Welcome to 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!
Don't worry if chemistry sometimes feels like learning a whole new language. In this chapter, we are going to break down everything step-by-step. By the end of these notes, you will understand exactly how atoms join together, how to draw their electrons, and how to avoid the classic traps that catch students out in exams.
1. Why Do Atoms Bond?
Atoms are just like us—they want to be stable and comfortable! For an atom, stability means having a full outer shell of electrons (often called an octet, meaning 8 electrons in the outer shell, matching the electronic configuration of a noble gas in Group 0/8).
Most atoms do not have a full outer shell on their own. To achieve one, they must interact with other atoms. There are three main types of bonding you need to know for CCEA GCSE Chemistry Unit C1:
• Ionic Bonding: Occurs between a metal and a non-metal (electrons are transferred).
• Covalent Bonding: Occurs between non-metal atoms (electrons are shared).
• Metallic Bonding: Occurs in metals and alloys (positive metal ions in a sea of delocalised electrons).
Quick Summary Trick to Identify the Bonding Type:
• Metal + Non-metal? \(\rightarrow\) Ionic
• Non-metal + Non-metal? \(\rightarrow\) Covalent
• Metal only? \(\rightarrow\) Metallic
2. Ionic Bonding
What is an Ion?
An ion is an atom (or group of atoms) that has gained an electrical charge by gaining or losing one or more electrons.
• Positive Ions (Cations): Metal atoms lose their outer shell electrons to form positive ions.
Memory tip: Think of the 't' in ca+ion as a plus sign (+).
Examples:
Sodium atom \(\text{Na}\ (2,8,1)\) loses \(1\) electron \(\rightarrow\) Sodium ion \(\text{Na}^+\ (2,8) + \text{e}^-\)
Magnesium atom \(\text{Mg}\ (2,8,2)\) loses \(2\) electrons \(\rightarrow\) Magnesium ion \(\text{Mg}^{2+}\ (2,8) + 2\text{e}^-\)
• Negative Ions (Anions): Non-metal atoms gain electrons into their outer shell to form negative ions. When a non-metal becomes a simple negative ion, its name changes to end in -ide.
Memory tip: A Negative Ion = Anion.
Examples:
Chlorine atom \(\text{Cl}\ (2,8,7)\) gains \(1\) electron \(\rightarrow\) Chloride ion \(\text{Cl}^-\ (2,8,8)\)
Oxygen atom \(\text{O}\ (2,6)\) gains \(2\) electrons \(\rightarrow\) Oxide ion \(\text{O}^{2-}\ (2,8)\)
The Official CCEA Definition of an Ionic Bond
An ionic bond is the strong electrostatic attraction between oppositely charged ions (positive cations and negative anions) arranged in a regular three-dimensional ionic lattice.
Drawing Dot-and-Cross Diagrams for Ionic Compounds
When drawing ionic dot-and-cross diagrams in your exam, remember these essential rules:
1. Show the transfer of electrons from the metal to the non-metal.
2. Use dots for one atom's electrons and crosses for the other.
3. Put the resulting ions inside square brackets: \([\ ]\).
4. Write the charge outside the top right of the bracket (e.g. \([\text{Na}]^+\) and \([\text{Cl}]^-\)).
Key Examples Required by CCEA:
• Sodium chloride (\(\text{NaCl}\)): One \(\text{Na}\) atom loses \(1\) electron to form \([\text{Na}]^+\). One \(\text{Cl}\) atom gains that \(1\) electron to form \([\text{Cl}]^-\).
• Magnesium oxide (\(\text{MgO}\)): One \(\text{Mg}\) atom loses \(2\) electrons to form \([\text{Mg}]^{2+}\). One \(\text{O}\) atom gains those \(2\) electrons to form \([\text{O}]^{2-}\).
• Magnesium chloride (\(\text{MgCl}_2\)): One \(\text{Mg}\) atom loses \(2\) electrons to form \([\text{Mg}]^{2+}\). Two separate \(\text{Cl}\) atoms each gain \(1\) electron to form two \([\text{Cl}]^-\) ions.
• Sodium oxide (\(\text{Na}_2\text{O}\)): Two \(\text{Na}\) atoms each lose \(1\) electron to form two \([\text{Na}]^+\) ions. One \(\text{O}\) atom gains those \(2\) electrons to form \([\text{O}]^{2-}\).
Molecular (Polyatomic) Ions You Must Memorise
Some ions are made of a small group of atoms joined together with an overall charge. You must know these formulas and charges off by heart:
• Hydroxide: \(\text{OH}^-\)
• Nitrate: \(\text{NO}_3^-\)
• Sulfate: \(\text{SO}_4^{2-}\)
• Carbonate: \(\text{CO}_3^{2-}\)
• Ammonium: \(\text{NH}_4^+\) (Note: this is a positive molecular ion!)
Key Takeaway for Ionic Bonding: Metals lose electrons to become positive cations, non-metals gain electrons to become negative anions, and opposite charges attract tightly in a giant lattice.
3. Covalent Bonding
The Official CCEA Definition of a Covalent Bond
A covalent bond is the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonding atoms.
Covalent bonds form when non-metal atoms share pairs of electrons to achieve full outer electron shells.
