Ionic Bonding

Ever wondered why salt dissolves in water but doesn't conduct electricity until it does? Ionic bonding happens when electrons jump from metal atoms to non-metal atoms, creating oppositely charged ions that attract each other like magnets.

These ionic compounds always form lattice structures - think of them as 3D networks of alternating positive and negative ions. They're always solid at room temperature because breaking apart that giant lattice structure requires loads of energy.

Here's the clever bit: ionic compounds only conduct electricity when molten or dissolved in water. In solid form, the ions are locked in place and can't move to carry charge. But melt them or dissolve them, and those charge-carrying ions become free to move around.

Quick Tip: Remember that ionic compounds are brittle - they shatter easily because when you hit them, like charges can line up and repel each other, causing the structure to break apart.

Covalent Bonding

Unlike ionic bonding where electrons transfer completely, covalent bonding is all about sharing. Non-metal atoms share pairs of electrons, creating strong attractions between the shared electrons and both atomic nuclei.

Single covalent bonds involve one shared pair of electrons. Think of methane (CH₄) - carbon shares one pair with each of four hydrogen atoms, creating four single bonds. It's like a molecular handshake between atoms.

But atoms can be greedy and share more! Multiple covalent bonds occur when atoms share two or three pairs of electrons. Oxygen gas (O₂) has a double bond $O=O$, whilst nitrogen gas (N₂) has an incredibly strong triple bond (N≡N).

Remember: The more electron pairs shared between atoms, the stronger and shorter the bond becomes.

Dative Covalent Bonding

Sometimes one atom is generous and provides both electrons in a shared pair - that's dative covalent bonding. The atom donating the pair has a lone pair of electrons, whilst the receiving atom is electron deficient.

The classic example is when ammonia (NH₃) meets a hydrogen ion (H⁺) to form NH₄⁺. Nitrogen donates its lone pair to the hydrogen, shown with an arrow (N→H). Once formed, this dative bond is identical to regular covalent bonds.

Think of it like lending someone your phone charger - once they're using it, it works just like their own charger would. The lone pair creates a concentrated negative charge region that attracts positive ions.

Average bond enthalpy measures how strong covalent bonds are. Higher values mean stronger bonds that require more energy to break.

Key Point: Dative bonds look different when forming (one atom gives both electrons) but behave identically to normal covalent bonds once established.

Electron Pair Repulsion Theory

Imagine trying to arrange magnets around a central point - they'd push apart to get as far from each other as possible. Electron pairs behave similarly, determining molecular shapes by minimising repulsion between negative charges.

The key insight is that lone pairs cause more repulsion than bonding pairs. Lone pairs are closer to the central atom and occupy more space because they're not stretched between two nuclei. This means they push other electron pairs away more strongly.

The repulsion strength order is: lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair. Each lone pair reduces bond angles by about 2.5°.

Steric number tells you the total number of atoms and lone pairs directly attached to a central atom. This number determines the basic molecular geometry before lone pairs modify the angles.

Memory Aid: Think "LP > BP" - Lone Pairs cause more repulsion than Bonding Pairs, so they dominate molecular shape.

Common Molecular Shapes

Linear molecules have two bonding pairs arranged at 180° - perfectly straight like BeF₂. Trigonal planar molecules have three bonding pairs at 120° angles, all in the same plane like BF₃.

Tetrahedral molecules are 3D with four electron pairs. Methane (CH₄) is perfectly tetrahedral with 109.5° bond angles. But add a lone pair like in ammonia (NH₃), and the angles squeeze down to 107° because that lone pair pushes the bonding pairs closer together.

Water (H₂O) has two lone pairs on oxygen, creating a bent shape with bond angles around 104.5°. The more lone pairs present, the more the bond angles get compressed from the ideal tetrahedral angle.

These shapes aren't just academic - they determine how molecules interact, fit together, and react with each other in biological systems and chemical reactions.

Exam Tip: Always count lone pairs carefully - they're invisible in molecular formulas but crucial for predicting shapes and angles.

Complex Molecular Geometries

When you have five electron pairs, you get a trigonal bipyramid - imagine three atoms around the equator with one above and one below. PF₅ is a perfect example with 120° angles around the middle and 90° angles to the top and bottom.

