Welcome to States of Matter: Unlocking the Secrets of Solids, Liquids, and Gases!

Hi there! This chapter, Topic 4 in Physical Chemistry, is all about understanding the differences between the three main states of matter: gases, liquids, and solids. Why does ice melt? Why does steam create pressure? The answers lie in the structure and the forces acting between particles.

Don't worry if this seems tricky at first! We will break down the complex structures of solids and the laws governing gases into simple, manageable pieces. Mastering this topic helps you understand everything from weather patterns to materials science!


4.1 The Gaseous State: Ideal and Real Gases

The Origin of Gas Pressure

When you inflate a balloon, where does the pressure come from? It comes from motion!

  • Gas molecules are constantly moving rapidly and randomly.
  • Pressure is simply the force exerted when these particles collide with the walls of the container.
  • More frequent or harder collisions mean higher pressure.

The Ideal Gas Model (The Perfect Gas)

In Chemistry, we often use models to simplify complex reality. The Ideal Gas Model is a set of two assumptions we make about gas particles. While no gas is truly "ideal," many gases behave close to this model at high temperatures and low pressures.

Key Assumptions of an Ideal Gas:
  1. Zero Particle Volume: The volume occupied by the gas molecules themselves is negligible (so small it can be ignored) compared to the total volume of the container.
  2. No Intermolecular Forces of Attraction: There are no attractive or repulsive forces between the particles. They move completely independently until they collide. (Think of them as perfectly non-sticky billiard balls).

Did You Know? Real gases deviate from ideal behaviour because their particles actually do have finite volume and small attractive forces (van der Waals forces). This deviation becomes significant at low temperatures (where forces matter more) and high pressures (where particle volume matters more).


The Ideal Gas Equation: \(pV = nRT\)

This is a fundamental equation used to link the four variables that define a gas sample: Pressure, Volume, Moles, and Temperature.

Understanding the Variables and Units

To use this equation correctly, you must use the appropriate SI units. This is a common exam mistake!

  • \(p\): Pressure (measured in Pascals, Pa). Note: You may be given kPa, so convert! \(1 \text{ kPa} = 1000 \text{ Pa}\)
  • \(V\): Volume (measured in cubic metres, \(m^3\)). Note: You may be given \(dm^3\) or \(cm^3\).
    • To convert from \(dm^3\) to \(m^3\): Divide by 1000.
    • To convert from \(cm^3\) to \(m^3\): Divide by 1,000,000.
  • \(n\): Amount of substance (measured in moles, mol).
  • \(R\): The Molar Gas Constant (constant value, usually given as \(8.31 \text{ J} \text{K}^{-1} \text{mol}^{-1}\)).
  • \(T\): Temperature (measured in Kelvin, K).
    • To convert from \(^{\circ}\text{C}\) to \(K\): Add 273. (The syllabus assumes standard conditions use 298 K).
Step-by-Step Calculation Guide: Finding \(M_r\)

The ideal gas equation is frequently used in AS Chemistry to determine the relative molecular mass, \(M_r\), of an unknown volatile liquid or gas. Here is the trick:

  1. Recall the definition of moles: \(n = \frac{\text{mass}}{\text{Mr}}\).
  2. Substitute this into the ideal gas equation: \(pV = \left(\frac{m}{M_r}\right)RT\).
  3. Rearrange to solve for \(M_r\): \[M_r = \frac{mRT}{pV}\]

Quick Review Tip: Unit Conversions
Always, always, always check your units before putting numbers into \(pV=nRT\). Convert to Pa, \(m^3\), and K first!


4.2 Bonding and Structure of Solids

Why are diamonds hard and dry ice easy to crush? It all comes down to how the particles are held together. Solids exist as crystalline lattices, and we categorize them into four main structures based on their bonding.

1. Giant Ionic Structures

These solids consist of a huge lattice of oppositely charged ions (cations and anions) held together by strong electrostatic forces (ionic bonds).

  • Structure: Regular arrangement of positive and negative ions (e.g., cubic lattice for sodium chloride, NaCl).
  • Bonding: Very strong ionic bonds extending throughout the entire structure.
  • Examples: Sodium Chloride (NaCl), Magnesium Oxide (MgO).
Properties Explained:
  • Melting/Boiling Points: Extremely high. A lot of energy is required to break the strong ionic bonds holding the lattice together.
  • Electrical Conductivity:
    • Solid: Poor (insulator). Ions are fixed in position and cannot move.
    • Molten or Aqueous: Good conductor. The ions are free to move and carry charge.
    • Solubility: Generally soluble in polar solvents like water, as water molecules can surround and stabilise the individual ions.

    2. Simple Molecular Structures

    These solids consist of individual molecules held together by weak forces between them. The strong forces (covalent bonds) exist *only* within the molecule.

