🌟 Welcome to the World of Covalent Bonding! 🌟
Hello future chemists! This chapter is part of your journey through the Principles of Chemistry, and it’s one of the most fundamental concepts you’ll learn: how atoms stick together to form almost everything around you.
Don't worry if this topic seems tricky at first. We will break down complex molecules into simple diagrams, and by the end of these notes, you’ll be drawing atoms sharing electrons like a professional!
What you will learn in this chapter:
- What a covalent bond is and why atoms form them.
- How to draw 'dot-and-cross' diagrams for molecules.
- The difference between simple molecules and giant covalent structures.
- How structure dictates the properties of materials like diamond and water.
1. The Basics of Covalent Bonding
1.1 What is a Covalent Bond?
Remember the driving force behind all chemical reactions? Stability! Atoms are happiest when their outer electron shells are full, usually with eight electrons (this is called the Octet Rule).
When non-metal atoms react together, they often can't simply transfer electrons (like in ionic bonding). Instead, they find a clever solution: they share electrons.
The Definition:
A covalent bond is the strong electrostatic attraction between the nuclei of two atoms and the pair (or pairs) of electrons that they share between them.
Analogy: Sharing is Caring!
Imagine you and a friend each have one shoe, and you need a pair to go running. Neither of you will give up your shoe entirely (that would be ionic bonding), so you agree to share the two shoes between you. You both get to wear the "shared pair" of shoes. In chemistry, the "shared shoes" are a pair of electrons.
1.2 Who Forms Covalent Bonds?
- Covalent bonds form primarily between non-metal atoms.
- Examples: Carbon (C), Hydrogen (H), Oxygen (O), Chlorine (Cl), Nitrogen (N).
When atoms share electrons, they form a group called a molecule. A molecule is the smallest unit of a covalently bonded substance. Example: Water, H2O, is a molecule.
Key Takeaway
Covalent bonds occur when non-metals share outer shell electrons to achieve a stable, full outer shell, forming a molecule.
2. Drawing Covalent Molecules (Dot-and-Cross Diagrams)
Dot-and-cross diagrams are the essential way we visualize covalent bonding. They show only the electrons in the outer shell (the valence electrons) and which ones are being shared.
2.1 Step-by-Step Guide to Drawing Diagrams
To draw a diagram, you must know the group number of the elements, as this tells you how many outer shell electrons they have (their valency).
Step 1: Count Electrons. Determine the number of valence electrons for each atom (e.g., Hydrogen is Group 1, so 1 electron; Oxygen is Group 6, so 6 electrons).
Step 2: Assign Symbols. Draw the atoms overlapping where the electrons will be shared. Use dots (•) for the electrons from one atom and crosses (x) for the electrons from the other.
Step 3: Share Electrons. Place electrons in the overlapping area (the shared region) so that every atom achieves a full outer shell (2 for H; 8 for C, O, N, Cl).
Step 4: Place Unshared Electrons. Put any remaining electrons (lone pairs) back into the outer shells of their respective atoms.
2.2 Examples of Covalent Bonds
Example A: Hydrogen Chloride (HCl) - Single Bond
Hydrogen (H) needs 1 electron. Chlorine (Cl) needs 1 electron (it has 7 already).
They share one pair of electrons (1 from H, 1 from Cl).
A single shared pair is called a single covalent bond.
Example B: Oxygen (\(O_2\)) - Double Bond
Oxygen is in Group 6, so it has 6 outer electrons. It needs 2 more to make 8.
Since both O atoms need 2 electrons, they must share two pairs of electrons (a total of four shared electrons).
A double shared pair is called a double covalent bond.
Did you know? Carbon Dioxide (\(CO_2\)) also contains two double bonds!
Example C: Nitrogen (\(N_2\)) - Triple Bond
Nitrogen is in Group 5, so it has 5 outer electrons. It needs 3 more to make 8.
Since both N atoms need 3 electrons, they must share three pairs of electrons (a total of six shared electrons).
A triple shared pair is called a triple covalent bond.
! Common Mistake Alert !
Always double-check that every atom in your diagram has achieved a full shell! Count the electrons in the shared region for BOTH atoms. For Chlorine, the total must be 8 (6 lone + 2 shared).
Quick Review: Types of Covalent Bonds
Single Bond: 1 shared pair (e.g., \(H_2\), \(HCl\))
Double Bond: 2 shared pairs (e.g., \(O_2\), \(CO_2\))
Triple Bond: 3 shared pairs (e.g., \(N_2\))
3. Properties of Simple Molecular Substances
Most covalently bonded substances, like water, sugar, and carbon dioxide, exist as individual, small molecules. These are called Simple Molecular Structures.
3.1 The Key to Properties: Forces
This is the most important concept when discussing properties:
- Intramolecular Forces: The strong covalent bonds within the molecule (holding H to O in \(H_2O\)). These are very strong!
