Chemistry starts with understanding the building blocks of matter -...
Chemistry250 views·Updated 13 Jul 2026·4 pagesA-Level Chemistry Chapter 2: Understanding Atoms, Ions, and Compounds
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Atomic Structure and Isotopes
Atoms are like tiny solar systems with a dense centre called the nucleus containing protons (positive charge) and neutrons (no charge). Electrons (negative charge) whizz around the nucleus in shells, but they're incredibly light - about 1/1836th the mass of a proton!
Here's what makes atoms tick: protons and neutrons have roughly the same mass, whilst electrons are practically weightless. The number of protons equals the number of electrons, so atoms are electrically neutral overall. Neutrons act like nuclear glue, preventing the positively charged protons from repelling each other.
The atomic number is simply the number of protons in an atom - it's like an element's ID card. Every carbon atom has 6 protons, every oxygen atom has 8 protons, and so on.
Isotopes are atoms of the same element with different numbers of neutrons. They react identically in chemical reactions because reactions involve electrons, not neutrons. However, heavier isotopes might have slightly different physical properties like higher melting points.
💡 Quick Tip: Remember that ions are just atoms that have gained or lost electrons - cations are positive (lost electrons), anions are negative (gained electrons).
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Relative Mass
Since atoms are ridiculously tiny, scientists needed a clever way to compare their masses. Enter carbon-12 - the international standard that's defined as exactly 12 atomic mass units (12u). Everything else gets measured relative to this.
Relative isotopic mass compares an isotope's mass to 1/12th of a carbon-12 atom's mass. For most practical purposes, this equals the mass number (protons + neutrons). No units needed since it's just a ratio!
Relative atomic mass (Ar) is trickier because most elements contain multiple isotopes. It's a weighted mean that considers both the percentage abundance and mass of each isotope. Think of it like calculating your overall exam grade when different tests count for different percentages.
Mass spectrometers determine these values by vaporising samples, ionising them, then separating ions based on their mass-to-charge ratio (m/z). Heavier ions move slower and deflect less than lighter ones.
💡 Calculation Hack: To find relative atomic mass, multiply each isotope's abundance by its mass number, add them up, then divide by 100. Easy!
3
of 4
Formulae and Equations
Writing chemical formulae might seem daunting, but there's a logical system behind it. Ionic compounds form when metals lose electrons (becoming cations) and non-metals gain electrons (becoming anions). Both try to achieve the stable electronic structure of the nearest noble gas.
Binary compounds contain just two elements. Name the metal first, then change the non-metal's ending to '-ide' (like sodium chloride). Some metals, especially transition metals, can form multiple ions - that's why we see copper(II) for Cu²⁺.
Polyatomic ions are groups of atoms that stick together and behave as single units. Key ones include ammonium (NH₄⁺), hydroxide (OH⁻), nitrate (NO₃⁻), carbonate (CO₃²⁻), and sulfate (SO₄²⁻). These pop up everywhere in chemistry!
Writing formulae is like solving a puzzle - the positive and negative charges must balance to zero. Use subscripts to show how many of each ion you need, and brackets around polyatomic ions when there's more than one.
💡 Memory Aid: Think "positive + negative = zero" - the charges must always cancel out completely in ionic compounds.
4
of 4
Chemical Equations
Chemical equations are chemistry's shorthand for describing reactions. Most elements appear as single atoms (like Na or Mg), but some exist as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂. Remember "HOFBrINCl" to recall these pairs!
State symbols tell you the physical state: for solid, for liquid, for gas, and (aq) for dissolved in water (aqueous). These matter because the same substance can behave differently in different states.
Balancing equations follows one golden rule: you can't change the chemical formulae, only add numbers in front of them. It's like solving a mathematical puzzle where atoms on both sides must match exactly. Start with the most complex molecule and work systematically through each element.
The key is patience and practice. Check your work by counting atoms of each element on both sides - they must be equal. Use brackets carefully when dealing with polyatomic ions, remembering that subscripts outside brackets multiply everything inside.
💡 Pro Strategy: Always balance the most complicated molecule first, then work through hydrogen and oxygen last - they're usually the easiest to adjust.
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Chemistry250 views·Updated 13 Jul 2026·4 pagesA-Level Chemistry Chapter 2: Understanding Atoms, Ions, and Compounds
Chemistry starts with understanding the building blocks of matter - atoms, ions, and how they combine to form compounds. This topic covers everything from the tiny particles inside atoms to writing chemical equations that describe reactions.
