Covalent Bonding Basics

Covalent bonds form when atoms share electrons to create strong connections. The shared electrons have opposite spins to minimise repulsion, and the bond stays strong because positive nuclei are attracted to the negative electron cloud between them.

Co-ordinate bonding is a special type where both electrons in the bond come from the same atom - always involving lone pairs. Think of how ammonia (NH₃) bonds with a hydrogen ion to form NH₄⁺.

💡 Quick Tip: If atoms get too close, their inner electrons and nuclei repel each other, which determines the bond length.

Intermolecular Forces and Electronegativity

Intermolecular forces are weak attractions between molecules (not within them like covalent bonds). There are three types: induced dipoles, permanent dipoles, and hydrogen bonding.

Induced dipoles happen when electrons "bunch up" for milliseconds as they whizz around, creating temporary positive and negative charges that attract neighbouring molecules. Larger atoms have more electrons, so stronger intermolecular forces.

Electronegativity is how strongly an atom pulls bonding electrons towards itself. Fluorine is the most electronegative element, and it increases across the periodic table due to more protons and less shielding.

💡 Remember: Polar compounds dissolve in polar solvents, non-polar dissolve in non-polar!

Polar Bonds and Hydrogen Bonding

Permanent dipoles form when sharing isn't equal - like in HCl where chlorine's 17 protons pull electrons more than hydrogen's single proton. This creates permanent positive and negative charges on different atoms.

You can predict bond types using Pauling electronegativity numbers: differences less than 0.3 mean non-polar covalent, 0.4-1.7 means polar covalent, and over 1.7 means ionic bonding.

Hydrogen bonding is the strongest intermolecular force, occurring when hydrogen bonds directly to nitrogen, oxygen, or fluorine atoms. This explains why water, ammonia, and hydrogen fluoride have unusually high boiling points.

💡 Exam Tip: Compounds with hydrogen bonds dissolve well in water because they can form hydrogen bonds with water molecules.

Giant vs Molecular Covalent Structures

Giant covalent structures like diamond and graphite have completely different properties despite both being carbon. Diamond bonds each carbon to four others in a tetrahedral structure - making it incredibly hard with a very high melting point, but it doesn't conduct electricity.

Graphite has layers where each carbon bonds to three others in hexagons. The layers slide over each other (great for pencils!), and it conducts electricity because of delocalised electrons that can move freely.

Molecular covalent structures like iodine and ice have strong bonds within molecules but weak intermolecular forces between them. This gives them relatively low melting points and poor electrical conductivity.

💡 Key Difference: Graphite conducts electricity because electrons are delocalised; iodine doesn't because all electrons are localised in bonds.

Electrical Conductivity in Covalent Compounds

The difference between graphite and iodine's electrical properties comes down to electron freedom. In graphite, each carbon bonds to only three others, leaving one outer electron delocalised and free to move and carry charge.

In iodine, both atoms are covalently bonded together with no free electrons available. All electrons are locked up in the covalent bond between the two iodine atoms.

This principle applies broadly - for electrical conduction, you need mobile charge carriers, whether they're delocalised electrons or free ions.

💡 Remember: Delocalised = free to move = conducts electricity!

Ionic Bonding Formation

Ionic bonds form between metals and non-metals through electron transfer, not sharing. Draw dot-and-cross diagrams showing the 'before' stage (separate atoms), the 'after' stage (ions formed), then remember the final key point.

The actual ionic bond is the electrostatic forces of attraction between positive cations and negative anions. Different ionic compounds form different crystal structures depending on ion sizes.

Crystal structures vary - NaCl has each ion surrounded by 6 of the opposite type (6:6 coordination), while CsCl has 8:8 coordination because the ions are different sizes.

💡 Exam Focus: Always state that ionic bonds are electrostatic forces of attraction between oppositely charged ions.

Ionic Properties

Ionic compounds conduct electricity when molten or dissolved because ions become free to move and carry charge. In solid form, ions are locked in position so no conduction occurs.

They're brittle because applying pressure forces similarly charged ions closer together - the electrostatic repulsion then splits the crystal structure. Think of it like trying to push the wrong ends of magnets together.

High melting points result from the strong electrostatic forces between ions. Solubility depends on whether the attraction between polar water molecules and ions is stronger than the crystal's lattice energy.

💡 Quick Check: Can you explain why salt dissolves in water but not in oil?

Ionic Properties Summary

The key ionic properties are straightforward to remember: they conduct electricity when molten or dissolved (but not as solids), they're brittle under pressure, have high melting points due to strong electrostatic forces, and their solubility depends on lattice energy vs water attraction.

These properties all link back to the fundamental nature of ionic bonding - the electrostatic forces between charged ions in a crystal lattice structure.

💡 Memory Aid: Think of ionic compounds as rigid crystal structures that only become mobile when the crystal breaks down $melting/dissolving$.

Metallic Bonding Structure

Metallic bonding consists of layers of positive cations surrounded by a "sea" of delocalised electrons. The metallic bond itself is the electrostatic forces of attraction between the positive cations and the negative electron sea.

This unique structure explains all metallic properties. Electrical conductivity comes from delocalised electrons that can move freely. Malleability and ductility occur because cation layers can slide over each other without breaking bonds.

High melting points result from the strong electrostatic forces throughout the structure. The more valence electrons an atom has, the stronger the metallic bond becomes.

💡 Visual Aid: Imagine positive charges floating in a sea of negative electrons - they can move but stay attracted!

Metallic Properties Summary

The four key metallic properties - electrical conductivity, malleability, ductility, and high melting points - all stem from the same electron sea model of metallic bonding.

Understanding this connection helps you explain why metals behave so differently from ionic or covalent compounds, and why they're so useful for electrical wiring, construction, and manufacturing.

💡 Exam Success: Always link metallic properties back to the electron sea model and sliding cation layers!