Solutions and Their Properties

Fundamental Definitions and Characteristics of Solutions

• Solution Definition: A solution is a homogeneous mixture consisting of two or more substances.

  • Homogeneous Properties: In a homogeneous mixture, the individual parts of the solution cannot be distinguished from one another. The mixture exists entirely in one phase.

  • Examples: Salt water and tea are common examples of solutions.

• Parts of a Solution:

  • Solute: The substance that is being dissolved (e.g., salt).

  • Solvent: The dissolving medium (e.g., water).

  • Aqueous Solutions: These are solutions where water is the solvent. Water is widely known as the "universal solvent."

  • Visual Characteristic: Liquid solutions are typically clear.

• Phases of Solutions:

  • Solutions can exist in all phases of matter, not just liquids.

  • Gaseous Solution: Air is a primary example.

  • Solid Solution: An alloy is a solid solution. An example is Brass, which is a combination of copper ($Cu$) and zinc ($Zn$).

  • Composition Rule: In solutions that are entirely gaseous or entirely liquid, the substance present in the smaller amount is designated as the solute.

  • Specific Classification: The transcript notes that seawater is considered a heterogeneous solution.

Heterogeneous Mixtures

• Suspension: A mixture that appears to be uniform while it is being actively stirred or agitated, but separates into different phases once the agitation ceases.

• Colloid: A heterogeneous mixture containing intermediate-sized particles that are evenly distributed through a dispersion medium.

• Brownian Motion:

  • Definition: The jerky, erratic movement observed in dispersed particles.

  • Function: This motion prevents colloids from settling out over time.

• Tyndall Effect:

  • Definition: The scattering of light by dispersed particles.

  • Occurrence: This is observable in clouds.

Solubility Terms and Principles

• Solubility Classifications:

  • Soluble: A substance that dissolves in another substance (e.g., sugar in water).

  • Insoluble: A substance that does not dissolve in another substance (e.g., sand in water).

  • Miscible: A term describing two liquids that are soluble in each other (e.g., alcohol and water).

  • Immiscible: A term describing two liquids that are insoluble in each other, which resultingly form distinct layers (e.g., oil and water).

• Solvation Mechanisms:

  • Solvation: The process by which solvent particles surround solute particles to form a solution.

  • Hydration: A specific type of solvation where the solvent used is water.

• The Principle of "Likes Dissolve Likes":

  • Polar substances dissolve in other polar substances (e.g., water and alcohol).

  • Ionic substances dissolve in polar substances (e.g., $NaCl$ dissolves in water).

  • Nonpolar substances dissolve in other nonpolar substances (e.g., oil and gasoline).

  • Predictive Examples:

    • Sugar (polar) and Sodium Fluoride ($NaF$) (ionic) are soluble in water (polar).

    • Octane (nonpolar) and benzene (nonpolar) are soluble in each other.

Factors Affecting the Rate of Solvation

• General Rule: To increase the rate at which a solute dissolves, one must increase the collisions between solute and solvent particles.

• Methods to Increase Rate:

  1. Stirring/Shaking: Because dissolving occurs at the surface, stirring allows more solvent to come into contact with the solute.

  2. Increasing Surface Area (Crushing): This exposes more of the solute to the solvent simultaneously.

  3. Increasing Solvent Temperature: As temperature increases, solvent particles move faster. This causes more frequent contact with the solute and provides the solvent particles with more energy to remove particles from the solid solute.

Thermodynamics of Solutions

• Heat of Solution:

  • Exothermic: A process that releases heat, causing the solution to get warm. An example is mixing Sodium Hydroxide ($NaOH$) with water.

  • Endothermic: A process that absorbs heat, causing the solution to feel cool. An example is the mixture of Barium Hydroxide and Ammonium Chloride.

Solubility Levels and Curves

• Solubility Definition: The maximum amount of solute that will dissolve in a given amount of solvent at a specific temperature.

• Saturation States:

  • Saturated: The solution contains the maximum amount of dissolved solute for the given amount of solvent at a specific temperature and pressure. These solutions exist in equilibrium ($NaCl(aq) \rightleftharpoons NaCl(s)$). Additional solute added will simply fall to the bottom.

