Periodic Trends

The periodic table isn't just a random arrangement - it reveals clear patterns that explain how elements behave. These trends help you predict everything from atomic size to how easily electrons can be removed.

Covalent radius measures atomic size, electronegativity shows how strongly atoms attract electrons in bonds, and ionisation energy tells you how much energy is needed to remove an electron. All these properties follow predictable patterns across periods and down groups.

Moving across a period, atoms get smaller, more electronegative, and harder to ionise because the nuclear charge increases. Going down a group does the opposite - atoms get larger and less electronegative because outer electrons are further from the nucleus and shielded by inner electron shells.

Key insight: The number of protons and electron shells determines all these trends - master this and you can predict how any element will behave!

Bonding and Structure of Elements

Elements don't all exist in the same way - some form giant metal lattices, others exist as single atoms, and some create complex networks. Understanding these structures explains why copper conducts electricity but sulfur doesn't.

Metallic bonding occurs when metal atoms lose electrons to create a "sea" of delocalised electrons around positive metal ions. The more outer electrons an atom has, the stronger these metallic bonds become. This explains why aluminium is harder than sodium.

Covalent bonding involves atoms sharing electrons rather than losing them. Elements like carbon can form either small molecules (like carbon dioxide) or giant networks (like diamond). Monatomic elements like noble gases exist as single atoms held together only by weak forces.

The key difference is electron behaviour - metals have mobile electrons that can conduct electricity, whilst covalent structures keep electrons localised between specific atoms.

Remember: The number of outer electrons determines bonding type and strength - this affects everything from melting points to electrical conductivity!

Bonding in Compounds

When different elements combine, the type of bond formed depends entirely on their electronegativity difference. This determines whether electrons are shared equally, shared unequally, or transferred completely.

Non-polar covalent bonds form between atoms with identical electronegativity values. Polar covalent bonds occur with small electronegativity differences, creating slightly positive and negative ends. Ionic bonds result from large electronegativity differences where electrons transfer completely.

Ionic compounds consist of positive metal ions and negative non-metal ions in giant lattice structures. They don't conduct when solid because ions can't move, but they do conduct when molten or dissolved because ions become mobile.

The greater the electronegativity difference, the more ionic character a compound has. This affects properties like melting point, electrical conductivity, and solubility.

Top tip: Electronegativity difference is your key to predicting bond type - memorise the boundaries between covalent, polar covalent, and ionic!

Covalent Network vs Molecular Compounds

Covalent network compounds and covalent molecular compounds couldn't be more different despite both involving covalent bonding. The key difference lies in their structure and the forces holding them together.

Covalent network compounds like diamond have continuous networks of strong covalent bonds throughout their entire structure. This means extremely high melting points because you'd need to break countless strong bonds to melt them.

Covalent molecular compounds contain small, discrete molecules held together by weak intermolecular forces. These weak forces mean low melting points - that's why most are gases, liquids, or soft solids at room temperature.

Neither type conducts electricity because electrons aren't free to move, but their physical properties are vastly different due to the forces between particles.

Quick check: If it's a gas or liquid at room temperature, it's definitely molecular, not network!

Polar vs Non-polar Molecules

Polarity determines how molecules interact with each other and affects everything from solubility to boiling points. A molecule's polarity depends on both its bonds and its shape.

Non-polar molecules arise in two situations: when all bonds are non-polar (like in hydrogen gas), or when polar bonds cancel out due to symmetrical molecular shape (like carbon tetrachloride). These molecules have no permanent positive or negative ends.

Polar molecules have polar bonds AND asymmetrical shapes, creating permanent positive and negative sides. Water is the classic example - its bent shape means the polar O-H bonds don't cancel out.

The crucial principle is "like dissolves like" - polar substances dissolve in polar solvents (like salt in water), whilst non-polar substances dissolve in non-polar solvents (like oil in petrol).

Shape matters: Even polar bonds can create non-polar molecules if the shape is symmetrical enough for dipoles to cancel!

Van der Waals Forces - London Dispersion Forces

Van der Waals forces are the weak attractions between all atoms and molecules. When you boil water or melt ice, you're breaking these forces, not the strong covalent bonds within molecules.

London dispersion forces (LDF) are the weakest type but exist between ALL particles. They arise from temporary, uneven electron distribution creating temporary dipoles. One atom's temporary positive side attracts neighbouring electrons, creating an induced dipole.

LDF strength depends entirely on electron count - more electrons mean stronger temporary dipoles and stronger attractions. This explains why larger atoms and molecules have higher boiling points than smaller ones with the same type of bonding.

These forces might be weak individually, but in large molecules with many electrons, they can add up to create surprisingly strong attractions between molecules.

Electron count is key: More electrons = stronger LDF = higher melting and boiling points!

Stronger Van der Waals Forces

Permanent dipole-dipole attractions occur between polar molecules where the positive end of one molecule attracts the negative end of another. These are stronger than London dispersion forces when comparing molecules of similar size.

Hydrogen bonding is a special, extra-strong type of dipole-dipole attraction. It only occurs when hydrogen is directly bonded to nitrogen, oxygen, or fluorine - the three most electronegative elements. This creates extremely polar bonds.

Hydrogen bonds explain many of water's unusual properties, including why ice floats. When water freezes, hydrogen bonding creates an open structure that's less dense than liquid water.

The strength order is always: hydrogen bonds > dipole-dipole attractions > London dispersion forces. However, all van der Waals forces are much weaker than covalent or ionic bonds.

H-bonding rule: Only H attached directly to N, O, or F can form hydrogen bonds - this creates water's unique properties!

Physical Properties and Van der Waals Forces

The type of van der Waals forces present determines physical properties like melting point, boiling point, viscosity, and solubility. Understanding these connections helps you predict how substances will behave.

For fair comparisons between molecules, size and electron count should be similar so London dispersion forces are roughly equal. Then you can see how polarity and hydrogen bonding affect properties.

Hydrogen bonding creates the highest melting/boiling points, followed by dipole-dipole attractions, then London dispersion forces alone. Polar molecules dissolve in polar solvents, whilst non-polar molecules dissolve in non-polar solvents.

When answering exam questions, remember the key distinction: periodic trends (like atomic radius) relate to nuclear charge and electron shells, whilst physical properties (like boiling point) relate to bonding and intermolecular forces.

Exam strategy: Periodic table questions need nuclear charge explanations; physical property questions need bonding and van der Waals explanations!

Quick Reference Guide

This flowchart approach helps you systematically work out what type of forces exist between molecules and predict their properties.

Start with electronegativity difference to determine if bonds are polar. Then consider molecular shape - do the polar bonds create clear positive and negative sides, or do they cancel out due to symmetry?

For non-polar molecules, only London dispersion forces exist. For polar molecules, check if hydrogen is directly attached to nitrogen, oxygen, or fluorine to determine if hydrogen bonding occurs.

This systematic approach ensures you don't miss any forces and can confidently predict properties like boiling points, solubility, and conductivity.

Follow the flowchart: Start with electronegativity, then shape, then check for H-bonding - this covers all possibilities!