Introduction to Solutions

Introduction to Solutions

• A solution is a homogeneous mixture of two or more substances.
• Homogeneous mixtures have a uniform composition throughout.
• Heterogeneous mixtures have a non-uniform composition.
• Common examples include sugar dissolved in water, salt dissolved in water, and lemonade.

Components of Solutions

• Solute: The component present in lesser quantity.
• Solvent: The component present in greater quantity.
• Examples of solute/solvent combinations include sugar and water (sugar is the solute, water is the solvent).

Types of Solutions

• Binary Solutions: Solutions made of two components.
• Ternary Solutions: Solutions made of three components.
• Quaternary Solutions: Solutions made of four components.
• Poly Solutions: Solutions made of multiple components.
• Diluted Solutions: Solutions with a low concentration of solute.
• Concentrated Solutions: Solutions with a high concentration of solute.
• Unsaturated Solutions: Solutions where more solute can be dissolved at a given temperature.
• Saturated Solutions: Solutions where no more solute can be dissolved at a given temperature.
• Aqueous Solutions: Solutions where water is the solvent.
• Non-Aqueous Solutions: Solutions where water is not the solvent.

Classifying Solutions: Solute and Solvent States

• Solid in Liquid: A solid solute dissolved in a liquid solvent. Examples include sugar in water and salt in water.
• Liquid in Liquid: A liquid solute dissolved in a liquid solvent. Examples include alcohol in water and vinegar.
• Gas in Liquid: A gaseous solute dissolved in a liquid solvent. Examples include carbon dioxide in water (soda) and oxygen in water (aquatic life).
• Solid in Gas: A solid solute dissolved in a gaseous solvent. Examples include smoke.
• Liquid in Gas: A liquid solute dissolved in a gaseous solvent. Examples include fog and clouds.
• Gas in Gas: A gaseous solute dissolved in a gaseous solvent. Example: Air.
• Solid in Solid: A solid solute dissolved in a solid solvent. Examples include alloys, such as brass and bronze (brass is an alloy of copper and zinc).
• Liquid in Solid: A liquid solute dissolved in a solid solvent. Examples include amalgams, such as mercury in sodium.
• Gas in Solid: A gaseous solute dissolved in a solid solvent. Examples include hydrogen dissolved in palladium.

Expressing Concentration Qualitatively

• Qualitative: Describing a solution's concentration based on visual observations.
• Quantitative: Describing a solution's concentration with numerical values, units, and formulas.

Expressing Concentration Quantitatively

• Mass Percentage: The mass of the component divided by the mass of the solution, multiplied by 100.
• Volume Percentage: The volume of the component divided by the volume of the solution, multiplied by 100.

Concentration Terms

• Volume by Volume Concentration (v/v): Expressed as the volume of solute per 100 mL of solution. Example: 15 mL of solute in 100 mL of solution.
• Mass by Volume Concentration (m/v): Represents the mass of solute (in grams) per 100 mL of solution. Example: 20 g of solute in 100 mL of solution.
• Volume by Mass Concentration (v/m): Indicates the volume of solute (in mL) per 100 g of the solution. Example: 20 mL of solute in 100 g of solution.

Mole Fraction

• Defined as the ratio of the number of moles of a component to the total number of moles in the solution.
• Represented by the symbol 'χ'.
• Formula: χ_A = Number of moles of A / Total number of moles in the solution.
• Key Properties:
  • Unitless (moles cancel out in the ratio).
  • The mole fraction of the solution is always equal to one (sum of mole fractions of each component).
  • Independent of temperature (as it's a mass-by-mass concept).

Molarity

• Defined as the number of moles of solute dissolved per liter of solution.
• Represented by the symbol 'M'.
• Formula: Molarity (M) = Number of moles of solute / Volume of solution (in liters)
• Unit: Moles per liter (mol/L).
• Memorization Tip: "Washing Machine 1000 Milliliter" (W = mass of solute, M = molar mass of solute, 1000 = conversion factor).
• Dilution Equation: m₁v₁ = m₂v₂ (initial molarity & volume vs. final molarity & volume)
• Addition Equation: m₁v₁ + m₂v₂ = m₃(v₁ + v₂) (molarities of solutions and their volumes)

Molality

• Defined as the number of moles of solute dissolved per kilogram of solvent.
• Represented by the symbol 'm'.
• Formula: Molality (m) = Number of moles of solute / Mass of solvent (in kilograms).
• Unit: Moles per kilogram (mol/kg).
• Memorization Tip: "Washing Machine 1000 Washing Ton" (mass of solute, molar mass of solute, conversion factor, mass of solvent, grams to kg conversion).
• Key Properties:
  • A mass-by-mass concept and hence independent of temperature.

Key Differences between Molarity and Molality

• Definition: Molarity considers the volume of the solution, while molality focuses on the mass of the solvent.
• Units: Molarity is expressed in mol/L, while molality is in mol/kg.
• Temperature Dependence: Molarity varies with temperature due to changes in volume, while molality remains unchanged.

Types of Solutions (Solid Solution)

• Solid Solution: A homogeneous mixture where the solvent is a solid. Examples include alloys like brass (copper and zinc).