  • Unsaturated: The solution contains less than the maximum amount of solute; more solute can still be dissolved.

  • Supersaturated: Contains more dissolved solute than a saturated solution at the same temperature. These are unstable and may recrystallize if disturbed. Sweet tea is an example (made by heating water, adding excess solute, and then cooling it down). Sodium acetate is another example.

• Factors Affecting Solubility Calculations:

  • Temperature and Solids: Solubility generally increases as temperature increases (e.g., sugar in water).

  • Temperature and Gases: Solubility decreases as temperature increases (e.g., Oxygen in water).

  • Solubility Curve Usage (Example based on graph data):

    1. At $90^{\circ}\text{C}$, approximately $54\,g$ of $KCl$ will dissolve in $100\,g$ water.

    2. $10\,g$ of $KClO_3$ will dissolve in $100\,g$ water at exactly $30^{\circ}\text{C}$.

    3. If $50\,g$ of $KCl$ is dissolved in $100\,g$ water at $90^{\circ}\text{C}$, the solution is considered unsaturated (because the limit is $54\,g$).

    4. Environmental impact example: If a company dumps warm water into a lake, the solubility of dissolved oxygen ($O_2$) decreases, which can suffocate fish and disrupt the food chain.

Henry's Law and Pressure

• Henry's Law: Relates the solubility of a gas to the partial pressure of that gas above the liquid.

• Soft Drink Example: To maximize solubility while making a soft drink:

  • To dissolve more sugar (solid solute): Increase temperature.

  • To dissolve more $CO_2$ (gaseous solute): Decrease temperature AND increase pressure.

Colligative Properties of Solutions

• Definition: Physical properties of solutions that are affected by the number of dissolved solute particles, rather than the identity of those particles.

• Four Main Colligative Properties:

  1. Vapor Pressure Lowering

  2. Boiling Point Elevation

  3. Freezing Point Depression

  4. Osmotic Pressure

• Electrolytes vs. Nonelectrolytes:

  • Electrolytes: Form ions in solution (ionic compounds, acids) and conduct electricity. They have a greater effect on colligative properties because they produce more particles.

    • Example: $1\,mol$ of $NaCl(s) \rightarrow 1\,mol\,Na^{+}(aq) + 1\,mol\,Cl^{-}(aq)$. Thus, $1\,mole$ of $NaCl$ yields $2\,moles$ of particles.

  • Nonelectrolytes: Do not ionize; they stay as molecules (covalent compounds) and do not conduct electricity.

    • Example: $1\,mol$ of Glucose ($C_6H_{12}O_6$) yields only $1\,mol$ of particles.

  • Comparison of Effect (1 Molar solutions):

    • Sucrose ($C_{12}H_{22}O_{11}$): $1$ particle.

    • Magnesium Nitrate ($Mg(NO_3)_2$): Forms $1\,Mg^{2+}$ and $2\,NO_3^{-}$ ($3$ particles).

    • Aluminum Bromide ($AlBr_3$): Forms $1\,Al^{3+}$ and $3\,Br^{-}$ ($4$ particles). This has the greatest effect because it yields the most particles.

• Vapor Pressure Lowering:

  • Adding a solute lowers the solvent's vapor pressure. Solute molecules occupy space near the surface, decreasing the number of solvent molecules that can escape into the gas phase.

• Boiling Point Elevation:

  • To boil, vapor pressure must equal atmospheric pressure. Because adding a solute lowers vapor pressure, the solution must be heated to a higher temperature to reach the boiling threshold ($T_b$). Example: Adding salt to water when cooking.

• Freezing Point Depression:

  • Adding a solute interferes with the attraction between solvent particles, which prevents them from shifting into the solid phase at the normal freezing point temperature. This results in a lower freezing point ($T_f$). Examples: Using salt on icy roads or making ice cream.

Measuring Solution Concentration Quantitatively

• Concentration: The measure of how much solute is dissolved in a specific amount of solvent or solution.

  • Concentrated: Large amount of solute.

  • Dilute: Small amount of solute.

• Molarity (M):

  • Formula: $\text{Molarity (M)} = \frac{\text{moles of solute (n)}}{\text{liters of solution (L)}}$

  • Units can be expressed as Molar, M, or $mol/L$.