Diatomic Molecules
A diatomic molecule consists of two atoms chemically bonded together. The seven diatomic non-metal elements you must know are:
\(\text{H}_2\), \(\text{N}_2\), \(\text{O}_2\), \(\text{F}_2\), \(\text{Cl}_2\), \(\text{Br}_2\), \(\text{I}_2\)
Memory trick: Have No Fear Of Ice Cold Beer (\(\text{H}_2, \text{N}_2, \text{F}_2, \text{O}_2, \text{I}_2, \text{Cl}_2, \text{Br}_2\)).
Types of Covalent Bonds & Required Dot-and-Cross Diagrams
1. Single Covalent Bonds (1 shared pair = 2 electrons shared)
• Hydrogen (\(\text{H}_2\)): Each \(\text{H}\) shares \(1\) electron, giving \(1\) shared pair (\(\text{H}-\text{H}\)). Both achieve a stable outer shell of \(2\) electrons.
• Chlorine (\(\text{Cl}_2\)): Each \(\text{Cl}\) shares \(1\) electron, giving \(1\) shared pair and \(3\) unshared pairs (lone pairs) on each atom.
• Hydrogen chloride (\(\text{HCl}\)): \(\text{H}\) shares \(1\) electron with \(\text{Cl}\), forming \(1\) shared pair.
• Water (\(\text{H}_2\text{O}\)): Central \(\text{O}\) atom forms two single covalent bonds (one with each \(\text{H}\)), leaving \(2\) lone pairs on the oxygen.
• Ammonia (\(\text{NH}_3\)): Central \(\text{N}\) atom forms three single covalent bonds (one with each \(\text{H}\)), leaving \(1\) lone pair on the nitrogen.
• Methane (\(\text{CH}_4\)): Central \(\text{C}\) atom forms four single covalent bonds (one with each \(\text{H}\)).
2. Double Covalent Bonds (2 shared pairs = 4 electrons shared)
• Oxygen molecule (\(\text{O}_2\)): Each oxygen atom needs \(2\) electrons to complete its shell, so they share \(2\) pairs of electrons (\(\text{O}=\text{O}\)).
• Carbon dioxide (\(\text{CO}_2\)): Central \(\text{C}\) atom shares \(2\) pairs of electrons with each of the two \(\text{O}\) atoms (\(\text{O}=\text{C}=\text{O}\)).
3. Triple Covalent Bonds (3 shared pairs = 6 electrons shared)
• Nitrogen molecule (\(\text{N}_2\)): Each nitrogen atom has \(5\) outer electrons and needs \(3\) more. They share \(3\) pairs of electrons (\(\text{N}\equiv\text{N}\)).
CCEA Drawing Conventions for Covalent Molecules:
• Use dots for electrons from one atom and crosses for electrons from the other atom.
• Only draw outer shell electrons unless the question explicitly asks for all shells.
• Check your finished diagram: count all electrons in each atom's overlapping circle to ensure every atom has achieved a complete shell (2 for \(\text{H}\), 8 for \(\text{C}, \text{N}, \text{O}, \text{Cl}\)).
Key Takeaway for Covalent Bonding: Non-metals share pairs of electrons. A single bond is \(1\) shared pair, a double bond is \(2\) shared pairs, and a triple bond is \(3\) shared pairs.
4. Metallic Bonding
The Official CCEA Definition of a Metallic Bond
A metallic bond is the electrostatic attraction between a lattice of positive metal ions (cations) and a "sea" of delocalised electrons.
What Does "Delocalised" Mean?
In a metal structure, the metal atoms lose their outer shell electrons. These electrons are not stuck to any single atom; instead, they are delocalised, meaning they are free to move throughout the entire metallic crystal lattice.
Why is this important? Because these delocalised electrons are free to move and carry electrical charge, metals are excellent conductors of electricity and heat!
Key Takeaway for Metallic Bonding: Positive ions in regular layers surrounded by a mobile sea of delocalised electrons.
5. Examiner Pitfalls & Common Mistakes to Avoid
CCEA examiners highlight several recurring errors every year. Watch out for these:
• Vague Definitions:
Incorrect: "An ionic bond is a metal and a non-metal joined together."
Correct: "An ionic bond is the electrostatic attraction between oppositely charged ions."
Incorrect: "A covalent bond is atoms sharing electrons."
Correct: "A covalent bond is the electrostatic attraction between a shared pair of electrons and the nuclei."
• Missing Brackets and Charges on Ionic Diagrams:
Always place brackets around each ion (e.g., \([\text{Mg}]^{2+}\) and \([\text{O}]^{2-}\)) and write the charge clearly in the top right corner.
• Miscounting Shared Pairs in \(\text{O}_2\) and \(\text{N}_2\):
Remember: \(\text{O}_2\) has a double bond (\(4\) shared electrons), and \(\text{N}_2\) has a triple bond (\(6\) shared electrons). Drawing a single bond for these is a very common mistake.
• The Ultimate Covalent Misconception:
When simple covalent molecular substances melt or boil (like ice turning to water, or liquid methane evaporating), covalent bonds DO NOT break! Only the weak intermolecular forces between molecules are overcome. The strong covalent bonds inside each molecule remain completely intact.
Quick Review Summary Table
• Ionic: Metal + Non-metal | Electron Transfer | Electrostatic attraction between positive cations & negative anions.
• Covalent: Non-metal + Non-metal | Electron Sharing | Electrostatic attraction between shared pair(s) of electrons & nuclei.
• Metallic: Metals / Alloys | Delocalised Sea | Electrostatic attraction between positive metal ions & delocalised electrons.