Sub-unit 1.2: Bonding Study Notes (CCEA GCSE Chemistry 1110)

Welcome to your complete revision guide for Bonding! Everything in the universe is made of tiny atoms. Except for the noble gases, most atoms are unstable on their own because their outer electron shells are not full. To achieve stability (a full outer shell of electrons, just like a noble gas), atoms join together by forming chemical bonds.

In this chapter, we will explore the three main types of chemical bonding: Ionic, Covalent, and Metallic. Don't worry if this seems tricky at first — we will break down each concept step-by-step with clear definitions, diagrams, and examiner tips to help you secure top marks.

1. Core Concepts & Building Blocks

What is an Ion?

An ion is a charged particle formed when an atom (or group of atoms) gains or loses electrons.

• Cation: A positively charged ion formed when an atom loses electron(s). Metals typically form cations.
Memory Trick: The letter t in cation looks like a plus sign (\(+\)) for positive!

• Anion: A negatively charged ion formed when an atom gains electron(s). Non-metals form anions. When non-metal atoms become simple anions, their names change to end in -ide (for example, chlorine becomes chloride, oxygen becomes oxide, sulfur becomes sulfide).
Memory Trick: A Negative Ion = Anion.

Molecular Ions

A molecular ion is a charged particle containing more than one atom covalently bonded together.

Standard Ion Formulae and Charges to Memorise

To construct correct ionic formulas in your exam, you must know the charges of standard ions from the Periodic Table and the common molecular ions:

• Group 1 Metals: form \(+1\) ions (e.g., \(\text{Li}^+\), \(\text{Na}^+\), \(\text{K}^+\))
• Group 2 Metals: form \(+2\) ions (e.g., \(\text{Mg}^{2+}\), \(\text{Ca}^{2+}\))
• Group 3 / Aluminium: forms \(+3\) ions (\(\text{Al}^{3+}\))
• Group 6 Non-metals: form \(-2\) ions (e.g., \(\text{O}^{2-}\), \(\text{S}^{2-}\))
• Group 7 Non-metals (Halides): form \(-1\) ions (e.g., \(\text{F}^-\), \(\text{Cl}^-\), \(\text{Br}^-\), \(\text{I}^-\))

Common Molecular Ions:

• Ammonium: \(\text{NH}_4^+\)
• Hydroxide: \(\text{OH}^-\)
• Nitrate: \(\text{NO}_3^-\)
• Carbonate: \(\text{CO}_3^{2-}\)
• Sulfate: \(\text{SO}_4^{2-}\)

2. Ionic Bonding

What is an Ionic Bond?

An ionic bond is the electrostatic attraction between oppositely charged ions (\(\text{M}^+\) and \(\text{X}^-\)) formed by electron transfer from a metal atom to a non-metal atom.

Analogy: Imagine a generous friend giving away an item they do not need to someone who really wants it. The metal atom wants to lose outer electrons, and the non-metal atom wants to gain them. Once transferred, both have stable full outer shells and stick tightly together because opposite charges attract!

How Electron Transfer Works

1. The metal atom loses its outer shell electron(s) to form a positive cation with a full outer shell.
2. The non-metal atom gains these electron(s) into its outer shell to form a negative anion with a full outer shell.
3. The oppositely charged ions attract each other strongly in all directions.

Dot-and-Cross Conventions for Ionic Compounds

In CCEA examinations, follow these exact drawing rules for dot-and-cross diagrams:

• Draw the outer electron shells of the resulting ions (unless the question specifically asks for all electron shells).
• Use dots for electrons from one atom and crosses for electrons from the other.
• Enclose every ion in square brackets \([\dots]\).
• Place the charge in the top-right corner outside the bracket (e.g., \([\text{Na}]^+\), \([\text{Cl}]^-\)).

Key Required Ionic Examples:

• Sodium Chloride (\(\text{NaCl}\)):
Sodium atom (\(2,8,1\)) loses \(1\) electron \(\rightarrow\) Sodium ion \([\text{Na}]^+\) (\(2,8\)).
Chlorine atom (\(2,8,7\)) gains \(1\) electron \(\rightarrow\) Chloride ion \([\text{Cl}]^-\) (\(2,8,8\)).
Diagram: \([\text{Na}]^+\) \([\text{Cl}]^-\)

• Magnesium Oxide (\(\text{MgO}\)):
Magnesium atom (\(2,8,2\)) loses \(2\) electrons \(\rightarrow\) \([\text{Mg}]^{2+}\) (\(2,8\)).
Oxygen atom (\(2,6\)) gains \(2\) electrons \(\rightarrow\) \([\text{O}]^{2-}\) (\(2,8\)).
Diagram: \([\text{Mg}]^{2+}\) \([\text{O}]^{2-}\)

• Magnesium Chloride (\(\text{MgCl}_2\)):
One magnesium atom loses \(2\) electrons. Two chlorine atoms each gain \(1\) electron.
Diagram: One \([\text{Mg}]^{2+}\) ion and two \([\text{Cl}]^-\) ions.

• Sodium Oxide (\(\text{Na}_2\text{O}\)):
Two sodium atoms each lose \(1\) electron. One oxygen atom gains \(2\) electrons.
Diagram: Two \([\text{Na}]^+\) ions and one \([\text{O}]^{2-}\) ion.

Electrical Conductivity of Ionic Substances

• Solid State: Does not conduct electricity because the ions are held in fixed positions within the lattice and cannot move.
• Molten (Liquid) or Aqueous Solution (Dissolved in Water): Does conduct electricity because the ionic lattice breaks down, allowing ions to become free to move and carry the charge.

Key Takeaway: Ionic bonds occur between metals and non-metals via electron transfer. Always draw resulting ions inside square brackets with charges on the top right!