  • Molarity is dependent on temperature.

  • Precision: Volumetric flasks are the best tools for preparing precise molar solutions.

  • Problem 1: Find molarity of $100\,mL$ solution with $0.075\,mol$ $NaCl$.

    • $0.100\,L$ is the volume.

    • $\text{M} = \frac{0.075\,mol}{0.100\,L} = 0.75\text{ M}$.

  • Problem 2: Find molarity of $45.0\,g$ $NaCl$ in $1\,L$.

    • Molar mass of $NaCl = 58.45\,g/mol$.

    • $\text{Moles} = \frac{45.0\,g}{58.45\,g/mol} = 0.769\,mol$.

    • $\text{M} = \frac{0.769\,mol}{1\,L} = 0.769\text{ M}$.

• Diluting Solutions:

  • Formula: $M_1V_1 = M_2V_2$

  • Problem: Volume of $12.0\text{ M } HCl$ needed to make $0.500\,L$ of $3.00\text{ M } HCl$.

    • $(12.0\text{ M}) \times V_1 = (3.00\text{ M}) \times (0.500\,L)$

    • $V_1 = \frac{3.00 \times 0.500}{12.0} = 0.125\,L = 125\,mL$.

• Molality (m):

  • Formula: $\text{molality (m)} = \frac{\text{moles of solute}}{\text{kilograms of solvent}}$

  • Molality is independent of temperature.

  • Problem 1: Grams of $KI$ in $500.0\,g$ water for a $0.060\text{ molal}$ solution. Answer: $5.0\,g\text{ KI}$.

  • Problem 2: Calculate molality of $10.0\,g\text{ NaCl}$ in $600\,g$ water. Answer: $0.3\text{ m}$.

  • Problem 3: Molality of solid solution with $0.125\,g$ Chromium and $81.3\,g$ Iron. Answer: $0.3\text{ m}$.

• Mole Fraction (X):

  • Formula: $X = \frac{\text{moles of component}}{\text{total moles of solution}}$

  • Problem 1: Mole fractions of ethyl alcohol ($C_2H_5OH$) and water when mixing $50.0\,g$ of each.

  • Problem 2: A gas mixture has $CH_4$ ($0.510$), $C_2H_6$ ($0.431$), $C_3H_8$ ($0.011$), and $C_4H_{10}$ ($0.013$). To find the mole fraction of the remaining gas, acetylene ($C_2H_2$), subtract the sum from $1$.

Questions and Discussion (Self-Quiz Review)

• Questions from Slide 40: Basic Definitions

  • Q1: A solution is a mixture of 2 or more . Answer: Homogeneous / substances.

  • Q2: Homogeneous means that you cannot distinguish the _ and it is in one _. Answer: Parts / phase.

  • Q3: Name solutes and solvent in sweet tea. Answer: Solute = Sugar/Tea; Solvent = Water.

  • Q4: _ is the universal solvent. Answer: Water.

  • Q5: Example of gaseous solution. Answer: Air.

  • Q6: Alloy composition and example. Answer: Solid solution of metals / Brass.

  • Q7: Solvent in aqueous solutions. Answer: Water.

• Questions from Slide 41: Solubility Principles

  • Q1: Substance that dissolves in another. Answer: Soluble.

  • Q2: Sand and water are _. Answer: Insoluble.

  • Q3: Two liquids that are soluble. Answer: Miscible.

  • Q4: Oil and water are _. Answer: Immiscible.

  • Q5: Solvation with water. Answer: Hydration.

  • Q6: "Likes dissolve likes" meaning. Answer: Polar dissolves polar and ionic; nonpolar dissolves nonpolar. Oil/water insoluble because water is polar and oil is nonpolar.

• Questions from Slide 42: Processes and States

  • Q1: 3 things to increase solvation. Answer: Stirring, increasing surface area, increasing temperature.

  • Q2: Calcium chloride + water getting warm. Answer: Exothermic.

  • Q3: Solution with max solute. Answer: Saturated.

  • Q4: Solution that can dissolve more. Answer: Unsaturated.

  • Q5: How to make supersaturated solution. Answer: Heat it, add more solute, then cool it down.