Applied Science - Unit 1

You're diving into one of the most practical areas of science that connects directly to real-world applications. Applied science takes the theory you learn and shows you exactly how it works in industries, medicine, and technology around you.

This chemistry unit forms the foundation for understanding everything from why your phone battery works to how medicines are developed. Master these concepts and you'll have the tools to tackle any chemistry challenge.

Quick Tip: Applied science questions often ask you to explain real-world examples, so always think about practical applications as you learn!

Chemistry Fundamentals

Chemistry at this level is all about understanding patterns and predicting behaviour. You'll discover that atoms aren't random - they follow clear rules that make perfect sense once you get the hang of them.

The periodic table becomes your best friend here, showing you exactly how elements behave based on their position. Think of it as a massive cheat sheet that tells you everything about an element's personality.

Everything you learn builds on itself, so each concept connects to create a complete picture of how matter works at the atomic level.

Remember: Chemistry is about recognising patterns - once you spot them, predictions become much easier!

Sub Shells - Electronic Structure

Think of sub shells as different-sized car parks for electrons around the nucleus. Each type can hold a specific number of electrons, and they fill up in a predictable order.

S block holds 2 electrons, P block takes 6, D block accommodates 10, and F block fits 14. As you move through the periods, more sub shells become available: Period 1 just has S, Period 2 adds P, Period 3 includes D, and Period 4 gets the full set with F.

This filling pattern explains why the periodic table has its distinctive shape and why elements in the same group behave similarly.

Memory Trick: Remember "SPeeDy Fast" for the order S, P, D, F with their capacities 2, 6, 10, 14!

First Ionisation Energy Trends

First ionisation energy measures how much energy you need to remove an electron from an atom - think of it as how tightly the nucleus grips its electrons.

Down a group, ionisation energy decreases because electrons get further from the nucleus. More shells mean weaker nuclear attraction, making electrons easier to remove - like trying to hold onto something with a longer arm.

Across a period, ionisation energy increases as more protons pull on electrons in the same shell. The atomic radius shrinks, creating stronger nuclear attraction that holds electrons more tightly.

Exam Focus: Questions often ask you to compare elements - use these trends to predict which will have higher ionisation energy!

Atomic and Ionic Radius Patterns

Atomic radius follows the opposite pattern to ionisation energy, and understanding why makes perfect sense when you think about it.

Down a group, atoms get bigger as each period adds another electron shell. It's like adding floors to a building - the structure gets taller with each addition.

Across a period, atoms actually shrink despite gaining electrons. More protons create stronger nuclear attraction, pulling the same electron shell closer to the nucleus. More electrons don't mean more space - they mean tighter packing.

Quick Check: If you can predict size changes, you can predict most other periodic trends too!

Electron Affinity

Electron affinity measures how much non-metals "want" to gain an electron. Think of it as the opposite of ionisation energy - instead of removing electrons, you're adding them.

Non-metals typically have high electron affinity because gaining electrons helps them achieve a stable electron configuration. The atom must be in a gaseous state for this measurement to be standardised.

This property helps explain why non-metals form negative ions so readily and why they're found on the right side of the periodic table.

Key Point: High electron affinity means the atom really wants to gain electrons - perfect for forming ionic compounds!

Melting and Boiling Points

The melting and boiling points of substances depend entirely on what holds their particles together. Breaking stronger bonds requires more energy and higher temperatures.

Metallic bonding creates high melting points because you're breaking the electrostatic attraction between positive metal ions and delocalised electrons. Giant covalent structures also need lots of energy to overcome strong covalent bonds throughout the structure.

Simple covalent molecules have low melting points because you're only breaking weak intermolecular forces (Van der Waals forces), not the actual covalent bonds within molecules.

Exam Strategy: Always identify the bonding type first - it instantly tells you whether to expect high or low melting points!

Physical Properties of Metals

Metals have their distinctive properties because of metallic bonding and delocalised electrons. These "sea of electrons" can move freely, explaining most metallic behaviour.

Electrical and thermal conductivity work because delocalised electrons carry charge and energy through the metal structure. Malleability happens because metal ion layers can slide over each other while maintaining electrostatic attraction.

Ductility allows metals to be drawn into wires for the same reason - layers slide but the bonding holds everything together. Metals are also typically shiny and insoluble in water.

Real-World Link: Every time you use copper wires, aluminium foil, or steel tools, you're seeing these properties in action!

Periodic Trends - Key Patterns

Master these periodic trends and you'll be able to predict element behaviour across the entire periodic table. The patterns are surprisingly logical once you understand the underlying physics.

Down groups: More electron shells mean weaker nuclear attraction and larger atomic radius. This leads to lower ionisation energy and different chemical reactivity patterns.

Across periods: Same electron shell but more protons create stronger nuclear attraction and smaller atomic radius. This increases ionisation energy and affects how readily atoms form bonds.

Study Tip: Draw arrows on your periodic table showing these trends - visual memory makes exam questions much easier!

Hydrogen Bonding

Hydrogen bonding creates the strongest type of intermolecular force, but it's very selective about when it occurs. You only get hydrogen bonding when hydrogen connects to nitrogen, oxygen, or fluorine.

These N, O, F elements are highly electronegative, creating polar molecules where hydrogen becomes slightly positive. The positive hydrogen then attracts lone pairs on other N, O, or F atoms, forming the hydrogen bond.

This explains water's unusually high boiling point of 100°C - much higher than you'd expect for such a small molecule. Without hydrogen bonding, water would be a gas at room temperature!

Memory Aid: "NOF" (Nitrogen, Oxygen, Fluorine) - the only elements that create hydrogen bonding with hydrogen!