Atoms, Elements, and Compounds

Every single thing around you is made of atoms - they're literally the building blocks of everything. Think of them like incredibly tiny Lego pieces that can't be broken down any further.

Inside each atom, you'll find three key particles. Protons carry a positive charge and have a mass of 1, neutrons have no charge but the same mass, and electrons zip around with a negative charge and virtually no mass. It's dead simple: protons = positive, neutrons = neutral, electrons = negative.

Here's where it gets interesting - elements are pure substances made of just one type of atom. Sodium has only sodium atoms, oxygen has only oxygen atoms. But compounds are completely different beasts - they're formed when two or more different elements join forces chemically, like water (H₂O) where hydrogen and oxygen team up.

Key Tip: Remember that in any neutral atom, the number of protons always equals the number of electrons - they balance each other out perfectly.

Mixtures and Separation Techniques

Unlike compounds, mixtures are just different substances hanging out together without any chemical bonding - think of a fruit salad where each piece keeps its own identity. The brilliant thing about mixtures is that you can separate them using some pretty clever techniques.

Filtration works when you need to separate something that won't dissolve (like sand from water) - just pour it through filter paper. Crystallisation is your go-to method when you want to get salt from salt water - just heat it up and let the water evaporate.

Distillation is perfect for separating liquids from solutions by heating them until they turn to vapour, then cooling them back down. If you've got multiple liquids mixed together, fractional distillation uses their different boiling points to separate them one by one. Chromatography is brilliant for separating dyes and pigments - you'll definitely use this in practical work.

Remember: These separation techniques only work because mixtures aren't chemically bonded - the substances keep their individual properties.

Evolution of Atomic Models

The story of how scientists figured out what atoms actually look like is absolutely fascinating - and it's a classic 6-mark question topic that you need to nail.

John Dalton kicked things off in the early 1800s, imagining atoms as solid, indestructible spheres - basically tiny snooker balls. Then J.J. Thomson discovered electrons in 1897 and created the "plum pudding model" - a positive sphere with negative electrons dotted throughout like raisins in a pudding.

Everything changed when Ernest Rutherford fired alpha particles at gold foil in 1909. Most particles sailed straight through, but some bounced back - proving atoms have a tiny, dense nucleus with electrons whizzing around it. This was revolutionary because it showed atoms are mostly empty space.

Niels Bohr refined this further in 1913, proposing that electrons orbit in fixed energy levels or shells, which explained why atoms emit specific colours of light. Finally, James Chadwick discovered neutrons in 1932, completing our modern picture of the atom.

Exam Tip: For 6-mark questions, mention each scientist, their key discovery, and how it changed our understanding - this shows clear progression of ideas.

Relative Atomic Mass and the Periodic Table

Relative atomic mass sounds complicated, but it's actually straightforward once you get the hang of it. Most elements have different versions called isotopes - same number of protons, different neutrons. You calculate the average using this formula: Ar = (mass × % abundance) ÷ 100.

Take chlorine - it has two main isotopes: Cl-35 (75%) and Cl-37 (25%). So Ar = (35×75 + 37×25)/100 = 35.5. That's why chlorine's atomic mass isn't a whole number on the periodic table.

The periodic table is brilliantly organised by increasing atomic number (number of protons). Groups are the vertical columns - elements in the same group have the same number of electrons in their outer shell, which is why they behave similarly. Periods are the horizontal rows - elements in the same period have the same number of electron shells.

Quick Check: If an element is in Group 2, Period 3, it has 2 outer electrons and 3 electron shells - the periodic table tells you everything you need to know!

Group 1: Alkali Metals and Group 7: Halogens

Group 1 metals (alkali metals) are the drama queens of the periodic table - they're incredibly reactive because they desperately want to lose their single outer electron. Sodium, potassium, lithium - they all fizz violently with water, getting more reactive as you go down the group.

When alkali metals meet water, they form a metal hydroxide plus hydrogen gas. The reaction for sodium is: 2Na + 2H₂O → 2NaOH + H₂. You'll see the hydrogen bubbling off, and sometimes the reaction is so energetic that the hydrogen actually ignites.

Group 7 halogens are completely different - they're non-metals with 7 outer electrons, desperately wanting to gain one more. Interestingly, their reactivity decreases as you go down the group (opposite to Group 1). Fluorine is the most reactive, then chlorine, bromine, and iodine.

Halogens can kick each other out of compounds in displacement reactions - a more reactive halogen will displace a less reactive one. So chlorine can displace bromine: Cl₂ + 2KBr → 2KCl + Br₂.

Memory Trick: Group 1 gets MORE reactive going down, Group 7 gets LESS reactive going down - they're opposites in every way!

Noble Gases and Types of Bonding

Group 0 noble gases are the chilled-out elements of the periodic table. With complete outer electron shells, they're incredibly unreactive - they've got everything they need and aren't interested in bonding with anyone. Their boiling points increase down the group due to stronger intermolecular forces between larger atoms.

Now for the big three types of bonding - this is where chemistry gets really interesting. Ionic bonding happens between metals and non-metals when electrons get transferred. The metal loses electrons to become a positive ion, the non-metal gains them to become negative, and oppositely charged ions attract strongly.

Covalent bonding is completely different - non-metal atoms share pairs of electrons rather than transferring them. Think of it like sharing a pizza rather than giving it away completely. Water (H₂O) and methane (CH₄) are classic examples where atoms share electrons to fill their outer shells.

Metallic bonding creates that characteristic "sea of electrons" - positive metal ions sit in a cloud of delocalised electrons that can move freely. This explains why metals conduct electricity and can be hammered into shapes.

Key Insight: The type of bonding determines the properties - ionic compounds conduct when molten, covalent compounds often have low melting points, and metals are malleable and conductive.

Properties of Different Structures

Understanding how bonding affects properties is crucial for predicting how materials will behave. Ionic compounds form giant lattices with high melting points because those electrostatic forces between ions are incredibly strong. They only conduct electricity when the ions are free to move - either when molten or dissolved in water.

Simple molecular covalent substances like water or carbon dioxide have relatively low melting points. While the bonds within molecules are strong, the forces between separate molecules are weak, so they're easy to separate.

Giant covalent structures are in a league of their own. Diamond has each carbon atom bonded to four others in a 3D network, making it incredibly hard with an extremely high melting point. Graphite is different - each carbon bonds to only three others, creating layers that can slide past each other, plus it conducts electricity thanks to delocalised electrons.

Graphene is essentially one layer of graphite and is revolutionising electronics because it's incredibly strong yet flexible and conducts electricity brilliantly. Metallic structures give metals their unique properties - they conduct heat and electricity well, and you can hammer them into different shapes because the layers of atoms can slide over each other.

Exam Success: Always link structure to properties - if you can explain why a material behaves the way it does based on its bonding, you'll ace these questions every time.