💥 Atomic Structure
Atoms make up everything — every element, object, and living thing. They’re the smallest part of an element that still has its chemical properties.
⚡ Protons, Neutrons & Electrons
| Particle | Charge | Relative Mass | Location |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | -1 | ~0 | Shells around nucleus |
💡 Why protons are positive & electrons are negative:
- Protons contain a fundamental positive charge due to their internal quark structure.
- Electrons carry an equal but opposite negative charge.
- These opposite charges attract each other, holding the atom together.
🧲 The atom’s nucleus is positively charged (thanks to protons), and electrons are held in orbit by electrostatic forces of attraction.
🔬 Elements and Compounds
👩🔬 Element:
- A pure substance made of only one type of atom.
- Example: Oxygen (O₂), Gold (Au), Hydrogen (H₂)
🧪 Compound:
- When two or more elements chemically combine in fixed proportions.
- Example:
- Water (H₂O): Hydrogen + Oxygen
- Carbon dioxide (CO₂): Carbon + Oxygen
🧩 Compounds have completely different properties from the elements they’re made from.
🧾 Chemical Formula
A chemical formula tells you which elements are present in a substance and how many atoms of each are in a single unit of that substance.
🔤 Symbols & Numbers
- Element symbols come from the periodic table (e.g., H = hydrogen, O = oxygen).
- Subscripts (small numbers) show how many atoms of that element are present.
- H₂O → 2 H atoms, 1 O atom.
- CO₂ → 1 C atom, 2 O atoms.
- If no number appears after the symbol, it means 1 atom.
📐 Types of formula
- Molecular formula — actual number of atoms in a molecule (e.g., C₆H₁₂O₆ for glucose).
- Empirical formula — simplest whole-number ratio of atoms (e.g., CH₂O is the empirical formula for glucose).
- Ionic formula — shows ratio of ions in an ionic compound (e.g., NaCl, CaCl₂). Ionic formulas reflect charges balanced, not “molecules”.
🔗 How ionic formulae are made
- Metals lose electrons → form positive ions (cations).
- Non-metals gain electrons → form negative ions (anions).
- Swap charges and simplify to get the formula:
- Magnesium (Mg²⁺) + Chloride (Cl⁻) → MgCl₂ (because one Mg²⁺ needs two Cl⁻).
📎 Examples
- Water (molecular): H₂O
- Methane (molecular): CH₄
- Sodium chloride (ionic): NaCl
- Calcium oxide (ionic): CaO (Ca²⁺ and O²⁻ → 1:1)
- Aluminium oxide: Al₂O₃ (Al³⁺ needs three O²⁻ → balance to 2 Al and 3 O)
🔬 State symbols (often in equations):
- (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous (dissolved in water)
E.g. NaCl(s) → Na⁺(aq) + Cl⁻(aq)
🧠 Quick tips
- For ionic compounds: criss-cross ion charges to find the simplest formula then simplify if needed.
- For empirical vs molecular: molecular might be a multiple of empirical (e.g., C₂H₄ empirical = CH₂).
- Use the periodic table to find charge tendencies: group number often tells you common ion charge for simple ions (Group 1 → +1, Group 7 → −1).
📊 The Periodic Table
🧠 History – The Mendeleev Masterplan:
- Before 1869, scientists grouped elements by atomic weight and chemical properties.
- Dmitri Mendeleev organised them into a table and left gaps for undiscovered elements — predicting their properties!
- Later, the discovery of protons showed that elements are better arranged by atomic number (number of protons).
💡 Modern Periodic Table:
- Arranged in order of atomic number.
- Groups (columns): elements with similar properties (same number of outer shell electrons).
- Periods (rows): show the number of electron shells.
📍Examples:
- Group 1 – Alkali metals (very reactive, 1 outer electron)
- Group 7 – Halogens (reactive non-metals, 7 outer electrons)
- Group 0 – Noble gases (unreactive, full outer shell)
💫 Electron Energy Levels (Shells)
Electrons orbit the nucleus in energy levels called shells.
- 1st shell = holds up to 2 electrons
- 2nd shell = holds up to 8 electrons
- 3rd shell = holds up to 8 electrons
⚗️ Example: Sodium (Na)
- Atomic number = 11 → 2, 8, 1
- That single outer electron makes sodium highly reactive!
💡 Outer shell electrons determine an element’s reactivity and chemical behaviour.
🧫 Mixtures and Separation Techniques
A mixture contains two or more substances not chemically bonded together.
They can be separated by physical methods:
1️⃣ Filtration
- Separates insoluble solids from liquids.
- Example: Sand + Water → filter → sand stays, water passes through.
2️⃣ Crystallisation
- Separates a dissolved solid (solute) from a solution.
- Heat gently → evaporate solvent → crystals form.
3️⃣ Chromatography
- Separates mixtures of coloured compounds (e.g. inks).
- Dots of colour travel up paper at different speeds.
4️⃣ Distillation
- Separates a liquid from a solution based on boiling points.
- The solvent is heated → evaporates → condensed → collected.
- Example: Getting pure water from seawater.
5️⃣ Fractional Distillation
- Used when liquids have similar boiling points.
- Example: Separating ethanol from water or crude oil fractions.
🌡️ Simple Distillation in Action
- Heat the mixture until the solvent boils.
- Vapour rises and passes through a condenser.
- Cooler water runs around the condenser to cool the vapour.
- The solvent condenses and drips into a separate container.
💧 Result: Pure solvent separated from the solute!
📚 Useful Sources for Revision
BBC Bitesize – Atomic Structure & Separation Techniques
Clear visuals, interactive quizzes, and Mendeleev’s story explained.
Physics & Maths Tutor – AQA Combined Science Unit 1
Concise topic summaries and exam-style practice.
Save My Exams – Chemistry Paper 1 Notes
Printable guides with example questions.
Quizlet – Atomic Structure & Separation Flashcards
Perfect for memorising definitions and key processes.
YouTube – Freesciencelessons & Cognito
Visual walkthroughs of every concept (great for quick recap).
