Chemistry Study Notes: Molecular Shapes, Formulae & The Macro-Micro Link

Hello! Welcome to these study notes designed to help you master one of the most fundamental ideas in Chemistry: how the tiny, invisible world of atoms and molecules (the microscopic world) determines the properties of the substances we see and touch every day (the macroscopic world). We'll explore how atoms bond, how to write their formulae, what shapes they form, and how this all connects to create the world around us. Don't worry if it sounds complicated – we'll break it down step-by-step!


Part 1: The World of Chemical Bonds - Why Atoms Stick Together

Atoms are a bit like people – they want to be stable and happy! For most atoms, 'happy' means having a full outermost electron shell, just like the very stable noble gases (like Neon or Argon). To achieve this, they form chemical bonds with other atoms. There are three main ways they do this.

Ionic Bonding: The Giver and the Taker

This happens between a metal and a non-metal. The metal atom wants to give away its outer shell electrons, and the non-metal atom wants to take them!

  • The metal atom loses electrons and becomes a positively charged ion called a cation. (Think: Cats have paws... PAWS-itive charge!)
  • The non-metal atom gains electrons and becomes a negatively charged ion called an anion.

The strong electrostatic attraction between these opposite charges is called an ionic bond. It's like a super-strong magnetic pull holding the ions together.

Example: Sodium Chloride (NaCl)

Sodium (Na) has 1 outer electron. It really wants to give it away. Chlorine (Cl) has 7 outer electrons. It just needs one more to be full. So, Na gives its electron to Cl.

Na (2,8,1) → Na+ (2,8) + e-

Cl (2,8,7) + e- → Cl- (2,8,8)

The result is a positive sodium ion (Na+) and a negative chloride ion (Cl-), stuck together tightly!

Quick Review: Ionic Compounds

- Formed between: Metals and Non-metals
- Electrons are: Transferred
- Particles are: Positive and Negative Ions
- Force: Strong electrostatic attraction

Covalent Bonding: The Great Electron Sharers

This happens between two non-metal atoms. Since both atoms want to gain electrons, they can't just transfer them. Instead, they compromise by sharing electrons. Each shared pair of electrons forms one covalent bond.

  • Single Bond: One pair of electrons is shared (e.g., in H₂ or Cl₂).
  • Double Bond: Two pairs of electrons are shared (e.g., in O₂ or CO₂).
  • Triple Bond: Three pairs of electrons are shared (e.g., in N₂).

These shared electrons hold the two atoms together, forming a stable unit called a molecule.

Did you know? Special Covalent Bonds!

Sometimes, one atom provides BOTH electrons for the shared pair. This is called a dative covalent bond. You can see this in ions like the ammonium ion (NH₄⁺) and the hydronium ion (H₃O⁺).

Quick Review: Covalent Compounds

- Formed between: Non-metals only
- Electrons are: Shared
- Particles are: Molecules
- Force: Strong attraction between nucleus and shared electrons

Metallic Bonding: A Sea of Electrons

This is what happens inside a piece of metal. Imagine a grid of positive metal ions (cations) sitting in a "sea" of electrons. These electrons, called delocalised electrons, don't belong to any single atom; they are free to move throughout the entire metal structure. The attraction between the positive ions and the negative electron "sea" is the metallic bond.

Analogy Time!

Think of it like a community swimming pool. The people are the positive metal ions, fixed in their general spots, while the water is the sea of delocalised electrons, flowing freely around everyone. This freedom of movement for electrons is key to understanding the properties of metals!


Part 2: From Bonds to Giant Structures - The Macro View

Now that we know how atoms bond (micro), let's see what kind of large structures they build (macro) and how that affects their properties.

Giant Ionic Structures

Ionic compounds don't just form one-to-one pairs. They build a huge, repeating 3D lattice of positive and negative ions. Think of a giant, perfectly stacked pile of oranges (Na⁺) and apples (Cl⁻). This is a giant ionic lattice.

  • Properties & Why:
    • High Melting & Boiling Points: It takes a massive amount of energy to break the strong electrostatic forces between all the ions.
    • Do NOT conduct electricity when solid: The ions are locked in place and cannot move.
    • DO conduct electricity when molten or dissolved in water: The ions are free to move and carry charge.
    • Brittle: If you hit the lattice, layers of ions shift. Positive ions align with positive ions, and they repel, shattering the crystal.
  • Examples: Sodium chloride (NaCl), Caesium chloride (CsCl)

Simple Molecular Structures

Substances with covalent bonds often exist as individual, separate molecules. The covalent bonds inside each molecule are very strong, but the forces between the molecules are very weak. These weak forces are called intermolecular forces (we'll learn more about these later!).

