Understanding atomic structure is absolutely crucial for mastering chemistry -...
Chemistry155 views·Updated 19 Jul 2026·22 pagesWJEC AS-Level Chemistry: Understanding Atoms in Unit 1 Topic 2
Megan@megan_0306
1 / 10
1
of 10
Basic Atomic Structure
Every atom is made up of three fundamental particles that determine its identity and behaviour. Protons (positively charged) and neutrons (no charge) cluster together in the nucleus, whilst electrons (negatively charged) whiz around in shells outside.
The atomic number tells you how many protons an element has - this never changes and defines what element you're dealing with. The mass number is the total of protons plus neutrons, which gives you the atom's "weight".
For example, sodium (Na) has 11 protons, 12 neutrons, and 11 electrons arranged in shells as 2,8,1. The key thing to remember: in a neutral atom, protons always equal electrons, keeping the overall charge at zero.
Quick Tip: Mass number - atomic number = number of neutrons. This simple calculation will save you loads of time in exams!
2
of 10
Ions and Isotopes
When atoms gain or lose electrons, they become ions - charged particles that behave very differently from neutral atoms. Lose electrons and you get a positive cation; gain electrons and you get a negative anion.
Isotopes are atoms of the same element with different numbers of neutrons. They have identical chemical properties but different masses, which makes them incredibly useful in real-world applications.
Radioactivity occurs when there's an imbalance between protons and neutrons in the nucleus. This instability causes the atom to "decay" by releasing particles or energy to become more stable.
Remember: Isotopes = same protons, different neutrons. Ions = same protons, different electrons.
3
of 10
Types of Radioactive Decay
There are three main types of radiation you need to know, each with different properties and dangers. Alpha particles are helium nuclei (heavy and slow), beta particles are fast-moving electrons, and gamma rays are electromagnetic waves with no mass.
Alpha radiation is stopped by paper, beta by aluminium, and gamma requires thick lead for protection. Their ionising power decreases in the same order - alpha is most dangerous inside your body, gamma penetrates furthest.
Beta particles form when a neutron splits into a proton and electron, with the electron shot out at high speed. Decay equations must balance - the mass numbers and atomic numbers on both sides must add up correctly.
Exam Strategy: Always check your decay equations balance. Top numbers (mass) and bottom numbers (atomic) must equal on both sides!
4
of 10
More Decay Types and Effects
Beta decay increases the atomic number by one as a neutron becomes a proton. Positron emission does the opposite, decreasing the atomic number. Electron capture also decreases atomic number when the nucleus absorbs an inner electron.
Ionising radiation damages cells and DNA by knocking electrons off atoms, potentially causing mutations or cancer. This is why radiation safety is so important in medical and industrial applications.
Understanding these decay processes helps explain why some isotopes are useful for medical tracers whilst others are dangerous waste products. The type of radiation determines both the applications and the safety precautions needed.
Health Connection: Low ionising radiation with short half-lives makes the best medical tracers - effective but safe!
5
of 10
Practical Applications of Radiation
Radioactive tracers are incredibly useful in medicine, engineering, and research. Doctors inject safe isotopes to track blood flow or organ function, engineers find leaks in water pipes, and scientists study plant metabolism.
Carbon dating uses carbon-14 decay to determine the age of once-living materials. Living organisms maintain constant C-14 levels, but after death, the C-14 decays with a known half-life, creating a natural clock.
Half-life is the time taken for half the radioactive atoms in a sample to decay. This concept is fundamental to understanding both the usefulness and the dangers of radioactive materials.
Real World: Carbon dating can accurately date materials up to about 50,000 years old - perfect for archaeology!
6
of 10
Half-Life and Ionisation Energy
Half-life varies enormously between isotopes - from fractions of a second to billions of years. This determines whether an isotope is useful for medical procedures (short half-life) or problematic nuclear waste (long half-life).
Ionisation energy is the energy needed to remove electrons from gaseous atoms. This concept helps predict how elements will react and bond with each other.
First ionisation energy removes the outermost electron, second ionisation energy removes the next one, and so on. Each successive removal requires more energy as you're pulling electrons from increasingly positive ions.
Pattern Spotting: Massive jumps in successive ionisation energies reveal electron shell structure - really useful for identifying elements!
7
of 10
Successive Ionisation Energies
When removing electrons one by one from a single atom, each removal gets progressively harder. The first electron is easiest to remove because it's furthest from the nucleus with maximum shielding from inner electrons.
