Solutions & Colloids

Introduction: Mixtures and Dispersions

Many of the substances we encounter daily are mixtures, not pure compounds. Blood is a complex mixture; seawater contains dissolved salts and gases; soil is a colloid of particles suspended in air and water. Understanding solutions and colloids is essential for medicine, industry, and everyday applications. This chapter explores the nature of mixtures, how substances dissolve, how to measure concentration, and the fascinating properties of colloidal systems.

Part 1: Solutions

Definition and Components

A solution is a homogeneous mixture of two or more substances where one (solute) is dissolved in another (solvent).

Components:

• Solvent: The substance present in larger amount; dissolving medium (usually liquid)
• Solute: The substance present in smaller amount; dissolved substance (can be solid, liquid, or gas)

Example: Saltwater

• Solvent: Water (H₂O)
• Solute: Sodium chloride (NaCl)

Types of Solutions

Unsaturated Solution

• Contains less solute than the maximum amount that can dissolve
• More solute can still dissolve
• Example: 5g of salt dissolved in 100 mL of water (when maximum is 36g)

Saturated Solution

• Contains the maximum amount of solute that can dissolve at that temperature
• No more solute can dissolve without changing conditions
• Example: 36g of NaCl dissolved in 100 mL of water at 20°C

Supersaturated Solution

• Contains MORE solute than the maximum amount that should dissolve at that temperature
• Unstable; crystallization occurs easily when disturbed
• Example: Honey often becomes supersaturated; crystallizes over time

[Memory Hook] Unsaturated < saturated < supersaturated (in terms of solute concentration)

Solubility and Factors Affecting It

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

Factors affecting solubility:

1. Temperature

• Most solids: Solubility increases with temperature (endothermic dissolution)

  • Example: Sugar solubility at 20°C: 200g/100mL; at 100°C: 487g/100mL

• Some salts: Solubility decreases with temperature (Ca(OH)₂, Ce₂(SO₄)₃)

• Gases: Solubility decreases with temperature (inverse relationship)

  • Hot water holds less dissolved oxygen than cold water
  • Why: Gas molecules escape faster at high temperatures

2. Pressure

• Effect on solids/liquids: Negligible (incompressible)
• Effect on gases: Solubility increases with pressure (Henry's Law)
  • Carbonated drinks have CO₂ dissolved under high pressure
  • Opening the bottle reduces pressure; CO₂ escapes as bubbles

3. Nature of Solute and Solvent

• "Like dissolves like": Polar solvents dissolve polar/ionic solutes; nonpolar solvents dissolve nonpolar solutes
  • Water (polar) dissolves salt (ionic), sugar (polar) but not oil (nonpolar)
  • Oil (nonpolar) dissolves in nonpolar solvents (gasoline, ether)

[Exam Trap] Students often think "high temperature always increases solubility." Not true for gases and a few salts. Check the specific solute.

Part 2: Concentration Methods

1. Molarity (M)

Definition: Number of moles of solute per liter of solution.

Formula: M = moles of solute / liters of solution

Example:

• Dissolve 5.85 g of NaCl (molar mass = 58.5 g/mol) in water to make 500 mL of solution
• Moles of NaCl = 5.85 / 58.5 = 0.1 mol
• Volume = 500 mL = 0.5 L
• Molarity = 0.1 mol / 0.5 L = 0.2 M

Advantages: Most common in chemistry; used for stoichiometry

Disadvantage: Changes slightly with temperature (volume of solution changes)

2. Molality (m)

Definition: Number of moles of solute per kilogram of solvent.

Formula: m = moles of solute / kg of solvent

Example:

• 5.85 g NaCl dissolved in 100 g water
• Moles of NaCl = 0.1 mol
• Mass of water = 100 g = 0.1 kg
• Molality = 0.1 mol / 0.1 kg = 1 m

Advantages: Unaffected by temperature (mass doesn't change); useful for colligative properties

3. Normality (N)

Definition: Number of equivalents of solute per liter of solution.