3. Covalent Bonding

What is a Covalent Bond?

A covalent bond is the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded non-metal atoms.

A molecule is defined as two or more atoms covalently bonded together.

Analogy: Unlike ionic bonding where electrons are given away, covalent bonding is like two people holding hands across a shared space. Both non-metal atoms need electrons to fill their outer shells, so they share pairs between them.

Single, Double, and Triple Covalent Bonds

• Single Covalent Bond: One shared pair of electrons (\(2\) electrons total, e.g., \(\text{H}-\text{H}\)).
• Double Covalent Bond: Two shared pairs of electrons (\(4\) electrons total, e.g., \(\text{O}=\text{O}\)).
• Triple Covalent Bond: Three shared pairs of electrons (\(6\) electrons total, e.g., \(\text{N}\equiv\text{N}\)).

Diatomic Molecules

A diatomic molecule contains two atoms chemically bonded together. You must know these seven diatomic elements:
\(\text{H}_2\), \(\text{N}_2\), \(\text{O}_2\), \(\text{F}_2\), \(\text{Cl}_2\), \(\text{Br}_2\), \(\text{I}_2\)

Dot-and-Cross Diagrams for Covalent Molecules

When drawing covalent molecules, intersecting circles represent the overlapping outer electron shells. The shared electrons must be placed inside the overlapping area, using dots for one atom and crosses for the other.

Key Required Covalent Molecules:

• Hydrogen (\(\text{H}_2\)): Two hydrogen atoms share \(1\) pair of electrons (single bond). Each atom has \(2\) electrons in its outer shell.
• Chlorine (\(\text{Cl}_2\)): Two chlorine atoms share \(1\) pair of electrons (single bond). Each atom has an octet of \(8\) outer electrons.
• Hydrogen Chloride (\(\text{HCl}\)): Hydrogen shares \(1\) pair of electrons with chlorine.
• Water (\(\text{H}_2\text{O}\)): Central oxygen atom shares \(1\) pair of electrons with each of the two hydrogen atoms (two single bonds; oxygen also has two unshared lone pairs).
• Ammonia (\(\text{NH}_3\)): Central nitrogen atom shares \(1\) pair of electrons with each of three hydrogen atoms (three single bonds; nitrogen has one unshared lone pair).
• Methane (\(\text{CH}_4\)): Central carbon atom shares \(1\) pair of electrons with each of four hydrogen atoms (four single bonds).
• Oxygen (\(\text{O}_2\)): Two oxygen atoms share \(2\) pairs of electrons (double bond).
• Nitrogen (\(\text{N}_2\)): Two nitrogen atoms share \(3\) pairs of electrons (triple bond).

Key Takeaway: Covalent bonding occurs between non-metal atoms through shared electron pairs. Single = 1 pair, Double = 2 pairs, Triple = 3 pairs.

4. Metallic Bonding

What is a Metallic Bond?

A metallic bond is the electrostatic attraction between positive metal ions in a regular lattice and the "sea" of delocalised electrons.

What are Delocalised Electrons?

Delocalised electrons are free outer-shell electrons that are not bound to any individual atom or ion and are free to move throughout the structure.

The 2D Metallic Model

In examination diagrams, metallic bonding is represented in 2D as:
• A regular array (neat rows and layers) of identical circles with positive signs (\(+\)) representing metal cations.
• Randomly dispersed negative signs (\(e^-\)) between the ions representing the sea of delocalised electrons.

Electrical Conductivity of Metals

Metals are excellent conductors of electricity in both solid and liquid states because their delocalised electrons are free to move throughout the structure and carry electrical charge.

Key Takeaway: Metallic bonding consists of positive metal ions in neat rows surrounded by a sea of free-moving delocalised electrons.

5. Summary Table of the Three Bonding Types

• Ionic Bonding:
- Between: Metal and Non-metal.
- Mechanism: Transfer of electron(s).
- Attraction: Strong electrostatic attraction between oppositely charged ions (\(\text{M}^+\) and \(\text{X}^-\)).
- Conductivity: Only when molten or in aqueous solution (due to free-moving ions).

• Covalent Bonding:
- Between: Non-metal and Non-metal.
- Mechanism: Sharing of electron pairs.
- Attraction: Electrostatic attraction between shared pair of electrons and bonded nuclei.
- Conductivity: Generally do not conduct electricity.

• Metallic Bonding:
- Between: Metal atoms.
- Mechanism: Outer electrons become delocalised.
- Attraction: Electrostatic attraction between positive metal ions and delocalised electrons.
- Conductivity: Excellent conductors in both solid and liquid states (due to free-moving delocalised electrons).

6. Common Pitfalls & Examiner Warnings

Avoid these frequent exam mistakes highlighted in CCEA Chief Examiner reports:

1. Boiling Covalent Substances vs Breaking Bonds:
Wrong: "When water boils, covalent bonds break."
Correct: When simple covalent substances melt or boil, only the weak intermolecular forces between molecules are overcome. The strong covalent bonds inside the molecules remain completely intact!

2. Describing Particles in Metallic Bonding:
Wrong: "Metals consist of neutral metal atoms surrounded by electrons."
Correct: Metals consist of positive metal ions (cations) in a regular lattice surrounded by delocalised electrons.

3. Dot-and-Cross Formatting for Ions:
Always put square brackets around ions and write the charge in the top-right corner outside the bracket (e.g., \([\text{Cl}]^-\)). Omitting brackets or putting charges inside the brackets loses marks.

4. Explaining Electrical Conductivity:
• If explaining why molten ionic compounds conduct: state that ions are free to move and carry charge (never say electrons!).
• If explaining why metals conduct: state that delocalised electrons are free to move and carry charge (never say ions!).
• Remember: Solid ionic compounds do not conduct because ions are fixed in the lattice.