : of : Solute / Liters : of : Solution
Commonly used in chemistry because it relates directly to the number of molecules of solute in a given volume of solution.
Molality (m):
The number of moles of solute per kilogram of solvent.
Molality = Moles : of : Solute / Kilograms : of : Solvent
Useful for experiments where temperature changes are significant, as molality is independent of volume changes caused by temperature.
Parts per Million (ppm) and Parts per Billion (ppb):
Used for very dilute solutions, such as pollutants in water.
ppm = (Mass : of : Solute / Mass : of : Solution) * 10^6
ppb = (Mass : of : Solute / Mass : of : Solution) * 10^9
Colligative Properties
Definition:
Properties of solutions that depend on the concentration of solute particles, but not on the nature of the solute.
Types of Colligative Properties:
Vapor pressure lowering
Boiling point elevation
Freezing point depression
Osmotic pressure
Vapor Pressure Lowering:
The vapor pressure of a solution is lower than that of the pure solvent.
Caused by the presence of solute particles, which reduce the number of solvent molecules that can escape into the gas phase.
Boiling Point Elevation:
The boiling point of a solution is higher than that of the pure solvent.
The elevation is proportional to the concentration of solute particles.
\Delta Tb = Kb * m
Where \Delta Tb is the boiling point elevation, Kb is the ebullioscopic constant, and m is the molality of the solution.
Freezing Point Depression:
The freezing point of a solution is lower than that of the pure solvent.
The depression is proportional to the concentration of solute particles.
\Delta Tf = Kf * m
Where \Delta Tf is the freezing point depression, Kf is the cryoscopic constant, and m is the molality of the solution.
Osmotic Pressure:
The pressure required to prevent the flow of solvent across a semipermeable membrane.
Proportional to the concentration of solute particles.
\Pi = iMRT
Where \Pi is the