Welcome to Chemical Bonding!
Have you ever wondered why water sticks together as a liquid, why table salt forms neat little crystals, or why copper wire can bend without snapping? The secret behind all of these is chemical bonding!
In this chapter, we will look at how and why atoms join together. Don't worry if science has felt overwhelming before — we are going to break everything down into bite-sized, easy-to-remember steps.
Did you know? Almost everything you can touch in the universe is held together by chemical bonds. Without them, all matter would just drift apart as individual atoms!
The Golden Rule of Chemistry: Why Do Atoms Bond?
Atoms are just like us: they want to be stable and content! An atom achieves stability when it gets a full outer shell of electrons (usually \(8\) electrons, which is often called a stable octet, or \(2\) electrons for hydrogen and helium).
The Noble Gases in Group 0 (like helium, neon, and argon) already have full outer shells. They are chemically unreactive and stable. Every other atom on the Periodic Table wants to achieve that same noble gas electron arrangement.
To get a full outer shell, atoms can do one of three things:
• Lose electrons (metals)
• Gain electrons (non-metals)
• Share electrons (non-metals)
There are three main types of bonding you need to know for your exam:
1. Ionic Bonding: between a metal and a non-metal.
2. Covalent Bonding: between non-metal atoms only.
3. Metallic Bonding: inside metals only.
Quick Takeaway: Atoms bond to achieve a full outer shell of electrons, making them stable like noble gases.
1. Ionic Bonding: The Giving and Taking of Electrons
What is an Ionic Bond?
Ionic bonding happens between a metal atom and a non-metal atom.
• Metals have only a few electrons in their outer shell, so it is easiest for them to lose electrons to empty that shell.
• Non-metals have outer shells that are nearly full, so it is easiest for them to gain electrons to complete their shell.
Analogy: Imagine one person has an item they want to get rid of, and another person desperately wants it. When the transfer happens, both people are happy and linked together as friends!
Forming Ions
When an atom loses or gains electrons, it is no longer neutral; it becomes a charged particle called an ion.
• Positive ion (Cation): Formed when a metal atom loses negative electrons. (e.g., \(Na^+\), \(Mg^{2+}\)).
• Negative ion (Anion): Formed when a non-metal atom gains negative electrons. (e.g., \(Cl^-\), \(O^{2-}\)).
Memory Trick: A ca+ion has a "+" in the middle, so it is positive! An A-Negative-ION (anion) is negative!
Step-by-Step Example: Sodium Chloride (\(NaCl\))
Let's look at how table salt is made from sodium and chlorine:
1. A sodium atom has the electronic configuration \(2,8,1\). It has \(1\) outer electron.
2. A chlorine atom has the electronic configuration \(2,8,7\). It has \(7\) outer electrons.
3. The sodium atom transfers its \(1\) outer electron to the chlorine atom.
4. Sodium now has the electronic configuration \(2,8\) and becomes a positive sodium ion, written as \(Na^+\).
5. Chlorine now has the electronic configuration \(2,8,8\) and becomes a negative chloride ion, written as \(Cl^-\).
6. The oppositely charged ions (\(Na^+\) and \(Cl^-\)) strongly attract one another.
Official Definition: An ionic bond is the strong electrostatic attraction between oppositely charged ions.
Another Example: Magnesium Oxide (\(MgO\))
• Magnesium (\(2,8,2\)) loses \(2\) electrons to become \(Mg^{2+}\) (\(2,8\)).
• Oxygen (\(2,6\)) gains \(2\) electrons to become \(O^{2-}\) (\(2,8\)).
• The strong attraction between \(Mg^{2+}\) and \(O^{2-}\) forms the compound magnesium oxide (\(MgO\)).
Properties of Ionic Compounds
• High melting and boiling points: It takes a huge amount of heat energy to break the strong electrostatic forces holding the giant lattice of ions together.
• Conducting electricity: They cannot conduct electricity when solid because the ions are locked firmly in place. However, they can conduct electricity when molten (melted) or dissolved in water because the ions are free to move and carry the electric charge.
Common Mistake to Avoid: Never say that electrons move to carry current in an ionic solution! In ionic compounds, it is the free ions that move and carry the charge.
Quick Takeaway: Ionic bonding involves the transfer of electrons from metals to non-metals, forming positive and negative ions that attract each other strongly.
2. Covalent Bonding: Sharing is Caring!
What is a Covalent Bond?
Covalent bonding happens between non-metal atoms only. Non-metals all need to gain electrons to fill their outer shells, so none of them want to give electrons away. Instead, they share pairs of electrons.
Analogy: Imagine two students sitting next to each other in class. Neither has a complete set of coloured pencils, but by putting their pencils in the middle of the desk, both students have access to a full set!
Official Definition: A covalent bond is a shared pair of electrons.
