Chemical Bonding
Discover how atoms join together to form compounds through ionic, covalent and metallic bonds — and how electronegativity determines bond type.
Why Do Atoms Bond?
Atoms bond with each other to achieve a more stable electron arrangement. Most atoms are stable when their outermost shell is full — this is called the octet rule (8 electrons in the outer shell, except for hydrogen and helium which need 2). Atoms gain, lose or share electrons to reach this stable configuration.
Valence Electrons are the Key
Only valence electrons (electrons in the outermost shell) participate in bonding. The number of valence electrons an atom has determines how it bonds and with how many other atoms.
Valence Electrons of Common Elements
H
1 ve
C
4 ve
N
5 ve
O
6 ve
F
7 ve
Na
1 ve
Cl
7 ve
Ne
8 ve
ve = valence electrons. Noble gases (like Ne) already have 8 — they rarely bond.
Ionic Bonding
Ionic bonds form when one atom transfers one or more electrons to another atom. This produces oppositely charged ions that attract each other. Ionic bonding typically occurs between a metal and a non-metal.
Cations (positive ions)
Metals lose electrons to form positively charged cations. For example, sodium (Na) loses 1 electron to become Na+.
Anions (negative ions)
Non-metals gain electrons to form negatively charged anions. For example, chlorine (Cl) gains 1 electron to become Cl−.
Formation of Sodium Chloride (NaCl)
Na donates its single valence electron to Cl, forming Na+ and Cl−. The opposite charges attract, forming an ionic bond.
Properties of ionic compounds: high melting points, conduct electricity when dissolved in water or molten, form crystal lattice structures, and are generally hard and brittle.
Covalent Bonding
Covalent bonds form when two atoms share pairs of electrons. This occurs mainly between non-metal atoms. Each shared pair counts as one bond; atoms can form single, double, or triple bonds.
Single Bond
H—H
1 shared pair
(H2)
Double Bond
O=O
2 shared pairs
(O2)
Triple Bond
N≡N
3 shared pairs
(N2)
Properties of covalent compounds:
- Generally low melting and boiling points (small molecules)
- Do not conduct electricity (no free ions)
- Can be gases, liquids or solids at room temperature
- Examples: water (H2O), carbon dioxide (CO2), methane (CH4)
Metallic Bonding & Electronegativity
Metallic Bonding
In metals, valence electrons are delocalised — they move freely through a lattice of positive metal ions. This “sea of electrons” holds the metal together and explains why metals:
- Conduct electricity and heat
- Are malleable (can be shaped) and ductile (drawn into wire)
- Have high melting points
- Have a shiny appearance (lustre)
Electronegativity
Electronegativity is a measure of an atom’s ability to attract shared electrons towards itself. It increases across a period (left to right) and decreases down a group.
Bond type based on electronegativity difference (ΔEN):
- ΔEN < 0.4: Non-polar covalent
- 0.4 ≤ ΔEN < 1.7: Polar covalent
- ΔEN ≥ 1.7: Ionic
| Property | Ionic | Covalent | Metallic |
|---|---|---|---|
| Electron behaviour | Transferred | Shared | Delocalised |
| Conductivity | When molten/dissolved | Generally none | Yes (solid) |
| Melting point | High | Low–moderate | High |
| Example | NaCl, MgO | H2O, CO2 | Fe, Cu, Al |
Key Vocabulary
| Term | Definition |
|---|---|
| Ionic bond | A chemical bond formed by the transfer of electrons from a metal to a non-metal, producing oppositely charged ions that attract. |
| Covalent bond | A chemical bond formed by the sharing of one or more pairs of electrons between two non-metal atoms. |
| Electronegativity | A measure of the ability of an atom to attract a shared pair of electrons towards itself in a chemical bond. |
| Delocalised electrons | Electrons that are not associated with a single atom but are free to move throughout a metallic lattice structure. |
Worked Examples
Predicting the bond type formed between magnesium (Mg) and oxygen (O).
Step 1: Identify the elements. Mg is a metal (Group 2); O is a non-metal (Group 16).
Step 2: Metal + non-metal → ionic bond.
Step 3: Mg has 2 valence electrons; it loses both to become Mg2+. O has 6 valence electrons; it gains 2 to become O2−.
Answer: MgO is an ionic compound. The oppositely charged ions Mg2+ and O2− attract to form the ionic lattice.
Drawing the covalent structure of water (H2O).
Step 1: Count valence electrons. O has 6; each H has 1. Total = 6 + 1 + 1 = 8.
Step 2: O needs 2 more electrons to fill its outer shell, so it forms 2 bonds.
Step 3: Each H shares its 1 electron with O, forming two single covalent bonds.
Answer: H—O—H. O has 2 bonding pairs and 2 lone pairs. The bond angle is approximately 104.5° because of lone pair repulsion.
Using electronegativity to classify the HCl bond.
Given: H has electronegativity 2.1; Cl has electronegativity 3.0.
Step 1: Calculate the difference: ΔEN = 3.0 − 2.1 = 0.9
Step 2: 0.4 ≤ 0.9 < 1.7 → polar covalent bond.
Answer: The H—Cl bond is a polar covalent bond. The electrons are pulled more strongly towards Cl, giving Cl a slight negative charge (δ−) and H a slight positive charge (δ+).
Knowledge Check
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Key Concepts Summary
- •Atoms bond to achieve a full outer electron shell (octet rule), gaining stability.
- •Ionic bonds: electrons are transferred from a metal to a non-metal, forming positive and negative ions.
- •Covalent bonds: electrons are shared between non-metal atoms; can be single, double, or triple bonds.
- •Metallic bonds: delocalised electrons move freely through a lattice of metal cations, explaining conductivity and malleability.
- •Electronegativity difference determines bond type: <0.4 non-polar covalent; 0.4–1.7 polar covalent; >1.7 ionic.