GCSE WJEC Chemistry Unit 1 Guide

Basic Chemistry Concepts

Ever wondered how scientists separate salt from seawater or why some substances mix whilst others don't? Mixtures are simply two or more elements or compounds that aren't chemically stuck together - think of a bowl of cereal where you can still pick out individual pieces.

You can separate mixtures using clever techniques. Filtration works for insoluble solids (like sand from water), crystallisation recovers dissolved solids, distillation separates liquids with different boiling points, and chromatography splits up substances that travel at different speeds.

Isotopes are like identical twins with different weights - same element, but different numbers of neutrons. The relative atomic mass you see on the periodic table is actually an average of all these different versions. Scientists calculate this using abundance (how common each isotope is) and their individual masses.

Quick Tip: Remember that elements contain only one type of atom, whilst compounds mix different types together - like the difference between a bag of only red sweets versus a mixed bag!

Ionic Bonding

When metals meet non-metals, it's like a generous friend giving away their belongings! Ionic bonding happens when metal atoms transfer electrons to non-metal atoms instead of sharing them. The metal becomes a positive ion (lost electrons), whilst the non-metal becomes a negative ion (gained electrons).

These oppositely charged ions attract each other with strong electrostatic forces, creating ionic compounds. It's like powerful magnets sticking together - the attraction is incredibly strong.

Ionic compounds have some predictable properties that make perfect sense. They have high melting points because those electrostatic attractions are tough to break. They don't conduct electricity when solid (ions can't move), but they do conduct when melted or dissolved (ions become free to carry charge).

Understanding these properties helps explain why salt dissolves in water and why ionic compounds are often used in batteries and electronics.

Key Point: Think of ionic bonding as a permanent loan - the metal "lends" electrons to the non-metal, but they're so attracted to each other that they stick together!

The Periodic Table and Groups

Mendeleev was brilliant but had gaps in his periodic table because some elements hadn't been discovered yet. Today's periodic table orders elements by atomic number with no gaps - metals on the left, non-metals on the right.

Group 0 elements (noble gases) are the chemistry equivalent of content cats - they have full outer electron shells so they don't react with anything. They exist as single atoms and their boiling points increase as you go down the group.

Group 1 elements (alkali metals) are the complete opposite - they're incredibly reactive because they desperately want to lose that single outer electron. They react with oxygen, chlorine, and water to form predictable products. The name "alkali metals" comes from their reaction with water producing alkaline solutions.

Memory Trick: Noble gases are "noble" because they don't need to react with anyone - they're perfectly happy on their own!

Group Properties and Trends

Group 1 reactivity increases as you go down because the atoms get bigger and the outer electron gets further from the nucleus. It's like trying to hold onto a balloon on a longer string in the wind - much harder! This weaker attraction makes it easier to lose that outer electron.

Group 7 elements (halogens) are non-metals that exist as molecules made of atom pairs. Fluorine and chlorine are gases, bromine is liquid, and iodine is solid at room temperature. Unlike Group 1, their reactivity decreases down the group because it becomes harder to gain that extra electron.

The structure of substances affects their properties dramatically. Giant covalent structures have high melting points because you need lots of energy to break strong covalent bonds. Small molecules have low melting points due to weak forces between molecules, whilst large molecules fall somewhere in between.

Alloys solve the problem of pure metals being too soft - adding different atoms disrupts the regular arrangement so layers can't slide over each other easily.

Pattern Spotting: Group 1 gets more reactive going down, Group 7 gets less reactive - they're like opposite trends!

Covalent Bonding and States of Matter

Covalent bonding happens between non-metals that share electrons rather than transferring them. It's like sharing a pizza - each atom contributes electrons to make both atoms happy with full outer shells.

The particle model shows how substances change state through melting, freezing, boiling, and condensing. The energy needed for these changes depends on forces between particles - stronger forces mean higher melting and boiling points.

Covalent structures come in three main types. Small molecules have strong bonds within molecules but weak forces between them (low melting points). Large molecules have stronger intermolecular forces than small ones. Giant covalent structures like diamond have billions of atoms all connected by strong covalent bonds (very high melting points).

Real-World Connection: Diamond's giant covalent structure makes it incredibly hard, whilst graphite (also carbon) has a different structure that makes it soft enough to use in pencils!

Chemical Calculations with Moles

Group 7 reactivity decreases down the group because atoms get larger and it becomes harder to attract that extra electron needed to fill the outer shell. More reactive halogens can displace less reactive ones from compounds - it's like a stronger person taking someone's seat!

Moles are chemistry's counting system. Just like a dozen equals 12, a mole equals 6.02 × 10²³ particles. You can have a mole of anything - biscuits, atoms, or molecules! The genius is that one mole of any substance has a mass equal to its relative atomic or formula mass in grams.

Concentration in mol/dm³ makes calculations much easier: concentration = moles of solute ÷ volume of solution. For gases, you can find moles by dividing volume by 24 dm³ (or 24,000 cm³).

Why Moles Matter: Using moles lets chemists count atoms without actually counting billions of individual particles - it's like using dozens to count eggs!

Mole Calculations in Practice

Working out moles is straightforward once you get the hang of it: moles = mass ÷ molar mass. The molar mass is just the relative atomic mass (Ar) or relative formula mass (Mr) expressed in grams per mole.

Let's make this real with examples. One mole of lithium oxide (Li₂O) weighs 30g because you add up the masses: lithium $7 × 2 = 14$ plus oxygen $16 × 1 = 16$. For 250g of iron, you'd have 4.46 moles since iron's molar mass is 56g/mol.

These calculations aren't just academic exercises - they're essential for making sure chemical reactions have the right amounts of each substance. Too little of one reactant and you'll have leftovers; too much and you're wasting materials.

Mastering mole calculations gives you the power to predict exactly what will happen in chemical reactions and how much product you'll make.

Success Tip: Practice with simple examples first - once you can calculate moles for single elements, compounds become much easier to handle!