Octahedral molecules like SF₆ have six bonding pairs creating a symmetrical shape with all bond angles at 90°. Picture a central atom with four others around it in a square, plus one above and one below.

These complex shapes follow the same electron repulsion principles but in three dimensions. The electron pairs arrange themselves to minimise repulsion, creating these distinctive geometries.

Understanding these shapes helps predict molecular behaviour, especially in coordination chemistry and complex biological molecules where metal centres often adopt octahedral geometries.

Visual Tip: Use molecular models or draw 3D sketches - these complex shapes are much easier to understand when you can visualise them properly.

Electronegativity Fundamentals

Electronegativity measures an atom's pulling power for electrons in covalent bonds. Think of it as atomic greed - fluorine is the most electronegative element (4.0 on the Pauling scale) and desperately wants to hog electrons.

Three factors determine electronegativity: nuclear charge (more protons = stronger pull), atomic radius (smaller atoms hold electrons tighter), and shielding (inner electrons block the nuclear attraction).

Across a period, electronegativity increases because atoms get smaller whilst gaining more protons. Down a group, it decreases because atoms get bigger and inner electrons shield the outer ones from nuclear attraction.

The nuclear charge increases down groups, but this effect is overwhelmed by the increased distance and shielding from additional electron shells.

Pattern Recognition: Electronegativity increases going right and up the periodic table, with fluorine at the top right being the champion electron-hog.

Electronegativity Trends and Bond Prediction

Moving across a period, electronegativity increases because atoms shrink whilst gaining protons. The nuclear charge grows stronger, but shielding stays constant, creating a tighter grip on electrons.

Going down a group, electronegativity decreases despite increasing nuclear charge. The atoms get bigger, more inner shells create shielding, and the increased distance weakens the attraction between nucleus and bonding electrons.

You can predict bond types using electronegativity differences: equal values give pure covalent bonds, slight differences create polar covalent bonds, and large differences result in ionic bonds.

This prediction system helps you understand why some compounds conduct electricity, dissolve in water, or have particular melting points - it all comes down to the electronegativity differences between atoms.

Quick Check: If electronegativity difference is roughly >1.7, expect ionic character; <0.5 suggests pure covalent bonding.

Bond Polarity and Dipoles

Polar bonds form when atoms with different electronegativities share electrons unequally. The more electronegative atom becomes slightly negative $δ-$ whilst the other becomes slightly positive $δ+$.

Having polar bonds doesn't automatically make a polar molecule though! The molecule's shape matters crucially. Water (H₂O) is polar because its bent shape means the dipoles don't cancel out, but carbon dioxide (CO₂) is non-polar despite having polar bonds because its linear shape makes the dipoles cancel.

A permanent dipole exists when there's a persistent charge difference across a covalent bond. These dipoles can align with each other, creating weak electrostatic attractions between molecules - about 1/100th the strength of covalent bonds.

Non-polar molecules either have no polar bonds at all, or have polar bonds arranged symmetrically so the charges cancel out completely.

Shape Matters: Always consider molecular geometry - symmetrical molecules with polar bonds often end up non-polar overall.

Predicting Molecular Polarity

Linear, symmetrical molecules are typically non-polar because dipoles cancel out. However, molecules with OH, NH, or lone H atoms at the ends are usually polar due to the significant electronegativity differences.

Carbon-containing molecules are often non-polar, especially hydrocarbons, but there are important exceptions. Diatomic elements with identical atoms (like Cl₂ or O₂) are always non-polar since there's no electronegativity difference.

C-H and S-H bonds don't produce significant dipoles due to similar electronegativities. When predicting polarity, look for the dipole sum - if individual bond dipoles cancel out, the molecule is non-polar; if they don't cancel, it's polar.

The key is recognising that molecular polarity depends on both individual bond polarities and the overall molecular geometry. Symmetry is your friend for predicting non-polar molecules.

Rule of Thumb: Symmetrical molecules tend to be non-polar even with polar bonds, whilst asymmetrical molecules with polar bonds are usually polar overall.