Chemistry: Solutions, Solubility, and Concentration — Study Notes

Observations

  • i. Both groups A and B obtained homogeneous mixtures since the composition is uniform throughout.
  • ii. Both groups C and D obtained heterogeneous mixtures since they have physically distinct boundaries and their composition is not uniform.

Definitions and Core Concepts

  • Activity aim and materials indicate a focus on distinguishing homogeneous vs heterogeneous mixtures.
  • SOLUTION: When one substance dissolves or mixes well with another substance, the mixture is called a solution.
  • A solution may be defined as a homogeneous mixture of two or more non-reacting substances whose composition can be varied within certain limits.
  • Air is a mixture of gas in gas (an example of a mixture that is not necessarily a solution).
  • Alloys are homogeneous mixtures of metals with other metals or non-metals and cannot be separated into their components by physical methods, yet they are considered mixtures because they show the properties of their constituents and can have variable composition.
  • Examples of alloys: Brass (Copper, Zinc); Bronze (Copper, Tin); German Silver (Copper, Zinc, Nickel); Steel (Iron, Carbon, Manganese); Stainless Steel (Iron, Carbon, Manganese, Chromium, Nickel).
  • Bronze is noted as one of the first alloys created by humans.

Alloys: When and Why

  • Alloys are mixtures that are typically homogeneous in composition but can have variable compositions.
  • They often exhibit properties different from their constituent elements (e.g., hardness, strength) and are engineered for specific uses.

Components of a Solution

  • A solution has two components:
    • Solvent: The component that dissolves the other component; usually the larger component.
    • Solute: The component that dissolves in the solvent; usually the smaller component.
  • A solution is a binary solution when it has two components (solvent + solute).
  • Example: Tincture of iodine is a solution where iodine is the solute in alcohol (solvent).
  • Carbonated drinks: Carbon dioxide (CO₂) is the solute in water (the solvent).

Properties and Characteristics of Solutions

  • A solution is a homogeneous mixture with extremely small solute particles (< 1 nm in diameter).
  • The particles of a solution cannot be seen under a microscope.
  • The particles pass through filter paper; thus, a solution cannot be separated by filtration.
  • Solutions are very stable.
  • A true solution does not scatter light (no Tyndall effect) because the particles are too small.
  • The components of a solution do not chemically react with one another.
  • A solution is always transparent in nature.
  • From a true solution, the solute can be recovered by evaporation or crystallisation.

Types of Solutions

  • Based on the Nature of the Solvent:
    • Aqueous Solution: Water acts as the solvent (water is the universal solvent).
    • Alcohol Solution: Alcohol acts as the solvent.
    • Non-Aqueous Solution: Solvent other than water (example given: benzene with chloroform).
  • Based on the Amount of Solute Soluble in a Solvent: Saturated, Unsaturated, Supersaturated.

Solutions: Based on Solubility and Conditions

  • Saturated Solution: No more solute can be dissolved in the solvent at a given temperature.
  • Unsaturated Solution: More solute can be dissolved in the solvent at that temperature.
  • Supersaturated Solution: Contains more solute than required to form a saturated solution; can be unstable and crystals may form.
  • Solubility: The maximum amount of solute in grams that can be dissolved in 100 g of the solvent at a given temperature to form a saturated solution.
  • Example solubilities at 20°C (in g of solute per 100 g of water):
    • Copper sulfate: 21\;g
    • Potassium nitrate: 32\;g
    • Potassium chloride: 34\;g
    • Sodium chloride: 36\;g
    • Ammonium chloride: 37\;g
    • Sugar: 204\;g

Solubility Curves and Gas Solubility

  • Solubility curves show how much solute dissolves in a solvent at different temperatures for various salts (examples include NaNO₃, CaCl₂, Pb(NO₃)₂, NaCl, KNO₃, KCl, etc.).
  • Effect of Pressure on Solubility:
    • Solubility of solids in liquids is largely unaffected by pressure changes (solids and liquids are incompressible).
  • Solubility of Gases in Liquids:
    • Governed by Henry's Law: the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the solution at a given temperature.
    • Formula: S = kH P{gas}
    • Applications:
    • Soft drink bottles are sealed at high pressure to increase CO₂ solubility.
    • At high altitude, lower partial pressure of O₂ reduces its dissolved concentration in blood, risking hypoxia.
    • Scuba diving: gas mixtures in tanks reduce nitrogen dissolution to prevent bends; tanks may use He, N₂, and O₂ to optimize solubility.
  • Temperature effect on gas solubility: Solubility of gases in liquids generally decreases with increasing temperature.

Concentration of a Solution

  • The concentration of a solution is the amount of solute present in each quantity of the solution; the relative proportion of solute to solvent can vary.
  • Terms:
    • Dilute vs. Concentrated (qualitative indications of relative amounts).

