Viridian Kids School

Van der Waals forces are weak intermolecular forces that arise from attractive or repulsive interactions between molecules or atoms. They are not true chemical bonds (like covalent or ionic bonds) but rather physical interactions that play a big role in many natural and biological processes—such as why geckos stick to walls or how gases condense into liquids.
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🧲 Types of Van der Waals Forces

1. London Dispersion Forces (Instantaneous Dipole–Induced Dipole)

Occurs in: All molecules (even nonpolar ones like helium or oxygen gas).

How it happens:

At any instant, electrons in an atom or molecule may be unevenly distributed.

This creates a temporary dipole.

That dipole can induce a dipole in a neighboring atom or molecule.

The two temporary dipoles attract each other weakly.

Strength: Weakest type, but increases with:

The size (number of electrons) of the molecule

The surface area (more contact = stronger attraction)

💡 Example: Noble gases like argon can liquefy at low temperatures due to these forces.

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2. Dipole–Dipole Forces

Occurs in: Polar molecules (those with permanent dipoles).

How it happens:

One molecule’s positive end attracts another molecule’s negative end.

Strength: Stronger than London dispersion, but weaker than hydrogen bonds.

💡 Example: Between molecules of HCl — the positive hydrogen end of one HCl attracts the negative chlorine end of another.

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3. Dipole–Induced Dipole Forces

Occurs in: Between a polar and a nonpolar molecule.

How it happens:

The electric field of a polar molecule distorts the electron cloud of a nearby nonpolar molecule.

This creates an induced dipole in the nonpolar molecule → attraction forms.

💡 Example: Between water (polar) and oxygen gas (nonpolar).

Mazi Uchechukwu

Chemical Bonding Details

1. Definition
Chemical bonding refers to the attractive forces that hold atoms or ions together in compounds. These bonds form because atoms seek stability, often achieved by attaining a full outer electron shell.

2. Types of Chemical Bonds

Ionic Bond:
Formed when one atom transfers electrons to another, creating oppositely charged ions that attract each other.
Example: Sodium chloride (NaCl) — sodium donates an electron to chlorine.
Covalent Bond:
Formed when two atoms share one or more pairs of electrons.
Example: Water (H₂O) — oxygen shares electrons with two hydrogen atoms.
Single bond: One shared pair of electrons (e.g., H–H)
Double bond: Two shared pairs (e.g., O=O)
Triple bond: Three shared pairs (e.g., N≡N)
Metallic Bond:
Found in metals, where atoms share a “sea” of delocalized electrons that move freely, giving metals conductivity and malleability.
3. Bond Polarity
When atoms with different electronegativities form a covalent bond, electrons are shared unequally, creating a polar bond.
Example: In H₂O, oxygen attracts electrons more strongly than hydrogen, making the molecule polar.

4. Intermolecular Forces
These are weaker forces between molecules, not within them:

Hydrogen bonds (strongest intermolecular force)
Dipole–dipole interactions
London dispersion forces (weakest, present in all molecules)
5. Bond Strength and Length

Stronger bonds (like triple bonds) are shorter and require more energy to break.
Weaker bonds (like single bonds) are longer and easier to break.
6. Importance of Chemical Bonding
Chemical bonding determines the structure, stability, and properties of substances — from the hardness of diamonds to the solubility of salts and the behavior of biological molecules.

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