Ionic Bonding
2. Ionic Bonding
Learning outcomes
- I can explain how ionic bonds form between oppositely charged ions.
- I can describe the electrostatic attraction in ionic compounds.
- I can identify ionic compounds from their formulas.
- I can represent ionic bonding using diagrams.
- I can explain why ionic compounds are electrically neutral overall.
Introduction
In the previous lesson, you learned that atoms can gain or lose electrons to form ions. Positive ions (cations) and negative ions (anions) are attracted to each other because they carry opposite electrical charges.
This attraction forms one of the strongest types of chemical bonding, known as an ionic bond.
Ionic bonding is responsible for many familiar substances, including table salt (NaCl), magnesium oxide (MgO), and calcium chloride (CaCl₂).
What is an Ionic Bond?
An ionic bond is the strong electrostatic attraction between oppositely charged ions.
The bond forms after one atom transfers one or more electrons to another atom.
This creates:
- A positive ion (cation)
- A negative ion (anion)
These oppositely charged ions attract one another, holding the compound together.
Unlike covalent bonding, electrons are transferred, not shared.
How Does Ionic Bonding Form?
Ionic bonding usually occurs between:
- a metal
- a non-metal
The metal loses electrons.
The non-metal gains those electrons.
Example:
Sodium + Chlorine
Before:
- Sodium: 2,8,1
- Chlorine: 2,8,7
Electron transfer:
- Sodium loses one electron.
- Chlorine gains one electron.
After:
- Sodium ion: Na⁺
- Chloride ion: Cl⁻
The opposite charges attract, forming an ionic bond.
Electrostatic Attraction
The force holding ionic compounds together is called electrostatic attraction.
Opposite charges attract just like opposite poles of magnets.
Positive ions attract negative ions in all directions.
This attraction is very strong.
It is responsible for the strength and stability of ionic compounds.
Ionic Compounds
An ionic compound is a compound made from positive and negative ions.
Some common examples are:
| Compound | Formula |
|---|---|
| Sodium chloride | NaCl |
| Magnesium oxide | MgO |
| Calcium chloride | CaCl₂ |
| Potassium bromide | KBr |
| Lithium fluoride | LiF |
| Aluminium oxide | Al₂O₃ |
Notice that every ionic compound contains:
- Positive ions
- Negative ions
Why Are Ionic Compounds Neutral?
Although ions have charges, the overall compound has no net charge.
The total positive charge equals the total negative charge.
Examples:
Sodium chloride
Na⁺
Cl⁻
Total charge:
+1 + (−1) = 0
Magnesium oxide
Mg²⁺
O²⁻
Total charge:
+2 + (−2) = 0
Calcium chloride
Ca²⁺
2Cl⁻
Total charge:
+2 + (−1) + (−1) = 0
Every ionic compound must have a total charge of zero.
Representing Ionic Bonding
Chemists often use dot-and-cross diagrams to show ionic bonding.
These diagrams show:
- Valence electrons
- Electron transfer
- The charges on the ions
Example:
Before transfer:
Na • Cl ×××××××
After transfer:
[Na]⁺ [Cl]⁻
The chloride ion now has a complete outer shell.
The sodium ion also has a full outer shell beneath it.
Dot-and-cross diagrams help us see exactly how ionic bonds form.
Ionic Lattices
Individual ionic compounds do not exist as separate molecules.
Instead, millions of ions arrange themselves into a giant ionic lattice.
Each positive ion is surrounded by negative ions.
Each negative ion is surrounded by positive ions.
This arrangement:
- maximizes attraction
- minimizes repulsion
- makes ionic compounds very stable
Recognizing Ionic Compounds from Their Formulas
Many ionic compounds can be recognized because they contain:
- A metal first
- A non-metal second
Examples:
| Formula | Ionic? |
|---|---|
| NaCl | Yes |
| MgO | Yes |
| CaBr₂ | Yes |
| LiF | Yes |
| CO₂ | No (covalent) |
| H₂O | No (covalent) |
| NH₃ | No (covalent) |
A quick clue is that metal + non-metal usually indicates an ionic compound.
Real-World Applications
Ionic compounds are used in many everyday products and industries.
Examples include:
- Table salt for food
- Road salt for melting ice
- Fertilizers
- Medicines
- Ceramics
- Glass production
- Electrolyte solutions
- Water treatment
Worked Examples
Example 1
What type of bond forms between Na⁺ and Cl⁻?
Answer:
An ionic bond.
Example 2
Why are Na⁺ and Cl⁻ attracted to one another?
Answer:
Because they have opposite electrical charges.
Example 3
Is MgO electrically neutral?
Mg²⁺
O²⁻
Total charge:
+2 + (−2) = 0
Answer:
Yes.
Example 4
Which compound is ionic?
A. CO₂
B. H₂O
C. NaCl
D. CH₄
Answer:
NaCl
Example 5
Why is CaCl₂ written with two chlorine ions?
Answer:
One calcium ion has a charge of +2.
Each chloride ion has a charge of −1.
Two chloride ions are needed to balance the +2 charge and produce an overall neutral compound.
Did You Know?
A single grain of table salt contains an enormous number of ions arranged in a repeating crystal lattice. Instead of existing as separate NaCl molecules, the sodium and chloride ions form one continuous three-dimensional structure that extends throughout the entire crystal.
Key Terms
| Term | Definition |
|---|---|
| Ionic Bond | The strong electrostatic attraction between oppositely charged ions. |
| Electrostatic Attraction | The force of attraction between positive and negative charges. |
| Ionic Compound | A compound made of positive and negative ions held together by ionic bonds. |
| Cation | A positively charged ion formed by losing electrons. |
| Anion | A negatively charged ion formed by gaining electrons. |
| Electron Transfer | The movement of one or more electrons from one atom to another. |
| Ionic Lattice | A giant three-dimensional arrangement of alternating positive and negative ions. |
Key Takeaways
- Ionic bonds form when electrons are transferred from a metal to a non-metal.
- Positive and negative ions are held together by strong electrostatic attraction.
- Ionic compounds always contain both cations and anions.
- The total positive and negative charges balance, making ionic compounds electrically neutral overall.
- Dot-and-cross diagrams help illustrate electron transfer and ion formation.
- Ionic compounds exist as giant crystal lattices rather than individual molecules.