Chemistry (0620) | Ions and ionic bonds

Welcome to Section 2.4: Ions and Ionic Bonds!

Hello future chemists! This is one of the most fundamental topics in IGCSE Chemistry. Why are we studying ions and ionic bonds? Because these powerful bonds create everyday substances like salt (sodium chloride). Understanding how they form is key to predicting the behavior of many compounds!

Don't worry if this seems tricky at first. It's all about atoms trying to become stable, just like you trying to find the most comfortable position to relax!

1. Achieving Stability: The Formation of Ions (Core 2.4.1)

All atoms want to achieve the stable electron structure of a Noble Gas (Group VIII), meaning they want a full outer electron shell (usually 8 electrons, known as the octet rule).

How Atoms Become Ions

Atoms achieve stability by either losing or gaining electrons. When an atom loses or gains electrons, it is no longer neutral; it becomes an ion, which has an overall electrical charge.

A. Positive Ions: Cations (Metals)

Metals (found on the left side of the Periodic Table, Groups I, II, III) usually have 1, 2, or 3 electrons in their outer shell. It is easier for them to lose these few electrons to achieve a full inner shell.

• When a neutral atom loses electrons, it becomes a positive ion, known as a Cation.
• Why positive? Because it now has more positive protons in the nucleus than negative electrons outside it.
• Example: Sodium (Na) is in Group I (Electronic structure 2,8,1). It loses 1 electron to become stable (2,8). The ion formed is \(Na^+\).

Quick Memory Aid: Cations are posiTive, because the 'T' looks like a plus sign (+).

B. Negative Ions: Anions (Non-Metals)

Non-metals (found on the right side of the Periodic Table, Groups V, VI, VII) usually have 5, 6, or 7 electrons in their outer shell. It is easier for them to gain a few electrons to complete their shell.

• When a neutral atom gains electrons, it becomes a negative ion, known as an Anion.
• Why negative? Because it now has more negative electrons than positive protons.
• Example: Chlorine (Cl) is in Group VII (Electronic structure 2,8,7). It gains 1 electron to become stable (2,8,8). The ion formed is \(Cl^-\) (or \(Cl^{1-}\)).

Key Takeaway (Section 1): Ions form when atoms lose (metals become cations) or gain (non-metals become anions) electrons to achieve a stable Noble Gas configuration.

2. The Ionic Bond Defined (Core 2.4.2)

Once metals have given away electrons (becoming positive cations) and non-metals have accepted those electrons (becoming negative anions), the resulting ions have opposite charges.

The ionic bond is simply the strong electrostatic attraction between these oppositely charged ions.

Think of it like powerful magnets snapping together! This attraction is what holds the entire compound together.

Ionic compounds form between atoms from metallic elements (low electronegativity, usually Group I, II) and non-metallic elements (high electronegativity, usually Group VI, VII).

3. Visualising Ionic Bond Formation: Dot-and-Cross Diagrams (Core 2.4.3 & Supplement 2.4.6)

Dot-and-cross diagrams are a required way to show how electrons are transferred and how the final ions look.

Step-by-Step Example: Sodium Chloride (NaCl) (Group I and Group VII)

1. Start Atoms: Sodium (Na) has 1 outer electron (use a dot •). Chlorine (Cl) has 7 outer electrons (use a cross x).
2. Transfer: The sodium atom gives its single outer electron to the chlorine atom.
3. Resulting Ions:
   • The Na atom loses its outer shell, revealing a full inner shell. It becomes a \(Na^+\) ion.
   • The Cl atom fills its outer shell, taking the electron from Na (now shown as a dot). It becomes a \(Cl^-\) ion.

Crucially, when drawing the resulting ions:

• Show the full outer shells of both ions (8 electrons, unless it's a small ion like \(Li^+\)).
• Use square brackets around the ion structure.
• Write the charge outside the brackets (\(Na^+\) and \(Cl^-\)).

Extended Example: Magnesium Chloride (\(MgCl_2\)) (Metallic and Non-Metallic)

Magnesium (Group II, 2 outer electrons) needs to lose 2 electrons. Chlorine (Group VII, 7 outer electrons) needs to gain 1 electron.

Since Mg needs to lose two electrons, it must bond with two chlorine atoms, one electron going to each chlorine atom.

The final ionic compound is made of: one \(Mg^{2+}\) ion and two separate \(Cl^-\) ions. The overall charge of the compound is zero (\(+2 + (-1) + (-1) = 0\)).

4. Structure of Ionic Compounds: The Giant Ionic Lattice (Supplement 2.4.5)

In reality, ionic compounds don't exist as just one pair of ions (like \(Na^+Cl^-\)). Instead, billions of ions are packed together in a highly organised 3D structure.

The Giant Lattice Structure

• Ionic compounds form a Giant Ionic Lattice structure.
• This is a regular arrangement of alternating positive ions (cations) and negative ions (anions).
• Because the electrostatic attraction is strong and works in all directions, the ions are locked rigidly in place throughout the entire crystal.

Analogy: Imagine building a wall out of alternating LEGO bricks of two different colours. It's a massive, repeating, highly structured pattern.

Key Takeaway (Section 4): Ionic compounds form a massive, rigid structure where every positive ion is surrounded by negative ions, and vice versa. This structure is held together by incredibly strong forces.

5. Properties of Ionic Compounds (Core 2.4.4 & Supplement 2.4.7)

The giant lattice structure and the strong electrostatic forces explain all the key physical properties of ionic compounds.

A. High Melting Points and Boiling Points

Core Requirement (State): Ionic compounds have high melting points and boiling points.

Supplement Requirement (Explain):

• To melt or boil an ionic compound, you must provide enough energy to overcome the very strong electrostatic forces holding the ions together in the giant lattice.
• Since these forces are so strong, a large amount of thermal energy is needed.
• Therefore, melting points and boiling points are high.

B. Electrical Conductivity

When Solid (Poor Conductor)

Core Requirement (State): Ionic compounds are poor electrical conductors when solid.

Supplement Requirement (Explain):

• In the solid state, although ions are charged, they are held in fixed positions within the rigid lattice.
• There are no mobile charged particles (ions or electrons) that are free to move and carry the electric current.

When Molten or Aqueous (Good Conductor)

Core Requirement (State): Ionic compounds are good electrical conductors when aqueous (dissolved in water) or molten (melted).

Supplement Requirement (Explain):

• When melted or dissolved, the strong forces of attraction are overcome.
• The ions are now free to move (they become mobile ions).
• These mobile charged ions can travel towards the electrodes, thus carrying the electric current.

C. Solubility (Did you know?)

Many (though not all!) ionic compounds are soluble in water. This is because water molecules are polar and can break down the ionic lattice, pulling the individual ions away from the structure and allowing them to move freely in the solution.

Key Takeaway (Section 5): The strong forces in the giant lattice lead to high MP/BP. Conductivity depends entirely on the freedom of the charged ions: fixed = non-conductor; mobile (molten/aqueous) = conductor.