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Atomic Structure and Isotopes
Atoms are like tiny solar systems with a dense centre called the nucleus containing protons (positive charge) and neutrons (no charge). Electrons (negative charge) whizz around the nucleus in shells, but they're incredibly light - about 1/1836th the mass of a proton!
Here's what makes atoms tick: protons and neutrons have roughly the same mass, whilst electrons are practically weightless. The number of protons equals the number of electrons, so atoms are electrically neutral overall. Neutrons act like nuclear glue, preventing the positively charged protons from repelling each other.
The atomic number is simply the number of protons in an atom - it's like an element's ID card. Every carbon atom has 6 protons, every oxygen atom has 8 protons, and so on.
Isotopes are atoms of the same element with different numbers of neutrons. They react identically in chemical reactions because reactions involve electrons, not neutrons. However, heavier isotopes might have slightly different physical properties like higher melting points.
💡 Quick Tip: Remember that ions are just atoms that have gained or lost electrons - cations are positive (lost electrons), anions are negative (gained electrons).
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Relative Mass
Since atoms are ridiculously tiny, scientists needed a clever way to compare their masses. Enter carbon-12 - the international standard that's defined as exactly 12 atomic mass units (12u). Everything else gets measured relative to this.
Relative isotopic mass compares an isotope's mass to 1/12th of a carbon-12 atom's mass. For most practical purposes, this equals the mass number (protons + neutrons). No units needed since it's just a ratio!
Relative atomic mass (Ar) is trickier because most elements contain multiple isotopes. It's a weighted mean that considers both the percentage abundance and mass of each isotope. Think of it like calculating your overall exam grade when different tests count for different percentages.
Mass spectrometers determine these values by vaporising samples, ionising them, then separating ions based on their mass-to-charge ratio (m/z). Heavier ions move slower and deflect less than lighter ones.
💡 Calculation Hack: To find relative atomic mass, multiply each isotope's abundance by its mass number, add them up, then divide by 100. Easy!
3
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Formulae and Equations
Writing chemical formulae might seem daunting, but there's a logical system behind it. Ionic compounds form when metals lose electrons (becoming cations) and non-metals gain electrons (becoming anions). Both try to achieve the stable electronic structure of the nearest noble gas.
Binary compounds contain just two elements. Name the metal first, then change the non-metal's ending to '-ide' (like sodium chloride). Some metals, especially transition metals, can form multiple ions - that's why we see copper(II) for Cu²⁺.
Polyatomic ions are groups of atoms that stick together and behave as single units. Key ones include ammonium (NH₄⁺), hydroxide (OH⁻), nitrate (NO₃⁻), carbonate (CO₃²⁻), and sulfate (SO₄²⁻). These pop up everywhere in chemistry!
Writing formulae is like solving a puzzle - the positive and negative charges must balance to zero. Use subscripts to show how many of each ion you need, and brackets around polyatomic ions when there's more than one.
💡 Memory Aid: Think "positive + negative = zero" - the charges must always cancel out completely in ionic compounds.
4
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Chemical Equations
Chemical equations are chemistry's shorthand for describing reactions. Most elements appear as single atoms (like Na or Mg), but some exist as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂. Remember "HOFBrINCl" to recall these pairs!
State symbols tell you the physical state: for solid, for liquid, for gas, and (aq) for dissolved in water (aqueous). These matter because the same substance can behave differently in different states.
Balancing equations follows one golden rule: you can't change the chemical formulae, only add numbers in front of them. It's like solving a mathematical puzzle where atoms on both sides must match exactly. Start with the most complex molecule and work systematically through each element.
The key is patience and practice. Check your work by counting atoms of each element on both sides - they must be equal. Use brackets carefully when dealing with polyatomic ions, remembering that subscripts outside brackets multiply everything inside.
💡 Pro Strategy: Always balance the most complicated molecule first, then work through hydrogen and oxygen last - they're usually the easiest to adjust.
We thought you’d never ask...
Our AI Companion is a student-focused AI tool that offers more than just answers. Built on millions of Knowunity resources, it provides relevant information, personalised study plans, quizzes, and content directly in the chat, adapting to your individual learning journey.
You can download the app from Google Play Store and Apple App Store.
That's right! Enjoy free access to study content, connect with fellow students, and get instant help – all at your fingertips.
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