Concentration Terms (Quantitative)

  • Common concentration metrics include Mole Fraction, Molality, Molarity, and Percent Concentrations (% w/w, % w/V, % V/V, % w/V, Strength, ppm/ppb).
  • Qualitative composition terms:
    • Dilute: relatively very small quantity of solute compared to solvent.
    • Concentrated: relatively very large quantity of solute compared to solvent.
  • Quantitative concentration terms (definitions):
    • % w/w (Mass percentage):
    • Definition: \%\,\text{w/w} = \frac{m{\text{solute}}}{m{\text{solution}}} \times 100
    • Note: Mass percent is generally temperature independent.
    • % w/V (Mass by volume percentage):
    • Definition: \%\,\text{w/V} = \frac{m{\text{solute}}}{V{\text{solution}}} \times 100
    • Commonly used in pharmacy; volume is in mL, mass in g; temperature can affect volume measurements.
    • % V/V (Volume percentage):
    • Definition: \%\,\text{V/V} = \frac{V{\text{solute}}}{V{\text{solution}}} \times 100
    • Used for liquid-liquid solutions; volume is involved in the calculation.

Worked Examples and Practice Questions

  • Example 1: A solution contains sugar dissolved in 150 g of water. If the mass of the solution is 200 g, calculate the concentration in terms of the mass-by-mass percentage (% w/w).
    • Solution:
    • Mass of solute (sugar) = Mass of solution - Mass of solvent = 200 g - 150 g = 50 g
    • % w/w = (50 / 200) × 100 = 25%
  • Example 2: How many grams of NaCl must be dissolved in 54.0 g of H₂O to give a 10% weight (mass) solution?
    • Let ms be mass of NaCl. Then ms / (m_s + 54) = 0.10
    • ms = 0.10(ms + 54) → ms = 0.10 ms + 5.4 → 0.90 ms = 5.4 → ms = 6.0 g
  • Example 3: Find weights in 15% (w/v) solution with density 1.06 g/mL.
    • 15% (w/v) means 15 g solute per 100 mL of solution.
    • Mass of 100 mL of solution = density × volume = 1.06 g/mL × 100 mL = 106 g
    • Mass of solute = 15 g
    • Mass of solvent = total mass − solute = 106 g − 15 g = 91 g
  • Example 4: % w/v of NaOH solution is 10%, if the volume of solution is 2 dm³. Find the mass of solute.
    • 2 dm³ = 2000 mL
    • 10% w/v means 10 g solute per 100 mL solution
    • Mass of solute = (2000 mL / 100 mL) × 10 g = 20 × 10 g = 200 g

Summary of Key Points to Remember

  • Solutions are homogeneous mixtures with very small solute particles that cannot be separated by filtration and are typically transparent.
  • The solvent is usually the component present in the larger amount; the solute is the component dissolved in the solvent.
  • Solubility data at a given temperature helps determine whether a solution is saturated, unsaturated, or supersaturated.
  • Gas solubility in liquids is governed by Henry’s Law; gas solubility increases with pressure and generally decreases with temperature.
  • Concentration terms provide quantitative measures of how much solute is present in a solution and include mass percentages (% w/w), volume percentages (% V/V), and mass-by-volume percentages (% w/V).
  • Practical calculations often require using density to relate mass and volume, especially for % w/v calculations.

Important Formulas (for quick reference)

  • Solubility concept: maximum grams of solute per 100 g of solvent at a given temperature to form a saturated solution. For solids in water at 20°C, example solubilities include:
    • \text{CuSO}_4 \; \to\; 21\,\text{g}
    • \text{KNO}_3 \; \to\; 32\,\text{g}
    • \text{KCl} \; \to\; 34\,\text{g}
    • \text{NaCl} \; \to\; 36\,\text{g}
    • \text{NH}_4\text{Cl} \; \to\; 37\,\text{g}
    • Sugar \; \to\; 204\,\text{g}
  • Henry’s Law: S = kH P{gas}
  • Percentage concentration definitions:
    • \%\,\text{w/w} = \frac{m{\text{solute}}}{m{\text{solution}}} \times 100
    • \%\,\text{w/V} = \frac{m{\text{solute}}}{V{\text{solution}}} \times 100
    • \%\,\text{V/V} = \frac{V{\text{solute}}}{V{\text{solution}}} \times 100$$

Notes on Conversions and Applications

  • Converting saturated to unsaturated is achieved by adding solvent.
  • Converting between solution types (saturated, unsaturated, supersaturated) can be influenced by cooling or heating, which affects solubility and potential crystallisation.
  • Density data is essential when calculating mass-related concentrations for liquids.