Assertion and Reason

• Assertion: A statement proposed to be true.
• Reason: A statement providing explanation for the assertion.

Analyzing Assertion and Reason Questions:

• A: Both assertion and reason correct; reason is correct explanation of assertion.
• B: Both assertion and reason correct; reason is not the correct explanation of assertion.
• C: Assertion correct, reason incorrect.
• D: Assertion incorrect, reason correct.

Molality and Temperature

• Molality of a substance does not change with temperature.

Volume of Solution and Temperature

• Volume of solution changes with temperature.

Defining Key Terms

• Molality: Number of moles of solute dissolved in 1 kilogram of solvent.
• Mole Fraction: Ratio of moles of a component to the total moles in the solution.
• Binary Solution: Contains two components; water may be solvent.
• Molarity: Number of moles of solute dissolved in 1 liter of solution.

Distinguishing between Molarity and Molality

• Molarity: moles of solute per liter of solution. Molality: moles of solute per kilogram of solvent.
• Molarity is affected by temperature, molality is not.

Converting Molarity to Molality

• Formula: M = (d * m * B) / (1 + (m * MB))
  • M = Molarity
  • d = Density
  • m = Molality
  • MB = Molar Mass of the solute in kilograms

Comparing Molar and Molal Solutions

• 1 molar sugar solution has 1 mole of sugar in 1 liter of solution.
• 1 molal sugar solution has 1 mole of sugar in 1 kilogram of solvent.
• Molar solutions typically more concentrated than molal solutions due to temperature insensitivity.

Defining Solubility

• Solubility is the maximum amount of solute that can dissolve in 100 grams of solvent at a given temperature.

Factors Affecting Solubility

• Nature of Solute and Solvent: "Like dissolves like." Polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents. Example: Salt (NaCl) dissolves in water (H₂O) due to polarity but not in oil.
• Temperature: Solubility of solids in liquids generally increases with increasing temperature. Endothermic reactions increase solubility with temperature, exothermic reactions decrease.
• Pressure: Pressure has a negligible effect on the solubility of solids in liquids.

Solubility of Gases in Liquids

• Nature of the Gas: Gases easily liquefied (e.g., CO₂, NH₃, SO₂) tend to be more soluble than gases that don't easily liquefy.
• Pressure: The solubility of gases in liquids increases with increasing pressure (Henry's Law).
  • Henry's Law: Solubility of a gas is proportional to its partial pressure above the liquid.
  • Formula: P = kC
    • P = partial pressure of the gas.
    • k = Henry's Law constant.
    • C = concentration of the dissolved gas.
• Temperature: Solubility of gases in liquids generally decreases with increasing temperature.

Applications of Henry's Law

• High pressure in soda bottles increases CO₂ solubility. Opening a soda releases pressure, reducing CO₂ solubility and releasing bubbles.
• Deep-sea divers use gas mixtures without nitrogen to avoid nitrogen narcosis (excessive nitrogen solubility in blood).

Pressure and Solubility

• Solubility of gases decreases with decreasing pressure.
• Scuba divers experience increased gas solubility at higher pressures and rapid release at lower pressures.
• Use of helium in scuba diving equipment reduces blood nitrogen solubility, preventing decompression sickness (the bends).

High Altitude Affects

• Lower atmospheric pressure leads to lower oxygen intake & results in anoxia (lack of oxygen).
• Symptoms such as fatigue, impaired cognitive function, and difficulty breathing.

Vapor Pressure

• Vapor pressure is the pressure exerted by the vapor of a liquid in equilibrium with its liquid phase.
• Factors Affecting Vapor Pressure:
  • Intermolecular interactions (weaker = higher vapor pressure)
  • Temperature (higher = higher vapor pressure)
  • Boiling point (higher = lower vapor pressure)
  • Volatile solutes (increase vapor pressure)
  • Non-volatile solutes (decrease vapor pressure)

Raoult's Law

• Raoult's Law states that the partial vapor pressure of a component in a solution is directly proportional to its mole fraction.
• Partial Vapor Pressure of Component A: (P_A = P^o_A \times X_A)
  • (P_A) = Partial vapor pressure of component A in the solution.
  • (P^o_A) = Vapor pressure of pure component A.
  • (X_A) = Mole fraction of component A in the solution.

Types of Solutions

• Ideal Solutions: Obey Raoult's Law at all concentrations and temperatures; similar intermolecular interactions; no volume/enthalpy change upon mixing. Examples: Hexane and Heptane, Benzene and Toluene.
• Non-Ideal Solutions: Deviate from Raoult's Law; different intermolecular forces; volume/enthalpy change upon mixing. Examples: Water and Ethanol, Chloroform and Acetone.

Positive Deviation From Raoult's Law

• Solute-solvent interactions weaker than solute-solute and solvent-solvent interactions. Higher vapor pressure than predicted by Raoult's Law. Volume increases, enthalpy positive (endothermic). Example: Water and Ethanol.

Negative Deviation From Raoult's Law

• Solute-solvent interactions stronger than solute-solute and solvent-solvent interactions. Lower vapor pressure than predicted. Volume decreases, enthalpy negative (exothermic). Example: Acetone and Chloroform.