Formula: N = (equivalents of solute) / (liters of solution)

Equivalent: Number of moles of H⁺ (or OH⁻, or electrons) a substance can donate/accept

Example:

• H₂SO₄ has 2 H⁺ ions per molecule; 1 mole H₂SO₄ = 2 equivalents
• 0.5 M H₂SO₄ = 1 N H₂SO₄

Use: Acid-base titrations; redox reactions

4. Percent Composition by Mass

Definition: Mass of solute as percentage of total solution mass.

Formula: % mass = (mass of solute / mass of solution) × 100%

Example:

• 10 g salt dissolved in 90 g water (total = 100 g)
• % mass = (10 / 100) × 100% = 10%

5. Parts Per Million (ppm) and Parts Per Billion (ppb)

Definition: Concentration for very dilute solutions.

Formula:

• ppm = (mass of solute / mass of solution) × 10⁶
• ppb = (mass of solute / mass of solution) × 10⁹

Use: Measuring pollutants, trace elements in water

Example: If drinking water has 2 ppm lead, it means 2 grams of lead per 1,000,000 grams of water.

[Memory Hook] Molarity (M) = moles/liter. Molality (m) = moles/kg. Normality (N) = equivalents/liter. ppm = 1 part per million

Part 3: Colligative Properties

Colligative properties depend on the number of solute particles, not their identity.

1. Vapor Pressure Lowering

Raoult's Law: Vapor pressure of a solution < vapor pressure of pure solvent.

Explanation: Solute particles occupy surface; fewer solvent molecules can escape as vapor.

Effect: Adding solute → lower vapor pressure → solution evaporates more slowly than pure solvent

2. Boiling Point Elevation

Definition: Solution has a higher boiling point than pure solvent.

Reason: Lower vapor pressure means more heat is needed for solution to reach atmospheric pressure.

Formula: ΔT_b = K_b × m

Where:

• ΔT_b = elevation in boiling point
• K_b = ebullioscopic constant (specific to solvent)
• m = molality of solute

Example: 1 molar aqueous solution boils at ~100.5°C (not 100°C)

3. Freezing Point Depression

Definition: Solution has a lower freezing point than pure solvent.

Reason: Solute particles interfere with crystallization; more cooling is needed to freeze.

Formula: ΔT_f = K_f × m

Where:

• ΔT_f = depression in freezing point
• K_f = cryoscopic constant
• m = molality of solute

Example: 1 molar aqueous solution freezes at ~-1.86°C (not 0°C)

Application: Antifreeze in car radiators (ethylene glycol); salt on icy roads (lowers freezing point, preventing ice formation)

4. Osmotic Pressure

Definition: Pressure needed to prevent water from osmosing into a solution.

Osmosis: Water molecules move across a semipermeable membrane toward the solution (trying to dilute it).

Formula: Π = iMRT

Where:

• Π = osmotic pressure
• i = van 't Hoff factor (number of particles per molecule)
• M = molarity
• R = gas constant
• T = temperature in Kelvin

Biological application:

• Isotonic solution: Same solute concentration as cell cytoplasm (no net osmosis)
• Hypertonic solution: Higher solute concentration; water leaves cell (cell shrinks)
• Hypotonic solution: Lower solute concentration; water enters cell (cell swells)

[Memory Hook] Colligative properties: particle count matters, not particle identity. Vapor pressure ↓, boiling point ↑, freezing point ↓, osmotic pressure ↑

Part 4: Colloids

Definition and Particle Size

A colloidal solution (or colloid) is a heterogeneous mixture where particles (1 nm – 1000 nm) are dispersed in a medium but don't settle (unlike suspensions).

