Chemical bonding — ionic, covalent, metallic (KCSE Chemistry Form 2)

TL;DR

Chemical bonds are forces that hold atoms together to form compounds and molecules. Atoms bond to achieve a stable electron configuration, usually a full outer shell. The three main types are ionic (electron transfer), covalent (electron sharing), and metallic (delocalized electrons).

1. The Mental Model

Think of atoms as wanting to be "happy" and stable, like noble gases. They achieve this by gaining, losing, or sharing electrons in their outermost shell. This interaction creates a strong "glue" that holds them together.

2. The Core Material

Why Atoms Bond

Atoms bond to achieve a stable electron configuration, typically a full outermost electron shell (like the noble gases). This is often referred to as the octet rule (eight electrons in the outer shell), though for very small atoms like hydrogen, it's the duplet rule (two electrons). Achieving this stability lowers the atom's energy.

Types of Chemical Bonds

1. Ionic Bonding

Ionic bonding happens when there's a complete transfer of electrons from one atom to another. This usually occurs between a metal and a non-metal.

• How it works:
  • Metals tend to lose their valence electrons to achieve a stable electron configuration, forming positively charged ions called cations.
  • Non-metals tend to gain electrons to achieve a stable electron configuration, forming negatively charged ions called anions.
  • The strong electrostatic force of attraction between the oppositely charged ions forms the ionic bond.
• Properties of Ionic Compounds:
  • High melting and boiling points (strong forces of attraction).
  • Usually soluble in water.
  • Conduct electricity when molten or dissolved in water (ions are free to move), but not when solid (ions are fixed in a lattice).
  • Form crystal lattices (regular, repeating arrangements of ions).

2. Covalent Bonding

Covalent bonding occurs when atoms share electrons to achieve a stable electron configuration. This typically happens between two non-metals.

• How it works:
  • Each atom contributes one or more electrons to form a shared pair of electrons.
  • This shared pair is attracted to the nuclei of both atoms, holding them together.
  • A single bond involves one shared pair, a double bond involves two shared pairs, and a triple bond involves three shared pairs.
• Types of Covalent Bonds:
  • Polar covalent bond: Electrons are shared unequally because one atom has a stronger pull (higher electronegativity). This creates slight positive and negative charges (e.g., H₂O).
  • Non-polar covalent bond: Electrons are shared equally, usually between identical atoms (e.g., O₂, Cl₂).
• Properties of Covalent Compounds:
  • Generally lower melting and boiling points than ionic compounds (weaker intermolecular forces).
  • Often insoluble in water (unless polar).
  • Do not conduct electricity (no free ions or electrons).
  • Can exist as gases, liquids, or solids at room temperature.

3. Metallic Bonding

Metallic bonding is found in metals and alloys. It involves a "sea" of delocalized electrons.

• How it works:
  • Metal atoms lose their outermost electrons, becoming positively charged metal ions (cations).
  • These valence electrons are not associated with any one atom but are free to move throughout the entire metal structure. This forms a "sea" of delocalized electrons.
  • The strong electrostatic attraction between the positive metal ions and the delocalized electron sea holds the metal together.
• Properties of Metals (due to delocalized electrons):
  • Good conductors of electricity and heat: Free-moving electrons can carry charge and thermal energy.
  • Malleable and ductile: The "sea" of electrons allows metal ions to slide past each other without breaking the bond.
  • Lustrous (shiny): Free electrons absorb and re-emit light.
  • High melting and boiling points (strong metallic bonds).

Deciding the Type of Bond

graph TD
    A[Start: Consider the elements involved] --> B{Metal and Non-metal?};
    B -- Yes --> C[Ionic Bond];
    B -- No --> D{Two Non-metals?};
    D -- Yes --> E[Covalent Bond];
    D -- No --> F{Two Metals?};
    F -- Yes --> G[Metallic Bond];
    F -- No --> H[Review elements, check for exceptions];

3. Worked Example

Problem: Describe the bonding in Sodium Chloride (NaCl) and Oxygen gas (O₂).

Solution:

1. Sodium Chloride (NaCl):
* Elements involved: Sodium (Na) is a Group 1 metal. Chlorine (Cl) is a Group 17 non-metal.
* Electron configuration:
* Na: 2.8.1 (1 valence electron)
* Cl: 2.8.7 (7 valence electrons)
* Bonding process: Sodium, being a metal, will lose its one valence electron to achieve a stable 2.8 configuration, forming a Na⁺ ion. Chlorine, being a non-metal, will gain one electron to complete its outer shell (2.8.8), forming a Cl⁻ ion.
* Type of bond: This complete transfer of electrons leads to the formation of an ionic bond due to the strong electrostatic attraction between the Na⁺ and Cl⁻ ions.

2. Oxygen gas (O₂):
* Elements involved: Oxygen (O) is a Group 16 non-metal.
* Electron configuration:
* O: 2.6 (6 valence electrons)
* Bonding process: Each oxygen atom needs 2 more electrons to achieve a stable octet (8 electrons). They achieve this by sharing two pairs of electrons between them.
* Each oxygen atom contributes 2 electrons to the shared pool.
* This forms a double covalent bond (O=O).
* Type of bond: Since electrons are shared between two non-metal atoms, it's a covalent bond. Because both atoms are identical, the sharing is equal, making it a non-polar covalent bond.

4. Key Takeaways

• Atoms bond to achieve a stable electron configuration, usually a full outer shell.
• Ionic bonds involve the complete transfer of electrons, forming ions and strong electrostatic attraction.
• Covalent bonds involve the sharing of electrons between atoms.
• Metallic bonds involve a "sea" of delocalized electrons holding positive metal ions together.
• The type of elements (metal/non-metal) largely determines the type of bond formed.
• Ionic compounds conduct electricity when molten or dissolved, but not when solid.
• Covalent compounds generally do not conduct electricity.
• Metals are excellent conductors of heat and electricity due to their delocalized electrons.

Common mistakes to avoid:
- Confusing electron transfer (ionic) with electron sharing (covalent).
- Assuming all compounds conduct electricity; only ionic compounds in solution/molten state and metals do.
- Forgetting that metallic bonding is unique to metals and their alloys.
- Not considering the number of valence electrons when determining how many bonds an atom will form.

5. Now Try It

Choose three different compounds: Magnesium Oxide (MgO), Carbon Dioxide (CO₂), and a piece of Copper (Cu). For each, identify the type of bonding present, explain why that type of bonding occurs based on the elements involved and their electron configurations, and list two characteristic properties you would expect for each substance.

What success looks like: You'll correctly identify the bond type for each, clearly explain the electron interactions (transfer, sharing, or delocalization), and accurately state two properties linked to that bond type.