Solutions: Chemistry unit 1

Unit 1: Solutions

• Almost all bodily processes occur in liquid solutions.

Objectives

• After studying this unit, students should be able to perform these tasks:
• Describe the formation of different types of solutions.
• Express the concentration of a solution in different units.
• State and explain Henry’s and Raoult’s laws.
• Distinguish between ideal and non-ideal solutions and explain deviations from Raoult's law.
• Describe colligative properties and correlate these with molar masses of solutes.
• Explain abnormal colligative properties of some solutes.

Introduction to Solutions

• Mixtures containing two or more pure substances are more common than pure substances.
• The utility or importance of a mixture depends on composition.
• For instance, brass (copper and zinc) possesses distinct properties from German silver (copper, zinc, and nickel) or bronze (copper and tin).
• Fluoride ions in water can have variable effects; 1 ppm prevents tooth decay whereas 1.5 ppm causes teeth to become mottled; higher concentrations are poisonous.
• Intravenous injections must be dissolved in water with salt concentrations matching blood plasma.

Liquid Solutions

• Liquid solutions and their formation are the primary focus.
• Properties, like vapour pressure and colligative properties, are examined.
• Alternatives in expressing solute concentrations in liquid solutions are explored.

Types of Solutions

• Solutions are homogeneous mixtures containing two or more components with uniform composition and properties.
• The solvent is the component present in the greatest quantity and determines the solution's physical state.
• Solutes are one or more components other than the solvent in the solution.
• Focus is given to binary solutions, which consist of two components.
• Each component (solute and solvent) can exist as a solid, liquid, or gas.

Types of Solutions Table

• Gaseous Solutions:
  • Gas solute in gas solvent yields a mixture of oxygen and nitrogen gases.
  • Liquid solute in gas solvent yields chloroform mixed with nitrogen gas.
  • Solid solute in gas solvent yields camphor in nitrogen gas.
• Liquid Solutions:
  • Gas solute in liquid solvent yields oxygen dissolved in water.
  • Liquid solute in liquid solvent yields ethanol dissolved in water.
  • Solid solute in liquid solvent yields glucose dissolved in water.
• Solid Solutions:
  • Gas solute in solid solvent yields hydrogen in palladium
  • Liquid solute in solid solvent yields amalgam of mercury with sodium.
  • Solid solute in solid solvent yields copper dissolved in gold.

Expressing Solution Concentration

• Solution composition can be described qualitatively (dilute or concentrated) or quantitatively.
• Quantitative descriptions are more useful in real-world applications.
• Mass Percentage (w/w):
  • Mass % of a component = (Mass of the component in the solution / Total mass of the solution) * 100
  • A 10% glucose solution means 10g of glucose is in 90g of water, yielding 100g solution.
  • Mass percentage is used for commercial bleaching solution containing 3.62% sodium hypochlorite.
• Volume Percentage (V/V):
  • Volume % of a component = (Volume of the component / Total volume of the solution) * 100
  • 10% ethanol solution by volume signifies 10 mL ethanol dissolved in water, totaling 100 mL volume.
  • 35% (v/v) ethylene glycol solution acts as antifreeze.
• Mass by Volume Percentage (w/V):
  • Mass by volume percentage indicates the mass of solute dissolved in 100 mL of solution; commonly used in medicine and pharmacy.
• Parts per Million (ppm):
  • Parts per million = (Number of parts of the component / Total number of parts of all components) * 10^6
  • Concentration in parts per million can be mass to mass, volume to volume, or mass to volume.
  • Seawater (1030 g/L) contains about 6 x 10^-3 g dissolved oxygen, or 5.8 ppm.
  • Pollutant concentration in water or atmosphere is commonly expressed as µg/mL^-1 or ppm.
• Mole Fraction:
  • Denoted by x, the mole fraction is calculated as moles of the component divided by total moles.
  • Mole fraction of component = Number of moles of the component / Total number of moles of all the components
  • For binary mixtures, xA = nA / (nA + nB)
  • For a solution with i components: xi = ni / Σni
  • Sum of all mole fractions in a solution equals 1.
  • Mole fraction is valuable for relating physical properties to concentration. e.g. vapour pressure
• Molarity:
  • Molarity (M) is defined as the number of moles of solute dissolved in one litre (or one cubic decimetre) of solution.
  • Molarity = Moles of solute / Volume of solution in litres
  • 0.25 M NaOH indicates 0.25 mol NaOH per litre (or cubic decimetre).
• Molality:
  • Molality (m) expresses moles of solute per kilogram of solvent.
    • Molality (m) = Moles of solute / Mass of solvent in kg
    • 1.00 m KCl contains 1 mol (74.5 g) KCl per kg of water.
  • Mass %, ppm, mole fraction, and molality are temperature independent.
  • Molarity depends on temperature because volume varies with temperature.

Solubility

• Solubility is the maximum solute amount that can dissolve in a specific solvent amount at a given temperature.
• Solubility depends on solute and solvent nature and on temperature and pressure.

Solubility of a Solid in a Liquid

• Not every solid dissolves in a given liquid; sodium chloride and sugar dissolve readily in water while naphthalene and anthracene do not.
• Polar solutes dissolve in polar solvents, and non-polar solutes dissolve in non-polar solvents, thus like dissolves like.
• Dissolution is the process where some solute dissolves, increasing concentration, while crystallization involves solute particles separating out.
• Equilibrium is reached when dissolution and crystallization occur at the same rate.
• A saturated solution contains the maximum solute amount, while an unsaturated solution can dissolve more solute.
• Solubility is the solute concentration in a saturated solution, affected by substance nature, temperature, and pressure.
• Temperature Effects:
  • The solubility of a solid in a liquid is affected by temperature changes.
  • Dissolution processes follow Le Chatelier's Principle within a dynamic equilibrium.
  • Endothermic dissolution (∆sol H > 0) implies solubility rises with temperature, and exothermic (∆sol H < 0) implies solubility decreases.
• Pressure's negligible impact:
  • Pressure minimally affects solid solubility in liquids since both are nearly incompressible.

Solubility of a Gas in a Liquid

• Many gases can dissolve in water.
• Oxygen's limited water solubility sustains aquatic life.
• Hydrogen chloride (HCl) is highly soluble in water.
• Gases in liquids are greatly affected by pressure and temperature.
• Gas solubility increases with higher pressure.
• Henry's Law establishes a quantitative relationship between gas pressure and the gas solubility in a solvent.
  • The law states the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid surface at constant temperature.
    • p = KHx
      • p = partial pressure of the gas in the vapor phase
      • x = mole fraction of the gas in the solution
      • KH = Henry’s Law constant
• Gases have different Henry’s Law constants (KH) at a given temperature, a function of the gas's nature
• The lower is the solubility of the gas in the liquid, the higher the value of KH at a given pressure.
• KH values for nitrogen and oxygen increase with rising temperature, which indicates that gas solubility decreases.
• Aquatic species thrive more in colder waters.
• Applications of Henry's law:
  • High pressure is applied to seal bottles, increasing carbon dioxide (CO2) solubility in soda water and soft drinks.
  • Scuba divers experience increased gas solubility in blood at high underwater pressures, which releases upon ascent. Then nitrogen bubbles in blood creates "bends".
  • Tanks for divers have air diluted with helium to avoid bends and nitrogen toxicity,
• Elevated altitudes cause lower oxygen partial pressure, leading to lower oxygen levels in blood and tissues, causing weakness and anoxia.
• Gas solubility in liquids decreases as temperature rises.
• Dissolution mirrors condensation, releasing heat; solubility diminishes with temperature rise, as per Le Chatelier’s principle under dynamic equilibrium.