Comparison of mixtures:

Mixture Type	Particle Size	Visibility	Settlement	Passes Filter
Solution	< 1 nm	Transparent	No	Yes
Colloid	1–1000 nm	Opaque/turbid	No	No (usually)
Suspension	> 1000 nm	Opaque	Yes (eventually)	No

Types of Colloids

Based on dispersed phase and dispersing medium:

Dispersed	Medium	Type	Example
Solid	Liquid	Sol	Paint, ink, blood
Liquid	Liquid	Emulsion	Milk, mayonnaise, cream
Gas	Liquid	Foam	Whipped cream, sea foam
Solid	Gas	Aerosol	Smoke, dust
Liquid	Gas	Aerosol	Fog, mist
Liquid	Solid	Gel	Jelly, gelatin
Solid	Solid	Solid sol	Colored glass, alloys

Colloidal Properties

Tyndall Effect

• Definition: Colloidal particles scatter light, making the beam visible
• Observation: A light beam through a colloidal solution is visible (like a movie projector beam through fog)
• Contrast: Solutions are transparent (light passes straight through); suspensions are opaque

Real example: Fog (water droplets in air) shows Tyndall effect; you see the sun's rays

Brownian Motion

• Definition: Random, zigzag movement of colloidal particles
• Cause: Bombardment by molecules of the dispersing medium
• Observation: Under a microscope, colloid particles appear to move randomly
• Significance: Proves molecular motion; used to estimate particle size

Colloidal Stability

• Colloidal particles remain suspended indefinitely (unlike suspensions which settle)
• Reason: Repulsive electrostatic forces keep particles apart
• Most colloidal particles are charged (positively or negatively), so they repel each other

Coagulation and Peptization

Coagulation

• Definition: Colloidal particles aggregate and settle out (colloid destabilizes)
• Causes:
  • Adding salt (high ionic strength neutralizes charges)
  • Adding alcohol (reduces dielectric constant, weakens electrostatic repulsion)
  • Heating (increased molecular motion)
  • Adding opposite-charged ions (attract colloidal particles)

Example: Adding alum (Al₂(SO₄)₃) to water in water purification; coagulates colloidal particles; they settle as sludge

Peptization

• Definition: Converting a coagulated colloid back into a colloidal solution
• Method: Adding a peptizing agent (solvent that restores charge to particles)

Example: Adding water to coagulated clay; water washes away excess ions; clay particles re-disperse

[Memory Hook] Colloid = particles 1-1000 nm; Tyndall effect (scatters light); Brownian motion; stable due to electrostatic repulsion

Industrial Applications of Colloids

Water purification: Alum coagulates suspended impurities

Food industry:

• Milk: Emulsion of fat droplets in water
• Mayonnaise: Emulsion of oil in water (stabilized by egg yolk emulsifier)
• Ice cream: Foam with fat droplets and air bubbles

Medicine:

• Blood: Colloid of proteins and cells
• Vaccine adjuvants: Colloidal gold used in medical diagnostics

Environmental:

• Smoke: Aerosol colloid; scrubbers coagulate particles
• Air pollution: Acid rain forms when aerosol pollutants coagulate

Conclusion

Solutions and colloids represent the spectrum of mixture types encountered in chemistry and daily life. Understanding concentration methods, colligative properties, and colloidal behavior is essential for competitive exams, industrial applications, and scientific research.

---

23 MCQ Questions

Q1: In a solution, the substance present in the larger amount is called:

• A) Solute
• B) Solvent
• C) Solubility
• D) Saturant

Q2: A saturated solution is one that:

• A) Contains less solute than it can dissolve
• B) Contains exactly the maximum amount of solute that can dissolve
• C) Contains more solute than it can dissolve (unstable)
• D) Contains no dissolved solute

Q3: How does temperature typically affect the solubility of most solid solutes?