  • Properties & Why:
    • Low Melting & Boiling Points: It's very easy to overcome the weak intermolecular forces and separate the molecules. You are NOT breaking the strong covalent bonds!
    • Do NOT conduct electricity: There are no free-moving ions or electrons.
    • Often soft: The molecules can slide past each other easily.
  • Examples: Carbon dioxide (CO₂), Iodine (I₂), Water (H₂O)

Giant Covalent Structures (or Giant Molecular Structures)

Imagine a substance where every atom is joined to its neighbours by strong covalent bonds, forming one enormous molecule or network. There are no separate, individual molecules.

  • Properties & Why:
    • VERY High Melting & Boiling Points: You have to break millions of strong covalent bonds to melt or boil it, which requires a huge amount of energy.
    • Insoluble in water: The strong bonds are too hard to break.
    • Usually do NOT conduct electricity: No delocalised electrons. (The exception is graphite!).
  • Key Examples:
    • Diamond: Each carbon atom is bonded to 4 others. It's extremely hard – used in cutting tools.
    • Graphite: Each carbon atom is bonded to 3 others, forming flat layers. The layers are held by weak forces, so they can slide – making it a good lubricant. It has delocalised electrons between the layers, so it can conduct electricity!
    • Quartz (SiO₂): Similar structure to diamond, very hard and has a high melting point.

Giant Metallic Structures

This is the structure we discussed with metallic bonding: a regular lattice of positive ions in a sea of delocalised electrons.

  • Properties & Why:
    • Good conductors of electricity and heat: The delocalised electrons are free to move and carry charge or heat energy through the metal.
    • Malleable (can be hammered into shape) and Ductile (can be drawn into wires): When you apply force, the layers of ions can slide over each other without breaking the metallic bond, because the electron "sea" flows around them.
    • High Melting & Boiling Points: Strong attraction between positive ions and the electron sea requires a lot of energy to overcome.
Key Takeaway

The link between micro and macro is all about the type of bonding and the structure it creates. Strong forces holding particles together (like in giant ionic, covalent, and metallic structures) lead to high melting points. Weak forces between particles (like in simple molecular structures) lead to low melting points.


Part 3: Chemical Formulae & Moles - The Language of Chemistry

Writing Chemical Formulae

A chemical formula is a shorthand way to show which elements are in a compound and how many atoms of each there are.

For Ionic Compounds:

The key is that the overall charge must be zero. Use the "cross-over" method!

Step-by-step example: Aluminium Oxide

  1. Write the symbols and charges of the ions. Aluminium is in Group III, so its ion is Al³⁺. Oxygen is in Group VI, so its ion is O²⁻.
  2. "Cross over" the numbers (not the signs). The '3' from Al³⁺ goes to O. The '2' from O²⁻ goes to Al.
  3. Write the formula: Al₂O₃.
  4. Check: (2 x +3) + (3 x -2) = +6 - 6 = 0. Perfect!
For Covalent Compounds:

The name often tells you the formula using prefixes (mono-, di-, tri-, tetra-, etc.).

  • Carbon dioxide = C with 2 oxygens = CO₂
  • Dinitrogen tetraoxide = 2 nitrogens with 4 oxygens = N₂O₄

Formula Mass and Relative Molecular Mass

These terms tell you the "weight" of a formula unit or a molecule compared to a standard (the ¹²C atom). The process is the same for both!

  • Relative Molecular Mass (Mr): Used for covalent molecules.
  • Formula Mass: Used for ionic compounds (since they don't have individual molecules).

How to calculate: Add up the relative atomic masses (Ar) of all the atoms in the formula.

Example: Water (H₂O)

Ar of H = 1.0; Ar of O = 16.0
Mr of H₂O = (2 × 1.0) + 16.0 = 18.0

Example: Magnesium Chloride (MgCl₂)

Ar of Mg = 24.3; Ar of Cl = 35.5
Formula Mass of MgCl₂ = 24.3 + (2 × 35.5) = 24.3 + 71.0 = 95.3


Part 4: Diving Deeper - Molecular Shapes and Polarity

Let's zoom back in on those simple molecules. It turns out their 3D shape is super important! The shape is determined by the electron pairs around the central atom, which repel each other to get as far apart as possible. This is the core idea of the VSEPR theory (Valence Shell Electron Pair Repulsion theory).

Predicting Shapes

Think of it like tying balloons together. The balloons (electron pairs) will naturally spread out to give each other space.

  • 2 Electron Pairs: Spread out to 180°. Shape = Linear. (e.g., CO₂)
  • 3 Electron Pairs: Spread out to 120°. Shape = Trigonal Planar. (e.g., BF₃)
  • 4 Electron Pairs: Spread out to 109.5°. Shape = Tetrahedral. (e.g., CH₄)
  • 5 Electron Pairs: Shape = Trigonal Bipyramidal. (e.g., PCl₅)
  • 6 Electron Pairs: Shape = Octahedral. (e.g., SF₆)

Note: Lone pairs of electrons (non-bonding pairs) also repel, and they actually repel more strongly than bonding pairs. This can alter the shape, like in Water (bent) and Ammonia (trigonal pyramidal).