Key Covalent Molecules to Know
1. Hydrogen (\(H_2\)) — Single Bond:
Each hydrogen atom has \(1\) electron and needs \(1\) more to fill its first shell (which holds \(2\)). Two hydrogen atoms share \(1\) pair of electrons, forming a single covalent bond (\(H-H\)).
2. Chlorine (\(Cl_2\)) — Single Bond:
Each chlorine atom has \(7\) outer electrons (\(2,8,7\)) and needs \(1\) more. Two chlorine atoms share \(1\) pair of electrons, so both reach a stable \(2,8,8\) structure (\(Cl-Cl\)).
3. Water (\(H_2O\)):
Oxygen has \(6\) outer electrons and needs \(2\) more. It shares \(1\) pair of electrons with one hydrogen atom, and another pair with a second hydrogen atom (\(H-O-H\)).
4. Methane (\(CH_4\)):
Carbon has \(4\) outer electrons and needs \(4\) more. It forms \(4\) single covalent bonds by sharing electrons with four separate hydrogen atoms.
5. Ammonia (\(NH_3\)):
Nitrogen has \(5\) outer electrons and needs \(3\) more. It shares electrons with three separate hydrogen atoms, forming \(3\) single covalent bonds.
6. Oxygen (\(O_2\)) — Double Bond:
Each oxygen atom has \(6\) outer electrons and needs \(2\) more. They share two pairs of electrons (a total of \(4\) electrons), forming a double covalent bond (\(O=O\)).
7. Nitrogen (\(N_2\)) — Triple Bond:
Each nitrogen atom has \(5\) outer electrons and needs \(3\) more. They share three pairs of electrons (a total of \(6\) electrons), forming a very strong triple covalent bond (\(N \equiv N\)).
Properties of Simple Covalent Molecules
• Low melting and boiling points: Although the covalent bonds inside the molecule are very strong, the forces between the separate molecules (called weak intermolecular forces) are very weak. It takes very little heat energy to pull the molecules apart from each other. That is why many covalent substances (like oxygen, nitrogen, and carbon dioxide) are gases at room temperature.
• Do not conduct electricity: Covalent molecules do not have any free electrons or ions to carry an electric charge.
Common Mistake to Avoid: When ice melts or water boils, the strong covalent bonds between \(H\) and \(O\) do not break! Only the weak forces between the water molecules are overcome.
Quick Takeaway: Covalent bonding involves sharing pairs of electrons between non-metal atoms. Simple molecules have low melting points and do not conduct electricity.
3. Metallic Bonding: A Sea of Electrons
What is Metallic Bonding?
Metallic bonding occurs inside metals (such as copper, iron, magnesium, and aluminium). Metal atoms pack closely together in a neat, regular pattern called a lattice.
The atoms release their outer shell electrons. These electrons are no longer tied to one specific atom — they are free to move throughout the entire structure. We call these delocalised electrons.
This leaves behind a regular lattice of positive metal ions surrounded by a "sea" of delocalised electrons.
Official Definition: A metallic bond is the attraction between the lattice of positive metal ions and the sea of delocalised electrons.
Properties of Metals Explained
• Good conductors of electricity and heat: The delocalised electrons are free to move throughout the structure and carry electrical charge or thermal energy quickly.
• High melting and boiling points: There is a strong attraction between the positive ions and the delocalised electrons, which requires large amounts of energy to overcome.
• Malleable (can be hammered into shape) and ductile (can be drawn into wires): Because the metal ions are arranged in neat layers, these layers can smoothly slide over each other when a force is applied, without breaking the metallic bond.
Quick Takeaway: Metallic bonding consists of regular layers of positive metal ions in a sea of delocalised electrons, making metals great conductors and easily bendable.
Summary Comparison: The Three Types of Bonding
Use this quick guide to review before an exam:
1. Ionic Bonding:
• Between: Metal + Non-metal
• Process: Electrons are transferred (lost and gained)
• Particles formed: Positive and negative ions
• Melting/Boiling Points: High
• Electrical Conductivity: Conducts only when molten or dissolved in water
2. Covalent Bonding:
• Between: Non-metal + Non-metal
• Process: Electrons are shared in pairs
• Particles formed: Neutral molecules
• Melting/Boiling Points: Low (for simple molecules)
• Electrical Conductivity: Does not conduct (poor/insulator)
3. Metallic Bonding:
• Between: Metal atoms only
• Process: Outer electrons become delocalised
• Particles formed: Positive metal ions in a sea of delocalised electrons
• Melting/Boiling Points: High
• Electrical Conductivity: Always conducts (solid and liquid)
Top Exam Tips
• Always look at the elements in the formula first! Metal + Non-metal = Ionic. Non-metal + Non-metal = Covalent. Metal alone = Metallic.
• If a question asks why metals conduct electricity, the key phrase is: delocalised electrons are free to move and carry charge.
• If a question asks why ionic solutions conduct electricity, the key phrase is: ions are free to move and carry charge.
• Remember to state that simple covalent substances have low boiling points due to weak forces between molecules (intermolecular forces), not weak covalent bonds!