As electrons are removed, the remaining ones experience stronger attraction to the nucleus. The effective nuclear charge increases because there are fewer electrons to share the positive charge from the protons.
Inner electrons are much harder to remove because they're closer to the nucleus with minimal shielding. The dramatic energy increases reveal the shell structure of atoms.
Exam Insight: Look for huge jumps in ionisation energy data - they show you've moved to a new, inner electron shell!
8
of 10
Trends in Ionisation Energy
Down a group, ionisation energy decreases because electrons are further from the nucleus and experience more shielding. This explains why caesium is more reactive than lithium - its outer electron is much easier to remove.
Across a period, ionisation energy generally increases as nuclear charge increases whilst shielding stays roughly constant. However, there are important exceptions that reveal electron arrangements.
The dips at boron and oxygen show that p-orbital electrons are slightly easier to remove than expected due to their different shapes and electron-electron repulsion in paired orbitals.
Trend Tip: Group 1 metals get more reactive going down because ionisation energy decreases - easier to lose that outer electron!
9
of 10
The Quantum Model and Electron Arrangement
Quantum theory describes electrons as existing in specific energy levels or shells around the nucleus. These shells are numbered using principal quantum numbers .
Lower quantum numbers mean closer to the nucleus and lower energy. As you move away from the nucleus, energy levels get closer together, explaining why larger atoms can accommodate so many electrons.
Positive ions always have higher ionisation energies than their neutral atoms because there are more protons than electrons, creating a stronger effective nuclear charge that holds remaining electrons more tightly.
Key Concept: Energy levels aren't fixed orbits like planets - they're probability zones where electrons are most likely to be found!
10
of 10
Orbital Shapes and Types
Orbitals are 3D regions where you're most likely to find electrons. Each orbital holds a maximum of two electrons, and different types have characteristic shapes that affect how atoms bond.
S-orbitals are spherical and exist in all energy levels. They get larger as the shell number increases but maintain their round shape. There's only one s-orbital per shell.
P-orbitals are dumbbell-shaped and appear from the second shell onwards. They also increase in size with shell number but keep their distinctive elongated shape.
Visual Memory: Think of s-orbitals as perfect spheres and p-orbitals as dumbbells or figure-8 shapes pointing in different directions!
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Chemistry155 views·Updated 19 Jul 2026·22 pagesWJEC AS-Level Chemistry: Understanding Atoms in Unit 1 Topic 2
Megan@megan_0306
Understanding atomic structure is absolutely crucial for mastering chemistry - it's the foundation that everything else builds on. You'll explore how atoms are organised, why they behave the way they do, and how radioactive decay works in the real world.
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Basic Atomic Structure
Every atom is made up of three fundamental particles that determine its identity and behaviour. Protons (positively charged) and neutrons (no charge) cluster together in the nucleus, whilst electrons (negatively charged) whiz around in shells outside.
The atomic number tells you how many protons an element has - this never changes and defines what element you're dealing with. The mass number is the total of protons plus neutrons, which gives you the atom's "weight".
For example, sodium (Na) has 11 protons, 12 neutrons, and 11 electrons arranged in shells as 2,8,1. The key thing to remember: in a neutral atom, protons always equal electrons, keeping the overall charge at zero.
Quick Tip: Mass number - atomic number = number of neutrons. This simple calculation will save you loads of time in exams!
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Ions and Isotopes
When atoms gain or lose electrons, they become ions - charged particles that behave very differently from neutral atoms. Lose electrons and you get a positive cation; gain electrons and you get a negative anion.
Isotopes are atoms of the same element with different numbers of neutrons. They have identical chemical properties but different masses, which makes them incredibly useful in real-world applications.
Radioactivity occurs when there's an imbalance between protons and neutrons in the nucleus. This instability causes the atom to "decay" by releasing particles or energy to become more stable.
Remember: Isotopes = same protons, different neutrons. Ions = same protons, different electrons.
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Types of Radioactive Decay
There are three main types of radiation you need to know, each with different properties and dangers. Alpha particles are helium nuclei (heavy and slow), beta particles are fast-moving electrons, and gamma rays are electromagnetic waves with no mass.
Alpha radiation is stopped by paper, beta by aluminium, and gamma requires thick lead for protection. Their ionising power decreases in the same order - alpha is most dangerous inside your body, gamma penetrates furthest.
Beta particles form when a neutron splits into a proton and electron, with the electron shot out at high speed. Decay equations must balance - the mass numbers and atomic numbers on both sides must add up correctly.