• A) Decreases
• B) Increases
• C) Remains constant
• D) First increases then decreases

Q4: Gas solubility in liquids:

• A) Increases with temperature
• B) Decreases with temperature
• C) Is unaffected by temperature
• D) Increases at high pressure and low temperature

Q5: "Like dissolves like" means:

• A) Polar solvents dissolve nonpolar solutes
• B) Nonpolar solvents dissolve polar solutes
• C) Polar solvents dissolve polar/ionic solutes; nonpolar solvents dissolve nonpolar solutes
• D) All substances dissolve equally

Q6: Molarity is defined as:

• A) Moles of solute per kilogram of solvent
• B) Moles of solute per liter of solution
• C) Grams of solute per liter of solvent
• D) Moles of solute per mole of solvent

Q7: If 2.93 g of NaCl (molar mass = 58.5 g/mol) is dissolved in 250 mL of water, what is the molarity?

• A) 0.1 M
• B) 0.2 M
• C) 0.5 M
• D) 1 M

Q8: Molality differs from molarity in that molality:

• A) Uses solvent mass instead of solution volume
• B) Uses solution volume
• C) Changes with temperature
• D) Is always higher than molarity

Q9: Normality is particularly useful for:

• A) Measuring gas concentrations
• B) Acid-base titrations and redox reactions (equivalents matter)
• C) Measuring solution density
• D) Determining vapor pressure

Q10: Vapor pressure lowering is caused by:

• A) Increased temperature
• B) Solute particles at the surface reducing solvent molecule escape
• C) Increased solvent concentration
• D) Reduced molecular motion

Q11: Boiling point elevation occurs because:

• A) Solute adds heat
• B) Dissolved particles increase vapor pressure
• C) Lower vapor pressure requires higher temperature to reach atmospheric pressure
• D) Solute increases intermolecular forces between water molecules

Q12: The formula ΔT_b = K_b × m is used to calculate:

• A) Freezing point depression
• B) Vapor pressure lowering
• C) Boiling point elevation
• D) Osmotic pressure

Q13: Antifreeze in car radiators works by:

• A) Adding heat to the coolant
• B) Lowering the freezing point of water
• C) Increasing boiling point (and freezing point)
• D) Preventing all phase changes

Q14: Osmosis is the movement of:

• A) Solute particles across a membrane
• B) Water molecules across a semipermeable membrane toward higher solute concentration
• C) Both solute and water
• D) Ions only

Q15: An isotonic solution has:

• A) Higher solute concentration than the cell
• B) Lower solute concentration than the cell
• C) The same solute concentration as the cell
• D) No solute at all

Q16: Colloids have particle sizes of approximately:

• A) < 1 nm
• B) 1–1000 nm
• C) > 1000 nm
• D) Variable, depending on solute

Q17: The Tyndall effect demonstrates that:

• A) Solutions are transparent
• B) Colloidal particles scatter light (beam becomes visible)
• C) Suspensions don't scatter light
• D) All mixtures scatter light

Q18: Brownian motion is observed in:

• A) Solutions only
• B) Suspensions only
• C) Colloidal solutions (particles move randomly due to bombardment by medium molecules)
• D) All mixtures

Q19: Coagulation of a colloid occurs when:

• A) Solvent is added
• B) Temperature is decreased
• C) Salt is added or opposite-charged ions are introduced
• D) Colloid becomes more dilute

Q20: Which of the following is a colloid?

• A) Saltwater solution
• B) Sand in water (suspension)
• C) Milk (emulsion)
• D) Both B and C

Q21: Blood is an example of:

• A) Suspension
• B) Solution
• C) Colloid (contains proteins, blood cells, etc.)
• D) Pure substance

Q22: Alum is used in water purification to:

• A) Disinfect water
• B) Coagulate colloidal impurities so they settle
• C) Increase pH
• D) Remove dissolved salts

Q23: Parts per million (ppm) is used to measure:

• A) Concentration of very dilute solutions (pollutants, trace elements)
• B) Molarity of concentrated solutions
• C) Colloidal suspension density
• D) Osmotic pressure

---

Answer Key: 1-B, 2-B, 3-B, 4-B, 5-C, 6-B, 7-B, 8-A, 9-B, 10-B, 11-C, 12-C, 13-B, 14-B, 15-C, 16-B, 17-B, 18-C, 19-C, 20-C, 21-C, 22-B, 23-A