    • Structure: Lattice of small, individual molecules.
    • Bonding: Strong covalent bonds hold atoms within the molecule; weak intermolecular forces (van der Waals or Hydrogen bonds) hold molecules to each other.
    • Examples: Iodine (\(I_2\)), Buckminsterfullerene (\(C_{60}\)), and Ice (\(H_2O\)).
    Properties Explained:
    • Melting/Boiling Points: Low. When melting or boiling, you only need to overcome the weak intermolecular forces, which requires little energy.
    • Electrical Conductivity: Poor (insulator). They have no mobile charged particles (no ions or delocalised electrons).
    • Solubility:
      • Soluble in non-polar solvents (e.g., \(I_2\) in hexane).
      • Ice is a special case: its lattice is held together by relatively strong hydrogen bonds, explaining why its melting point is higher than other small molecules of similar \(M_r\).

      Analogy: Think of Simple Molecular structures like LEGO bricks held together by weak Velcro (intermolecular forces). It’s easy to pull the bricks apart (melt it), even though the bricks themselves (the molecules) are strong.


      3. Giant Molecular (Giant Covalent) Structures

      In these structures, every atom is linked to its neighbours by strong covalent bonds throughout the entire crystal, forming one massive molecule.

      • Structure: Infinite lattice where all atoms are covalently bonded.
      • Bonding: Extremely strong covalent bonds throughout.
      • Examples: Diamond, Silicon(IV) Oxide (\(SiO_2\)), Graphite.
      Properties Explained:
      • Melting/Boiling Points: Extremely high. To melt or boil, you must break countless strong covalent bonds.
      • Solubility: Insoluble in almost all solvents, as the bonds are too strong to be broken by solvent molecules.
      • Electrical Conductivity (Key Differences):
        • Diamond & \(SiO_2\) (Sand): Poor conductor (insulator). All valence electrons are fixed in covalent bonds; none are free to move.
        • Graphite: Good conductor (excellent along the layers). Graphite has a layer structure; each carbon atom is bonded to only three others. The remaining valence electron is delocalised and free to move within the layers, carrying charge.

      Memory Aid: Graphite is a conductor because it's only sp2 hybridised, leaving one electron free. Diamond is sp3, binding all electrons tightly.


      4. Giant Metallic Structures

      These solids are found in metals. They consist of a lattice of positive metal ions surrounded by a mobile "sea" of delocalised valence electrons.

      • Structure: Lattice of positive metal ions.
      • Bonding: Metallic bonding – the electrostatic attraction between the positive ions and the sea of delocalised electrons.
      • Examples: Copper (Cu), Iron, Magnesium.
      Properties Explained:
      • Melting/Boiling Points: Variable, usually high (depending on ion charge and size, which affects the strength of the attraction to the electron sea).
      • Electrical Conductivity: Excellent conductor (both solid and molten). The delocalised electrons are free to move when a potential difference is applied, carrying charge.
      • Malleability/Ductility: Excellent (can be hammered into sheets or drawn into wires). When stress is applied, the layers of positive ions can slide over one another without disrupting the metallic bond, as the electron sea holds the structure together regardless of the arrangement.

      Summary: Structure, Bonding, and Properties Comparison

      A crucial skill in this topic is being able to deduce the type of structure and bonding from given physical properties (M.P., conductivity, etc.).

      Quick Review Box: The Four Structures
      Structure Type Example Bonding/Force Overcome Melting Point Conductivity (Solid)
      Giant Ionic NaCl, MgO Ionic Bonds (Strong) Very High Poor
      Simple Molecular I₂, C₆₀, Ice Intermolecular Forces (Weak) Low Poor
      Giant Molecular Diamond, SiO₂ Covalent Bonds (Very Strong) Extremely High Poor (Graphite is an exception)
      Giant Metallic Copper, Fe Metallic Bonds (Strong) High/Variable Excellent

      Common Mistakes to Avoid

      • Confusing Strong Bonds with Strong Forces: When melting a simple molecular solid (like iodine), you are not breaking the strong covalent bonds inside the \(I_2\) molecules. You are only breaking the weak van der Waals forces between the molecules.
      • Forgetting Temperature Units: Always convert temperature to Kelvin (K) for the ideal gas equation (\(pV=nRT\)).
      • Graphite vs. Diamond: Remember that graphite's conductivity comes from delocalised electrons residing between its layers, a unique feature among giant covalent structures.

      Key Takeaway: The physical properties (like melting point and conductivity) of any substance are a direct result of the type and strength of the chemical bonds or intermolecular forces holding its particles in their specific lattice structure. Strong forces (ionic, metallic, covalent network) lead to high melting points; weak forces (intermolecular) lead to low melting points.