- Intermolecular Forces (IMFs): The very weak forces between neighbouring molecules (holding one \(H_2O\) molecule next to another \(H_2O\) molecule).
Analogy: The Lego Bricks
Think of a molecule as a Lego brick. The plastic holding the brick together (the covalent bond) is very strong. The weak force holding one Lego brick to another (the IMF) is easy to break.
When we melt or boil a substance, we are only breaking the weak intermolecular forces, not the strong covalent bonds inside the molecule.
3.2 Melting and Boiling Points
- Result: Simple molecular substances have very low melting and boiling points.
- Explanation: Because the intermolecular forces are weak, only a small amount of heat energy is needed to overcome them and separate the molecules, turning the liquid into a gas (or solid into liquid).
- Example: \(CO_2\) exists as a gas at room temperature, and water boils at 100 °C (relatively low).
3.3 Electrical Conductivity
- Result: Simple molecular substances do not conduct electricity, either as solids or liquids.
- Explanation: To conduct electricity, a substance needs mobile charged particles (ions or delocalised electrons). Covalent molecules are uncharged overall, and their electrons are locked tightly in the shared bonds (they are localised).
3.4 Solubility
This is a complex area, but often simplified as "like dissolves like."
- Most simple covalent molecules (like oil) are non-polar and will dissolve well in other non-polar solvents (like hexane).
- Small polar molecules (like sugar or ethanol) and highly polar molecules (like water) can dissolve in water, but only if the substance can form bonds or attractions with the water molecules.
Key Takeaway
Simple molecular substances are defined by having weak forces between molecules, leading to low melting/boiling points and a lack of electrical conductivity.
4. Giant Covalent Structures (Macromolecules)
Not all non-metals form small, simple molecules. Some form enormous structures where every atom is bonded covalently to its neighbours, creating a giant network. These are called Giant Covalent Structures (or Macromolecules).
Since they are held together entirely by strong covalent bonds, their properties are the exact opposite of simple molecules!
4.1 Diamond (Carbon)
Diamond is made purely of carbon atoms.
Structure:
- Each carbon atom is covalently bonded to four other carbon atoms in a rigid, 3D tetrahedral network.
- There are no weak intermolecular forces—the entire crystal is one giant molecule.
Properties and Uses:
- High Melting/Boiling Point: Extreme heat is required to break the vast number of strong covalent bonds. Diamond is incredibly hard to melt.
- Hardness: The rigid tetrahedral structure makes it the hardest known natural substance. (Used for cutting tools and drills).
- Electrical Conductivity: Does not conduct. All outer electrons are fixed in the four strong bonds (no free electrons).
4.2 Graphite (Carbon)
Graphite is another form of carbon, but its structure is completely different.
Structure:
- Each carbon atom is covalently bonded to only three other carbon atoms.
- These bonds form flat, hexagonal rings arranged in layers.
- The forces within the layers are strong covalent bonds, but the forces between the layers are weak intermolecular forces.
Properties and Uses:
- Soft/Slippery: The weak forces between layers mean the layers can easily slide over each other. (Used as a lubricant and in pencils).
- Electrical Conductivity: Conducts electricity. Since carbon only bonds to three neighbours, the fourth valence electron from each atom is delocalised (free to move) between the layers. These mobile electrons carry charge. (Used for electrodes).
- High Melting Point: It still has a high melting point because you need to break the strong covalent bonds within the layers.
4.3 Silicon Dioxide (\(SiO_2\))
Silicon dioxide (the main component of sand and quartz) has a structure very similar to diamond.
- Each Silicon atom is bonded to four Oxygen atoms, and each Oxygen atom is bonded to two Silicon atoms, forming a giant tetrahedral lattice.
- This results in very high melting points and extreme hardness (like diamond).
Did you know? Despite both being pure carbon, the dramatically different structures of diamond and graphite lead to dramatically different uses—one cuts metal, the other is used in pencil lead!
Key Takeaway
Giant covalent structures have strong covalent bonds throughout the entire structure. This results in very high melting points and hardness. Graphite is the unique exception because its layered structure and delocalised electrons allow it to conduct electricity and be soft.
📚 Chapter Review: Covalent Bonding Summary
Final Encouragement:
You've tackled the complexities of sharing electrons! Remember to always ask yourself: "Is it a strong bond or a weak force I need to break?" That question will unlock the properties of almost any substance! Keep practising those dot-and-cross diagrams—they are crucial for success in this topic!
| Feature | Simple Molecular (e.g., Water, CO2) | Giant Covalent (e.g., Diamond, SiO2) |
|---|---|---|
| Bonding | Covalent bonds within molecules | Covalent bonds throughout the entire structure |
| Forces Broken | Weak Intermolecular Forces (IMFs) | Strong Covalent Bonds |
| M.P. / B.P. | Low | Very High |
| Hardness | Soft liquids or gases | Very hard (except graphite) |
| Conductivity | No (localized electrons) | No (except Graphite, due to delocalized electrons) |