Exam Strategy: Always check your decay equations balance. Top numbers (mass) and bottom numbers (atomic) must equal on both sides!
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More Decay Types and Effects
Beta decay increases the atomic number by one as a neutron becomes a proton. Positron emission does the opposite, decreasing the atomic number. Electron capture also decreases atomic number when the nucleus absorbs an inner electron.
Ionising radiation damages cells and DNA by knocking electrons off atoms, potentially causing mutations or cancer. This is why radiation safety is so important in medical and industrial applications.
Understanding these decay processes helps explain why some isotopes are useful for medical tracers whilst others are dangerous waste products. The type of radiation determines both the applications and the safety precautions needed.
Health Connection: Low ionising radiation with short half-lives makes the best medical tracers - effective but safe!
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of 10Sign up to see the content. It's free!
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Practical Applications of Radiation
Radioactive tracers are incredibly useful in medicine, engineering, and research. Doctors inject safe isotopes to track blood flow or organ function, engineers find leaks in water pipes, and scientists study plant metabolism.
Carbon dating uses carbon-14 decay to determine the age of once-living materials. Living organisms maintain constant C-14 levels, but after death, the C-14 decays with a known half-life, creating a natural clock.
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Real World: Carbon dating can accurately date materials up to about 50,000 years old - perfect for archaeology!
6
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Half-Life and Ionisation Energy
Half-life varies enormously between isotopes - from fractions of a second to billions of years. This determines whether an isotope is useful for medical procedures (short half-life) or problematic nuclear waste (long half-life).
Ionisation energy is the energy needed to remove electrons from gaseous atoms. This concept helps predict how elements will react and bond with each other.
First ionisation energy removes the outermost electron, second ionisation energy removes the next one, and so on. Each successive removal requires more energy as you're pulling electrons from increasingly positive ions.
Pattern Spotting: Massive jumps in successive ionisation energies reveal electron shell structure - really useful for identifying elements!
7
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Successive Ionisation Energies
When removing electrons one by one from a single atom, each removal gets progressively harder. The first electron is easiest to remove because it's furthest from the nucleus with maximum shielding from inner electrons.
As electrons are removed, the remaining ones experience stronger attraction to the nucleus. The effective nuclear charge increases because there are fewer electrons to share the positive charge from the protons.
Inner electrons are much harder to remove because they're closer to the nucleus with minimal shielding. The dramatic energy increases reveal the shell structure of atoms.
Exam Insight: Look for huge jumps in ionisation energy data - they show you've moved to a new, inner electron shell!
8
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Trends in Ionisation Energy
Down a group, ionisation energy decreases because electrons are further from the nucleus and experience more shielding. This explains why caesium is more reactive than lithium - its outer electron is much easier to remove.
Across a period, ionisation energy generally increases as nuclear charge increases whilst shielding stays roughly constant. However, there are important exceptions that reveal electron arrangements.
The dips at boron and oxygen show that p-orbital electrons are slightly easier to remove than expected due to their different shapes and electron-electron repulsion in paired orbitals.
Trend Tip: Group 1 metals get more reactive going down because ionisation energy decreases - easier to lose that outer electron!
9
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The Quantum Model and Electron Arrangement
Quantum theory describes electrons as existing in specific energy levels or shells around the nucleus. These shells are numbered using principal quantum numbers .
Lower quantum numbers mean closer to the nucleus and lower energy. As you move away from the nucleus, energy levels get closer together, explaining why larger atoms can accommodate so many electrons.
Positive ions always have higher ionisation energies than their neutral atoms because there are more protons than electrons, creating a stronger effective nuclear charge that holds remaining electrons more tightly.
Key Concept: Energy levels aren't fixed orbits like planets - they're probability zones where electrons are most likely to be found!
10
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Orbital Shapes and Types
Orbitals are 3D regions where you're most likely to find electrons. Each orbital holds a maximum of two electrons, and different types have characteristic shapes that affect how atoms bond.
S-orbitals are spherical and exist in all energy levels. They get larger as the shell number increases but maintain their round shape. There's only one s-orbital per shell.
P-orbitals are dumbbell-shaped and appear from the second shell onwards. They also increase in size with shell number but keep their distinctive elongated shape.
Visual Memory: Think of s-orbitals as perfect spheres and p-orbitals as dumbbells or figure-8 shapes pointing in different directions!
We thought you’d never ask...
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