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Chapter 11. Solutions

Recommended Article : 【Chemistry】 Chemistry Table of Contents


1. Classification of Substances

2. Overview

3. Solubility

4. General Properties of Solutions

5. Colloids



## 1. Classification of Substances

⑴ Pure Substances : Classified into elements and compounds

⑵ Mixtures : Classified into homogeneous mixtures and heterogeneous mixtures

⑶ Elements

① Substances composed of only one element

② Example : Copper (Cu), Nitrogen (N2), Iron (Fe), Diamond (C), Aluminum (Al)

⑷ Compounds

① Substances composed of two or more different elements in a fixed ratio

② Example : Carbon Dioxide (CO2), Copper Sulfate (CuSO4), Water (H2O)

⑸ Homogeneous Mixtures (Solutions)

① Mixtures where two or more pure substances are uniformly mixed, and the composition is the same throughout

② Example : Air, Sugar Solution

⑹ Heterogeneous Mixtures

① Mixtures where two or more pure substances are unevenly mixed, and the composition varies depending on the portion taken

② Example : Muddy Water, Milk



2. Overview

⑴ Formation of Solutions

① Solvent : Dissolving substance

○ Example : Water in Saltwater

○ Example : Component with a larger quantity in Ethanol solution

② Solute : Dissolved substance

○ Example : Salt in Saltwater

○ Example : Component with a smaller quantity in Ethanol solution

③ Solvation : The phenomenon where solute dissolves in solvent

○ Solvent surrounds solute particles during solution formation

○ When water is the solvent, it’s called hydration

⑵ Concentration of Solutions

① Mass%, Volume%

② Molarity (M) : Moles of solute per liter of solvent

③ Molality (m) : Moles of solute per kilogram of solvent

④ Mole Fraction (x) : Ratio of the number of particles of solvent to that of solute

⑤ ppm, ppb

⑶ Types of Solutions

① Unsaturated Solution : A solution that can dissolve more solute

② Saturated Solution : A solution with the maximum amount of solute dissolved

③ Supersaturated Solution : A solution containing more solute than a saturated solution, causing precipitation of solute

⑷ Electrolytes and Non-Electrolytes

① Electrolyte : A substance that conducts electricity when dissolved in water

○ Electrolytes are composed of particles with opposite charges that dissociate in solution

○ Example : Salt, Copper(II) Sulfate

② Non-Electrolyte : A substance that does not conduct electricity when dissolved in water

○ Example : Distilled Water, Ethanol, Sugar Solution

③ Strong Electrolyte : A substance that ionizes extensively in solution

○ High ion concentration characterizes strong electrolytes

○ Example : Ionic compounds (NaCl), Strong acids (HCl), Strong bases (NaOH), etc.

④ Weak Electrolyte : A substance that ionizes to a lesser extent in solution

○ Low ion concentration characterizes weak electrolytes

○ Example : Weak acids (CH3COOH), Weak bases (NH4OH), etc.



3. Solubility

⑴ Basic Principles of Solubility

① Enthalpy of Solution

○ ΔHsolution = ΔHlattice + ΔHhydration : M+(g) + N-(g) → M+(aq) + N-(aq)

○ ΔHlattice, lattice (< 0) : M+(g) + N-(g) → MN(s)

○ Smaller metal ion radii lead to greater absolute values of lattice enthalpy

○ ΔHhydration, hydration (0 ± ) : MN(s) → M+(aq) + N-(aq)

② Hydration : Surrounding of solute particles by water molecules in a solution

○ Ionic solutes : Smaller ions and higher charges lead to stronger hydration

○ Hydration leads to ordered water molecules, resulting in a decrease in solvent entropy

○ Hydration disrupts the regular structure of solute particles, leading to an increase in solute entropy

○ Hydration is an entropy-increasing process

○ Entropy change upon evaporation of pure solvent = Gas entropy - Solvent entropy > Gas entropy - Solution entropy = Solution’s evaporation entropy

⑵ Temperature and Solubility

① Solubility of Solids : Increases with temperature, endothermic reaction (ΔHsolution > 0)

② Solubility of Gases (applies to some solids as well) : Decreases with temperature, exothermic reaction (ΔHsolution < 0)

③ Henry’s Law

○ **Formula: **Concentration (C) in solution = Henry’s constant (K) × Partial pressure (P) of gas

○ Gas solubility is directly proportional to the partial pressure of the gas

○ Applicable mostly to gases with low solubility

○ Derivation of Henry’s Law

⑶ Polarity and Solubility : “Like dissolves like”

① Polar solvents dissolve polar solutes effectively

② Nonpolar solvents dissolve nonpolar solutes effectively



4. General Properties of Solutions

⑴ Definition

① Properties related to the number of particles of solute, regardless of the type of solute

② Under the assumption of ideal solutions, entropy is the underlying factor for all general properties

⑵ Van’t Hoff Factor : denoted as i

① Definition

② Example : When NaCl(s) is dissolved in water, it completely dissociates into Na+(aq) and Cl-(aq), so the Van’t Hoff factor is 2

③ Real Solutions : As concentration becomes higher and the ion charges are larger, more ion pairs form in the solution, causing i to decrease

④ Ionization Degree and Van’t Hoff Factor

⑶ Lowering of Vapor Pressure

① Raoult’s Law

Content 1. In equilibrium, the partial pressure of each component is proportional to the mole fraction of the component in the liquid mixture

Content 2. The volume of the mixture is equal to the sum of the volumes of each component before mixing

Content 3. Intermolecular interactions in the mixture are the same as those between pure components

