Atomic Structure - The Building Blocks of Everything

Ever wondered what's inside an atom? Scientists have been figuring this out for centuries, starting with Dalton's solid sphere model, then moving through Thomson's plum pudding model to Rutherford's nuclear model, and finally Chadwick's discovery of neutrons.

Today we know atoms contain three main particles. Protons have a +1 charge and mass of 1, neutrons have no charge but the same mass, and electrons have a -1 charge with virtually no mass. The protons and neutrons pack tightly into the nucleus at the atom's centre, whilst electrons orbit around it.

Here's what makes atoms balanced: they always have equal numbers of protons and electrons, so the positive and negative charges cancel out. The mass number tells you the total protons plus neutrons, whilst the atomic number is just the number of protons (which defines what element it is).

Key insight: The nucleus is incredibly tiny compared to the whole atom, but contains nearly all the mass - imagine a marble in a football stadium!

Isotopes and the Periodic Table

Isotopes are like different versions of the same element - they have identical numbers of protons but different numbers of neutrons. This affects their mass, which is why we calculate relative atomic mass as an average of all isotope masses.

The periodic table arranges elements by increasing atomic number (not weight, as originally thought). Elements in the same group (column) have similar properties because they have the same number of outer electrons. Elements in the same period (row) have the same number of electron shells.

Electronic configuration follows a simple pattern: the first shell holds up to 2 electrons, the second and third shells hold up to 8 each. Your position on the periodic table tells you exactly how many electron shells (period number) and outer electrons (group number) an element has.

Quick tip: Metals tend to lose outer electrons whilst non-metals tend to gain them - this difference creates their contrasting properties!

Ionic Bonding - When Atoms Transfer Electrons

Ionic bonds form when atoms transfer electrons to create charged particles called ions. Metals $groups 1-3$ lose electrons to become positive cations, whilst non-metals $groups 5-7$ gain electrons to become negative anions.

When naming ionic compounds, remember the rules: -ide endings for two-element compounds (like NaCl), and -ate endings for compounds containing oxygen plus other elements (like CaCO₃). Always balance the charges so the overall compound is neutral.

These ions arrange themselves in a lattice structure - a regular, repeating pattern held together by strong electrostatic forces between oppositely charged ions. Think of it like a 3D puzzle where positive and negative pieces attract and lock together.

Real example: When sodium loses an electron and chlorine gains it, they form Na⁺ and Cl⁻ ions that create table salt (NaCl)!

Covalent Bonding - Sharing is Caring

Covalent bonds work completely differently from ionic bonds - instead of transferring electrons, atoms share pairs of electrons between them. This happens between non-metals and results in the formation of molecules.

The shared electrons create a strong bond that holds the atoms together within each molecule. You can represent this with dot and cross diagrams, showing how electron pairs are shared between atoms.

Remember: Covalent bonding = sharing electrons between non-metals to form molecules!

Types of Substances and Their Properties

Understanding ionic substances means knowing they have high melting and boiling points because you need lots of energy to break those strong electrostatic forces. They conduct electricity when molten or dissolved because the ions can move freely and carry charge.

Simple molecular substances (covalent) have strong bonds within molecules but weak forces between them. This gives them low melting and boiling points and poor electrical conductivity since there are no free electrons or ions.

Giant molecular structures like diamond create massive lattices of covalently bonded atoms, resulting in high melting points. Metallic substances have their own special bonding with delocalised electrons that can move freely, explaining why metals conduct electricity so well.

Pattern spotted: The type of bonding determines the properties - strong bonds mean high melting points, mobile charges mean electrical conductivity!

Carbon Allotropes - Same Element, Different Structures

Diamond arranges each carbon atom bonded to four others in a tetrahedral structure, creating incredibly strong covalent bonds throughout. This makes it extremely hard with a high melting point, perfect for cutting tools and jewellery, but it can't conduct electricity.

Graphite takes a different approach - each carbon has only three bonds, creating layered structures with weak forces between layers. These layers can slide over each other (making it soft and slippery), and it conducts electricity thanks to delocalised electrons.

Graphene is essentially a single layer of graphite - incredibly strong, flexible, and an excellent electrical conductor. Fullerenes like Buckminsterfullerene (C₆₀) form when graphene sheets roll into hollow balls or tubes, creating unique properties for different applications.

Mind-blowing fact: Diamond and graphite are both pure carbon, but their different structures give them completely opposite properties!

Polymers and Model Limitations

Polymers are massive molecules built from many smaller units called monomers - think of them like chemical Lego blocks that connect to form long chains.

Every model has limitations. Dot and cross diagrams miss 3D information, ball and stick models don't show electron numbers or bond types, 2D representations can't show actual shapes, and 3D models often can't display the complete structure.

Study tip: Use different models together to get the full picture - each one shows you something the others miss!

Metals vs Non-metals - The Great Divide

Metals have that distinctive shiny appearance and excel at conducting both electricity and heat through their delocalised electrons and atomic vibrations. Their atoms arrange in layers that can slide over each other, making them malleable (shapeable) and ductile (stretchable) whilst maintaining strength.

The metallic lattice structure consists of positive metal ions surrounded by a "sea" of delocalised electrons that can move freely. This explains their high density, strength, and excellent conductivity properties.

Non-metals are essentially the opposite - dull appearance, poor conductors of heat and electricity, lower density, and much lower melting and boiling points due to their different bonding arrangements.

Key difference: Metals have mobile electrons that can move freely, whilst non-metals keep their electrons locked in specific bonds!