Understanding Chemical Bonding

Chemical bonding is all about how atoms connect to form the substances around you. Whether it's the salt on your chips or the metal in your phone, everything depends on how atoms bond together.

There are three main types of bonding you need to master: ionic bonding (where electrons are transferred), covalent bonding (where electrons are shared), and metallic bonding (where electrons form a "sea"). Each type creates substances with completely different properties.

Understanding bonding helps explain why diamond is so hard, why metals conduct electricity, and why salt dissolves in water - it's the foundation of chemistry!

Remember: The type of bonding determines the properties of the substance.

Ionic Bonding Basics

Ionic bonding happens when electrons are donated from one atom to another - there's no sharing involved! This typically occurs between metals (which lose electrons) and non-metals (which gain electrons).

The magic happens through electrostatic attraction between oppositely charged ions. For example, in MgO, magnesium loses two electrons to become Mg²⁺, whilst oxygen gains two electrons to become O²⁻. These opposite charges attract strongly.

You'll need to memorise key ions like sulfate (SO₄²⁻), nitrate (NO₃⁻), carbonate (CO₃²⁻), phosphate (PO₄³⁻), and ammonium (NH₄⁺). These pop up everywhere in chemistry!

The ions arrange themselves in a giant ionic lattice - a repeating 3D structure held together by those strong electrostatic forces.

Top Tip: Use dots and crosses diagrams to show where electrons come from in ionic compounds.

Properties of Ionic Compounds

Ionic compounds behave in predictable ways because of their structure. Electrical conductivity only happens when ions can move freely - so ionic compounds conduct when molten or dissolved, but not when solid.

Melting points are typically high because those electrostatic forces in the lattice are incredibly strong. It takes loads of energy to break apart the giant ionic structure.

Solubility in water works because water molecules are polar. They surround the ions and literally pull them away from the lattice, causing it to dissolve. This is why salt dissolves so easily in your tea!

The key is remembering that ionic properties all come back to those strong electrostatic forces holding the lattice together.

Exam Hint: Always explain ionic properties by referring to the electrostatic forces and whether ions can move freely.

Covalent Bonding Fundamentals

Covalent bonding is all about sharing electrons so both atoms can achieve a full outer shell. Unlike ionic bonding, both positive nuclei are attracted to the same shared electrons, creating a strong bond.

Simple molecules like HCl involve sharing one pair of electrons. More complex molecules like propanol (C₃H₇OH) can be drawn as structural diagrams showing all the bonds between atoms.

Double bonds involve sharing two pairs of electrons, like in CO₂. You can represent these molecules using structural formulas that show every single bond as a line.

Drawing these structures becomes second nature with practice - start with the central atom and work outwards, making sure every atom has the right number of bonds.

Study Smart: Practice drawing structural formulas regularly - they're essential for organic chemistry later!

Dative Covalent Bonds

Dative covalent bonds (also called coordinate bonds) form when one atom donates both electrons to the shared pair. This happens when an atom has a lone pair of electrons to spare.

The classic example is the ammonium ion (NH₄⁺). Nitrogen donates its lone pair to a hydrogen ion (H⁺), which is just a single proton. The resulting ion has a positive charge because it's gained an extra proton.

You can spot coordinate bonds in diagrams - they're sometimes shown with an arrow pointing from the electron donor to the acceptor. Once formed, coordinate bonds behave exactly like normal covalent bonds.

This concept explains how many complex ions and molecules form, especially in transition metal chemistry.

Key Point: In dative bonding, one atom provides both electrons, but the bond strength is identical to regular covalent bonds.

Giant Covalent Structures

Giant covalent structures are massive networks of atoms held together by covalent bonds. Group 4 elements like carbon form these structures, creating materials with extraordinary properties.

Diamond consists of millions of carbon atoms, each bonded to four others in a tetrahedral shape. This makes it incredibly hard due to the strong covalent bonds throughout. However, it doesn't conduct electricity because there are no free electrons.

Graphite has a different structure - each carbon bonds to only three others, forming hexagonal layers. The layers are held together by weak intermolecular forces, making graphite soft and slippery. It conducts electricity because each carbon has one spare electron that can move freely.

These allotropes of carbon show how bonding arrangement determines properties completely.

Real World: Diamond's hardness makes it perfect for cutting tools, while graphite's conductivity and softness make it ideal for pencils and lubricants.

Graphene: The Wonder Material

Graphene is essentially one single layer of graphite - just one atom thick! It's the thinnest material known to science, yet it's 100 times stronger than steel.

The structure consists of carbon atoms arranged in hexagonal patterns, with weak van der Waals forces holding the layers together in graphite. When you isolate just one layer, you get graphene.

This wonder material is conductive, transparent, and incredibly flexible. It's revolutionising technology in aircraft, cars, and mobile phones because of this unique combination of properties.

Graphene perfectly demonstrates how atomic structure determines macroscopic properties - understanding bonding literally helps explain cutting-edge technology!

Amazing Fact: Graphene is so thin that it's technically two-dimensional, yet stronger than any bulk material!

Metallic Bonding Structure

Metallic bonding creates a completely different structure where metal atoms pack closely together. Each atom donates its outer electrons to form a "sea of delocalised electrons" that surrounds all the metal ions.

The bonding involves electrostatic forces of attraction between the positive metal ions and the negative electron sea. This creates another type of giant lattice structure.

The strength of metallic bonding depends on the charge of the metal ions. Higher charged ions (like Al³⁺) create stronger bonding than lower charged ones (like Na⁺), which explains why aluminium has a much higher melting point than sodium.

You can see this pattern clearly when comparing melting points across the periodic table - more electrons in the sea means stronger bonding.

Pattern Spotting: As you go down Group 1, ionic radius increases but melting point decreases - this shows weaker metallic bonding.

Properties of Metals

The sea of delocalised electrons explains all metallic properties brilliantly. Metals are malleable (can be shaped) and ductile (drawn into wires) because layers of atoms can slide over each other without breaking bonds.

Electrical conductivity happens because electrons are free to move through the metal structure. Unlike ionic compounds, metals conduct in the solid state because the electron sea is always mobile.

Metals are generally insoluble in water because metallic bonds are too strong to be disrupted by polar water molecules. The bonding is fundamentally different from ionic compounds.

These properties make metals incredibly useful for construction, wiring, and manufacturing - it's all down to that electron sea structure!

Memory Trick: Think of metallic bonding like a crowd of people (atoms) in a swimming pool (electron sea) - they can move around but stay connected.

Molecular Shapes and Electron Repulsion

Electron pair repulsion theory explains why molecules have specific shapes. Charge clouds (regions where electrons are likely to be found) repel each other and arrange themselves as far apart as possible around the central atom.

The key rule is that lone pairs repel more than bonding pairs. The order of repulsion strength is: Lp:Lp > Lp:Bp > Bp:Bp $where Lp = lone pair, Bp = bonding pair$.

BeCl₂ has 2 bonding pairs and 0 lone pairs, creating a linear shape with 180° bond angles. Water (H₂O) has 2 bonding pairs and 2 lone pairs, creating a bent shape with 104.5° bond angles (less than the expected 109.5° due to lone pair repulsion).

Understanding molecular shapes is crucial for predicting properties like polarity and reactivity.

Exam Success: Always count electron pairs around the central atom first, then consider how lone pairs affect the final shape.