Introduction: Why Atoms Stick Together (And Why It Matters!)

Hello! This chapter is all about how non-metal atoms join up to create the thousands of everyday substances around us, from the air we breathe to the plastic bottles we use. After learning about the intense electrostatic forces in ionic bonds (metals + non-metals), we now look at how non-metals bond with *each other*.

When you finish these notes, you'll be able to explain exactly why water boils easily, why diamond is the hardest substance, and why graphite can conduct electricity—all by looking at how their atoms share electrons!

Section 1: Covalent Bonds in Simple Molecules (2.5 Core)

What is a Covalent Bond?

A covalent bond forms when two atoms share a pair of electrons. This usually happens between two non-metal atoms.

Why do they share? Because sharing allows both atoms to achieve the stable electron structure of a noble gas (usually having a full outer shell, or 8 electrons—the octet rule).

Analogy: The Handshake

Think of atoms wanting eight electrons in their outer shell (except Hydrogen, which only wants two). If Atom A needs one electron and Atom B needs one electron, they don't fight over it; they "hold hands" or share that pair of electrons. This shared pair forms the covalent bond.

Drawing Covalent Bonds: Dot-and-Cross Diagrams (Core Examples)

We use dot-and-cross diagrams to show the movement and arrangement of the outer electrons. Remember to only draw the outer electron shell!

Step-by-step Guide to Drawing Covalent Bonds:

  1. Identify the atoms involved and their group numbers to find out how many outer (valence) electrons they start with.
  2. Determine how many electrons each atom needs to gain a full outer shell (usually 8).
  3. Draw the atoms overlapping, showing only the outer shells.
  4. Place the shared electrons (one dot, one cross per atom contributing) in the overlapping area.
  5. Fill in any remaining unshared electrons on the outer shells.
  6. Check: Count the total electrons around *each* atom to ensure they all have a stable noble gas configuration (2 for H, 8 for most others).
Example 1: Hydrogen Chloride (HCl)
  • Hydrogen (Group I) needs 1 electron.
  • Chlorine (Group VII) needs 1 electron.
  • They share one pair of electrons, forming a single covalent bond.
Example 2: Methane ($\text{CH}_4$)
  • Carbon (Group IV) needs 4 electrons.
  • Each of the four Hydrogen atoms needs 1 electron.
  • Carbon shares one electron with each Hydrogen, resulting in four single covalent bonds.

Quick Review: Core Dot-and-Cross Molecules

  • Hydrogen ($\text{H}_2$): 1 shared pair (single bond).
  • Chlorine ($\text{Cl}_2$): 1 shared pair (single bond).
  • Water ($\text{H}_2\text{O}$): Two single bonds. Oxygen shares 1 electron with each H.
  • Methane ($\text{CH}_4$): Four single bonds. Carbon shares 1 electron with each H.
  • Ammonia ($\text{NH}_3$): Three single bonds. Nitrogen shares 1 electron with each H.
  • Hydrogen Chloride ($\text{HCl}$): 1 shared pair (single bond).

Beyond Single Bonds (2.5 Supplement)

Some atoms need to share more than one pair of electrons to reach stability.

  • Double Bond: Two atoms share two pairs of electrons (total of 4 shared electrons).
    Example: Oxygen Gas ($\text{O}_2$). Each oxygen atom needs 2 electrons, so they share two pairs.
  • Triple Bond: Two atoms share three pairs of electrons (total of 6 shared electrons).
    Example: Nitrogen Gas ($\text{N}_2$). Each nitrogen atom needs 3 electrons, so they share three pairs.
  • Other Supplement examples include: Methanol ($\text{CH}_3\text{OH}$), Ethene ($\text{C}_2\text{H}_4$), and Carbon Dioxide ($\text{CO}_2$).
Did you know?

The nitrogen triple bond in $\text{N}_2$ is one of the strongest bonds known in chemistry! This is why nitrogen gas is very unreactive—it takes huge amounts of energy to break those three shared pairs.

Key Takeaway for Covalent Bonding: Covalent bonds involve the *sharing* of electrons between non-metals to achieve full outer shells. We use dot-and-cross diagrams to visualize this sharing.


Section 2: Properties of Simple Molecular Compounds (2.5 Core & Supplement)

Compounds formed by covalent bonds in small groups (like $\text{H}_2\text{O}$ or $\text{CO}_2$) are called simple molecular structures. Their properties are defined by the type of forces holding the structure together.

1. Low Melting Points and Boiling Points

Simple molecular compounds are usually gases, liquids, or low-melting point solids at room temperature.

  • Observation (Core): They have low melting points and boiling points.
  • Explanation (Supplement): Why?
    • Inside the molecule (e.g., the O-H bonds in water), the covalent bonds are very strong.
    • However, the forces between the individual molecules (these are called intermolecular forces) are very weak.