○ Relevant equation

② When both the solvent and solute form vapor pressure

③ Ideal Solutions (↔ Non-ideal Solutions) : Solutions that satisfy Raoult’s Law

Figure 1. Raoult’s Law and Positive Deviations, Negative Deviations [Footnote:1]

○ Solution with P = Super-saturated Solution with P: ΔH Solubilization = 0

○ Solution with P < Super-saturated Solution with P: Solvent dislikes vaporization, vapor pressure decreases ⇔ Strong solute-solvent interaction ⇔ ΔH Solubilization < 0

○ Solution with P > Super-saturated Solution with P: Solvent tends to vaporize, vapor pressure increases ⇔ Weak solute-solvent interaction ⇔ ΔH Solubilization > 0

④ Fractional Distillation of Ideal Solutions: Ideal solutions have different compositions in liquid and gas phases → Repetitive vaporization and condensation allow for the extraction of pure components

⑷ Boiling Point Elevation: 1st approximation, valid for dilute solutions and small temperature changes

⑸ Freezing Point Depression: 1st approximation, valid for dilute solutions and small temperature changes

① Molal Depression Constant Kf is a property of the solvent, not the solute

⑹ Osmotic Phenomenon (osmosis)

① Osmosis

○ Phenomenon where solvent molecules move due to differences in concentration between two solutions separated by a semipermeable membrane (diffusion of free molecules)

○ Osmotic Pressure: Pressure needed to prevent osmosis. Formulated using van ‘t Hoff’s law

π = CRT × i

○ Similarity of osmotic pressure formula to ideal gas equation: Due to the dilute nature of solute molecules in solution, they behave like an ideal gas

② Reverse Osmosis

○ When pressure greater than osmotic pressure is applied, water moves from high concentration to low concentration

○ Theoretical pressure needed in reverse osmosis = Pressure needed to establish equilibrium when some water is removed and osmotic pressure increases

Example



5. Colloids

⑴ Overview

① Definition: Particles ranging from 1 μm to 1,000 μm dispersed in a gas or liquid medium

② Particles dispersed in a gas are called aerosols

⑵ Properties: Particle size

① Tyndall Phenomenon

○ Definition: The path of light becomes visible due to particles within the colloid

Rayleigh Scattering and Tyndall Phenomenon both scatter specific wavelengths, but conditions and paths of scattering differ

○ No accurate mathematical formula describing Tyndall Phenomenon exists to date

② Dialysis: Diffusion of substances through a semipermeable membrane

○ Dialysis solution concentration > Solution concentration: Substance moves from dialysis solution to solution

○ Dialysis solution concentration = Solution concentration: No substance movement

○ Dialysis solution concentration < Solution concentration: Substance moves from solution to dialysis solution

○ Renal hemodialysis is a representative application

③ Adsorption

④ Brownian Motion: Random motion of colloidal particles in a liquid medium

○ Einstein analyzed it mathematically and received the Nobel Prize

○ Application: Dynamic Light Scattering (DLS) measures scattered light reflecting Brownian motion to determine particle size

○ Draws 1st time-intensity plot

Figure. 2. DLS setup and intensity-time plot

○ Compares two intensity-time plots using cross-correlation to draw delayed time-correlation function plot

Figure. 3. Delayed-correlation function plot

○ Analyzes exponential decay curve to determine translational diffusion coefficient Dt

○ Calculates hydrodynamic diameter Dh according to Stokes-Einstein law

⑶ Properties: Charge

Electrophoresis

② Zeta Potential

Figure. 4. Concept of Zeta Potential]

Figure. 5. Zeta potential patterns based on particle type

○ Background theory

○ Oppositely charged particles gather around the charged particle to form a primary shell

○ Stern Layer: Primary shell

○ Polar particles gather around the Stern Layer to form a secondary shell

○ Stern Layer moves with the particle

○ Double Layer: Secondary shell, also known as DEL (Double Electrode Layer)

○ Secondary shell can be of the same or different polarity as the initial particle

○ Movement influenced by solvent more than particles

○ Zeta Potential

○ Definition: Potential at the surface of the double layer

○ Measurable unlike surface potential or stern potential

○ Zeta potential can be measured by observing the difference in particle movement speed when applying potential

○ Utility 1: Measures particle polarity

○ Utility 2: Reflects both charge state and particle dispersion

○ If particles have the same charge and large zeta potential, they won’t aggregate

○ Absolute value of zeta potential exceeding 30 mV prevents aggregation between particles

○ Due to the emergence of repulsive forces

③ Aggregation (Flocculation)

○ Colloidal particles aggregate due to electrostatic forces between particles, forming small clumps

○ Differentiated from coagulation by the precipitation of solute and solvent together

○ Schulze-Hardy (S-H) rule: Aggregation strength proportional to electrostatic forces between solutes. Proposed around 125 years ago

○ Factor 1: Ion concentration: Aggregation occurs faster with more positive and negative solutes

○ Factor 2: Higher charge of solute increases aggregation strength

④ Salting-out

Figure. 6. Salting-out phenomenon

○ Differentiated from flocculation by the precipitation of solute only

○ Low concentration salt: Solubility increases with added salt (salting-in) due to salt-induced alteration aiding water penetration

○ High concentration salt: Solubility decreases with added salt (salting-out) as salts reduce interaction between substance and water

○ Peak solubility values increase with higher target substance concentration

○ Initial means of purification, ammonium sulfate used extensively

○ Application: Used in making tofu by adding a coagulant (MgCl2)



Input: 2018.12.30 20:39

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