Overview
This lecture covers the states of matter (solids, liquids, gases) for CIE Chemistry, focusing on their properties, transitions, key gas laws, and important structural examples, including allotropes and hydrogen bonding. The content is tailored for the Cambridge International syllabus and emphasizes the use of scientific terminology and understanding at an advanced level.
States of Matter: Basics
- Solids: Particles are arranged in a regular, tightly packed structure. Solids are incompressible and maintain a fixed shape due to strong forces holding the particles in place.
- Liquids: Particles are closely packed but arranged randomly, allowing them to flow and take the shape of their container. Liquids are also incompressible and can diffuse, though more slowly than gases.
- Gases: Particles are randomly arranged and widely spaced, resulting in low density. Gases are compressible, fill the entire container, and can move freely in all directions.
Properties of Gases
- Gases have low density because their particles are far apart.
- They are compressible due to the large spaces between particles.
- Gases have weak or negligible intermolecular forces, allowing particles to move independently.
- Diffusion: Gases diffuse rapidly from areas of high concentration to low concentration because of their random and rapid particle motion.
- Diffusion is a key property of gases, but it can also occur in liquids, though at a slower rate due to closer particle packing.
Ideal and Non-Ideal Gases
- Ideal Gases: Assumed to have particles with negligible volume, no intermolecular forces, random continuous motion, and perfectly elastic collisions (no energy loss).
- Boyle’s Law: For a fixed mass of gas at constant temperature, pressure × volume = constant (PV = constant).
- Charles’s Law: For a fixed mass of gas at constant pressure, volume/temperature = constant (V/T = constant).
- Ideal Gas Equation: PV = nRT, where:
- P = pressure (Pa)
- V = volume (m³)
- n = number of moles
- R = gas constant (8.31 J·mol⁻¹·K⁻¹)
- T = temperature (K)
- Units must be correct: pressure in pascals, volume in cubic meters, temperature in kelvin.
- Calculations: Rearranging the ideal gas equation is essential. For example, to find volume: V = nRT/P. Remember to convert units as needed (e.g., kPa to Pa, cm³ to m³).
- Real Gases: Deviate from ideal behavior at high pressures (particles are closer, volume matters) and low temperatures (intermolecular forces become significant). Real gases may condense or behave more like liquids under these conditions.
Liquids and Changes of State
- Liquids have a random arrangement but are tightly packed, making them incompressible and able to diffuse, though more slowly than gases.
- Melting: Solid to liquid; energy is absorbed to break bonds between particles.
- Freezing: Liquid to solid; energy is released as bonds form.
- Vaporization: Liquid to gas; energy is absorbed to separate particles, usually at the surface (evaporation) or throughout the liquid (boiling).
- Condensation: Gas to liquid; energy is released as particles come together.
- Evaporation: Occurs at the surface and below boiling point, does not require the entire liquid to reach boiling temperature.
- Boiling: Involves the entire liquid, with bubbles forming throughout as vaporization occurs at a specific temperature.
Heating and Cooling Curves
- Heating/Cooling Graphs: Show how temperature changes as a substance is heated or cooled.
- Flat (horizontal) sections: Indicate energy is being used to break or form bonds during state changes (melting or boiling), not to increase temperature.
- Melting Point: Temperature at which a solid becomes a liquid.
- Boiling Point: Temperature at which a liquid becomes a gas.
- During cooling, energy is released as bonds form, and the process reverses.
Solids & Bonding Types
- Solids have a regular, tightly packed arrangement of particles.
- Types of Bonding in Solids:
- Giant covalent (e.g., diamond, graphite)
- Simple molecular (e.g., iodine)
- Giant ionic (e.g., sodium chloride)
- Metallic (e.g., copper, magnesium)
- Hydrogen bonding (strong intermolecular force, not a true bond)
- Physical Properties to Know:
- State at room temperature and pressure
- Electrical conductivity (as solid and liquid)
- Solubility in water and other solvents
- Melting and boiling points
- Solubility: Non-polar substances dissolve in non-polar solvents; polar substances dissolve in polar solvents.
Allotropes of Carbon
- Graphite: Layers of carbon atoms with strong covalent bonds within layers and weak forces between layers. Delocalized electrons allow electrical conductivity. Soft, high melting point, low density, insoluble in water.
- Diamond: Each carbon atom bonded to four others in a tetrahedral structure. Very hard, high melting point, does not conduct electricity, insoluble in water, good heat conductor.
- Graphene: Single layer of graphite, one atom thick. Strong, lightweight, transparent, excellent electrical conductor. Used in electronics, screens, and materials requiring strength and low weight.
- Buckminsterfullerene (C₆₀, Buckyballs): Spherical molecules of 60 carbon atoms. Used in nanotechnology and medicine, especially for drug delivery due to their cage-like structure.
Hydrogen Bonding and Water
- Hydrogen Bonds: Form between hydrogen and highly electronegative atoms (N, O, F). Responsible for water’s unique properties.
- Effects of Hydrogen Bonding:
- Raises boiling and melting points (e.g., HF has a higher boiling point than HCl due to hydrogen bonding).
- Causes surface tension (water beads on surfaces, insects can walk on water).
- Increases viscosity in substances like honey and treacle.
- Explains why ice is less dense than liquid water (hydrogen bonds create an open structure in ice, making it float).
- Hydrogen Halides: Only HF can hydrogen bond, resulting in a much higher boiling point compared to HCl, HBr, and HI, which only have dipole-dipole or van der Waals forces.
Dissolving and Solubility
- Ionic Compounds in Water: Dissolve via hydration; water molecules surround and separate ions due to polarity and hydrogen bonding.
- For dissolution, the new interactions (hydration) must be as strong as or stronger than the original ionic bonds.
- Not all salts are soluble; sometimes heating is needed to dissolve them.
- Solubility Rules: Non-polar substances dissolve in non-polar solvents (e.g., oil in hexane), polar substances dissolve in polar solvents (e.g., ammonia in water).
Metallic Structure
- Metals consist of a lattice of positive ions surrounded by a sea of delocalized electrons.
- Properties:
- Good electrical and thermal conductivity due to free-moving electrons.
- High melting and boiling points from strong electrostatic attraction.
- Malleable and ductile; layers of atoms can slide over each other while electrons maintain the bond.
- Insoluble in water; metallic bonds are too strong to be broken by water.
- Strength increases with more delocalized electrons per atom (e.g., magnesium vs. sodium).
Key Terms & Definitions
- Diffusion: Movement of particles from an area of high concentration to low concentration.
- Ideal Gas: A gas that follows the assumptions of negligible particle volume, no intermolecular forces, random motion, and elastic collisions.
- PV = nRT: The ideal gas equation relating pressure, volume, moles, gas constant, and temperature.
- Allotropes: Different structural forms of the same element with distinct properties (e.g., diamond, graphite, graphene).
- Hydrogen Bond: A strong intermolecular force between hydrogen and N, O, or F atoms.
- Hydration: The process of water molecules surrounding and separating ions when dissolving ionic compounds.
Action Items / Next Steps
- Practice rearranging and using the ideal gas equation with correct SI units; ensure you can convert between cm³, dm³, and m³.
- Review the differences between ideal and real gases, and understand when deviations occur.
- Revise topic 3 for detailed bonding types, intermolecular forces, and solubility rules.
- Attempt exam questions on states of matter, changes of state, and calculations involving the ideal gas law.
- Be able to describe and compare the structures and properties of different allotropes of carbon and explain the effects of hydrogen bonding in water and other substances.