Analogy: Lego Bricks. A simple molecule is like a single Lego brick—it's very strong and hard to break. But the forces between one brick and the next (the intermolecular forces) are weak. When you heat water, you are only overcoming the weak forces that hold the molecules near each other, not the strong covalent bonds inside the water molecule itself. This takes very little energy, leading to low melting and boiling points.

2. Electrical Conductivity

  • Observation (Core): Simple molecular compounds are poor electrical conductors (they do not conduct electricity) in any state (solid, liquid, or gas).
  • Explanation (Supplement): For electricity to flow, you need mobile charged particles (either ions or delocalised electrons). Simple molecular structures have neither! Their electrons are stuck in fixed shared pairs between atoms, and the molecules themselves are neutral.

Quick Review: Simple Molecules

Structure: Strong covalent bonds *within* molecules; Weak intermolecular forces *between* molecules.

Properties: Low M.P./B.P. (easy to break weak forces); Poor conductor (no mobile ions or electrons).

Key Takeaway for Simple Molecular Properties: These compounds are defined by their weak intermolecular forces, which means they require little energy to change state and cannot conduct electricity.


Section 3: Giant Covalent Structures (Macromolecules) (2.6 Core & Supplement)

Not all covalent substances form simple little molecules. Some atoms bond together in massive, repeating structures called giant molecular structures or macromolecules.

In these structures, every single atom is held to its neighbours by strong covalent bonds, creating a huge network. This gives them very different properties compared to simple molecules.

3.1 Diamond (Core)

Diamond is an allotrope of carbon (different structural forms of the same element).

  • Structure: Each carbon atom is bonded to four other carbon atoms in a tetrahedral arrangement (a 3D pyramid shape). This creates a vast, rigid giant covalent lattice.
  • Bonding: All outer electrons are used up in the four strong covalent bonds. There are no delocalised electrons.
  • Properties and Uses:
    • Extremely hard and strong: Because of the strong bonds throughout the entire structure.
    • High Melting Point: It takes immense energy to break all those strong covalent bonds.
    • Does not conduct electricity: Because there are no free (delocalised) electrons.
    • Use: Used in cutting tools (e.g., drills and saws) due to its extreme hardness.

3.2 Graphite (Core)

Graphite is another allotrope of carbon, but its structure is completely different, giving it unique properties.

  • Structure: Carbon atoms are arranged in flat, hexagonal layers. In each layer, a carbon atom is bonded to only three other carbon atoms.
  • Bonding: Since carbon has four outer electrons, and only uses three to form strong bonds within the layer, the fourth electron is free to move between the layers. These are delocalised electrons.
  • Properties and Uses:
    • Conducts electricity: This is rare for a non-metal! The delocalised electrons act as charge carriers, allowing current to flow along the layers.
    • Soft and slippery: The strong covalent layers are held together only by weak intermolecular forces. These layers can easily slide over one another.
    • Uses:
      1. As a lubricant (due to its slipperiness, reducing friction).
      2. As electrodes (in batteries and electrolysis) because it conducts electricity and is chemically inert.

3.3 Silicon(IV) Oxide ($\text{SiO}_2$) (Supplement)

Silicon(IV) oxide, commonly known as silica or sand, is also a giant covalent structure.

  • Structure: Similar to diamond. Each silicon atom is covalently bonded to four oxygen atoms, and each oxygen atom bridges two silicon atoms. It forms a vast, network lattice.
  • Bonding: Strong covalent bonds throughout.
  • Similarity to Diamond (Supplement): Because $\text{SiO}_2$ has a giant, rigid network of strong covalent bonds just like diamond, it shares the properties of being very hard and having a very high melting point.

Common Mistake Alert!

Don't confuse the bonding in graphite with the bonding in metals! Both have delocalised electrons, but in graphite, the conduction only happens *along the layers*, and the forces holding the layers together are weak (not the strong metallic bond).

Key Takeaway for Giant Covalent Structures: They are defined by having strong covalent bonds linking *all* atoms in a massive lattice. This leads to exceptional hardness and very high melting points. Graphite is the unique exception as its layered structure gives it electrical conductivity and softness.


Study Summary: Comparing Structure and Properties

Understanding the bond type and structure is the key to predicting properties!

Covalent Bonding Types and Their Properties

Structure Type Example Forces Overcome by Melting/Boiling M.P. / B.P. Electrical Conductivity
Simple Molecular Water, $\text{CO}_2$ Weak intermolecular forces Low Poor (no mobile charged particles)
Giant Covalent Diamond, $\text{SiO}_2$ Strong covalent bonds (the whole lattice) Very High Poor (no mobile charged particles)
Giant Covalent (Graphite) Graphite Strong covalent bonds (within layers) Very High Good (mobile delocalised electrons)