Bond Polarity and Electronegativity

In a covalent bond, electrons are shared. But is the sharing always equal? No!

Electronegativity is a measure of how strongly an atom pulls shared electrons towards itself. Think of it as a "tug-of-war" for electrons.

  • If two identical atoms are bonded (e.g., Cl-Cl), the pull is equal. The bond is non-polar.
  • If two different atoms are bonded (e.g., H-Cl), one atom (Cl) pulls harder. The electrons spend more time near the chlorine. This creates a polar bond, with a small negative charge (δ-) on the Cl and a small positive charge (δ+) on the H.

Molecular Polarity: The Big Picture

A molecule can have polar bonds but still be non-polar overall! It all depends on the shape.

Think about two people pulling a box.

  • Carbon Dioxide (CO₂): It's linear (O=C=O). Both oxygens pull electrons away from the carbon with equal strength, but in opposite directions. The pulls cancel out! So, CO₂ has polar bonds but is a non-polar molecule.
  • Water (H₂O): It has a bent shape. The two H-O bonds are polar, and because of the shape, the pulls don't cancel out. The oxygen side is slightly negative (δ-) and the hydrogen side is slightly positive (δ+). Water is a polar molecule.

Key idea: If the polar bonds are arranged symmetrically, they cancel out, and the molecule is non-polar (like CH₄ or BF₃). If they are arranged asymmetrically, the molecule is polar (like H₂O, NH₃, or CHCl₃).

Key Takeaway

Molecular shape is crucial! It determines whether the effects of polar bonds cancel out, which in turn decides if the entire molecule is polar or non-polar. This property has a huge effect on how molecules interact with each other.


Part 5: The Forces BETWEEN Molecules

We mentioned these weak forces earlier. They are responsible for holding simple molecules together in a liquid or solid state. They are much, much weaker than the covalent bonds inside the molecules.

Van der Waals' Forces

This is the basic, weakest type of intermolecular force. It exists between ALL molecules, both polar and non-polar. It's caused by the random movement of electrons, which creates temporary, tiny dipoles that weakly attract other molecules.

  • Strength depends on: The number of electrons. More electrons (i.e., bigger molecules) mean stronger Van der Waals' forces.
  • This explains why boiling points increase as you go down a group (e.g., from F₂ to I₂). Iodine is a much bigger molecule with more electrons, so the forces between I₂ molecules are stronger and require more energy to break.

Hydrogen Bonding

This isn't a real bond! It's an extra-strong type of intermolecular force. It's like the "Velcro" of chemistry.

It only happens when: A hydrogen atom is bonded to a very electronegative atom (Nitrogen, Oxygen, or Fluorine) and is attracted to another N, O, or F atom on a nearby molecule.

  • Why it's so strong: The H-N, H-O, or H-F bond is extremely polar. The hydrogen is very δ+ and the N/O/F is very δ-. The attraction between these partial charges is very strong.
  • Consequences: Substances with hydrogen bonding (like water H₂O, ammonia NH₃, and hydrogen fluoride HF) have much higher boiling points than you would expect for their size. This is why water is a liquid at room temperature, while similarly sized molecules are gases!

Final Summary: The Big Picture

Let's tie it all together with a summary table. This is the ultimate macro-micro link!

| Property | Giant Ionic | Simple Molecular | Giant Covalent | Giant Metallic | | :--- | :--- | :--- | :--- | :--- | | Microscopic View | | | | | | Particles | Positive & Negative Ions | Molecules | Atoms | Positive Ions & Delocalised Electrons | | Bonding | Strong Ionic Bonds | Strong covalent bonds within molecules | Strong Covalent Bonds | Strong Metallic Bonds | | Forces to Overcome | Strong electrostatic forces | Weak intermolecular forces | Strong covalent bonds | Strong metallic bonds | | Macroscopic View | | | | | | M.P. / B.P. | High | Low | Very High | High | | Electrical Conductivity | Solid: No
Molten/Aq: Yes | No | No (except graphite) | Yes (solid & liquid) | | Hardness | Hard & Brittle | Soft | Very Hard (except graphite) | Malleable & Ductile | | Example | NaCl | CO₂, H₂O | Diamond, SiO₂ | Iron (Fe), Copper (Cu) |

Congratulations! You've just journeyed from the level of single electrons all the way up to the properties of everyday materials. By understanding the bonding, structure, and shape at the microscopic level, you can now predict and explain the macroscopic world. Keep reviewing these core ideas, and you'll build a fantastic foundation in chemistry!