Ever wondered why some electrons are harder to remove than...
Chemistry113 views·Updated Jun 14, 2026·5 pagesUnderstanding Atomic Structure: Electronic Configurations and Ionization Energies
Aasiyah Rahman@aasiyahrahman
1 / 5
1
of 5
Atomic Structure Basics
Think of atoms like a block of flats - electrons live in different floors (shells) with specific rooms (orbitals). Electron configuration tells us exactly where each electron lives, and it's not random at all.
Shells are divided into sub-shells (s, p, d, f), and each sub-shell contains orbitals - regions where you're most likely to find an electron. Each orbital can hold a maximum of 2 electrons, but they must have opposite spins (think of them as spinning in different directions).
The Aufbau principle is your best friend here - electrons always fill the lowest energy orbitals first. Remember the capacities: s holds 2 electrons, p holds 6, and d holds 10. It's like filling up the cheapest seats at a concert before moving to the expensive ones!
Quick Tip: Draw the energy level diagram - it shows you exactly which orbitals fill in which order, and you'll use this constantly in exams.
2
of 5
Electron Configuration Rules
Writing electron configurations is like giving someone directions to find each electron. Start by counting the total electrons (same as atomic number), then fill orbitals from lowest to highest energy using the pattern: 1s, 2s, 2p, 3s, 3p, 4s, 3d...
Hund's rule is the "bus seating rule" - electrons prefer to sit alone in orbitals of equal energy before pairing up. This reduces electrostatic repulsion between negatively charged electrons, making the atom more stable.
For ions, remember that 4s electrons are lost before 3d electrons even though 4s fills first. This catches many students out in exams! When an atom loses electrons to form a positive ion, the highest energy electrons go first.
Exam Alert: Watch out for chromium - it's an exception that fills 3d⁵ 4s¹ instead of the expected pattern because half-filled d orbitals are extra stable.
3
of 5
Ionisation Energy Fundamentals
First ionisation energy is the energy needed to remove one electron from a gaseous atom - imagine trying to pull an electron away from the nucleus. Four main factors affect how difficult this is: distance from nucleus, nuclear charge, shielding, and electron repulsion.
Down a group, ionisation energy decreases because atoms get bigger and there's more shielding from inner electrons. Across a period, it generally increases because nuclear charge increases, pulling outer electrons more strongly.
However, there are two important dips across period 3. The Mg→Al dip happens because Al's outer electron is in a 3p orbital (further from nucleus than Mg's 3s). The P→S dip occurs because S has a paired electron in one p orbital, creating repulsion that makes it easier to remove.
Memory Trick: Think of shielding like sunglasses - more electron layers between the nucleus and outer electrons make the nuclear "pull" feel weaker.
4
of 5
Successive Ionisation Energies
Successive ionisation energies show the energy needed to remove each electron one by one. They always increase because you're removing electrons from an increasingly positive ion - imagine trying to take something away from someone who's gripping it tighter each time!
The pattern reveals which group an element belongs to. Group 1 elements show a huge jump after the 1st ionisation energy, Group 2 after the 2nd, and so on. This happens when you start removing electrons from the next shell down, which are much closer to the nucleus.
For aluminium, the first three ionisation energies are relatively similar (577, 1820, 2740 kJ mol⁻¹), but the fourth jumps to 11,600 kJ mol⁻¹. That's why aluminium forms Al³⁺ ions but never Al⁴⁺ - the energy cost is just too high.
Exam Strategy: Use ionisation energy data to identify unknown elements by looking for the big jump - it tells you how many outer electrons the atom has.
5
of 5
Identifying Elements from Data
Looking at successive ionisation energy data is like being a detective - the patterns tell you exactly which element you're dealing with. The key is spotting where the massive jump occurs.
For the vanadium example, the ionisation energies increase gradually through the first five (648, 1370, 2870, 4600, 6280), then jump dramatically to 12,400 for the sixth. This confirms vanadium is in Group 5, with electron configuration [Ar]3d³4s².
The huge increase after the 5th ionisation happens because you're now removing an electron from the 3p orbital, which is much closer to the nucleus and has less shielding. Transition metals can be tricky because they have both 3d and 4s electrons to lose before hitting the inner shell.
Pro Tip: Count how many "reasonable" ionisation energies there are before the massive jump - this tells you the group number and helps predict chemical behaviour.
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Chemistry113 views·Updated Jun 14, 2026·5 pagesUnderstanding Atomic Structure: Electronic Configurations and Ionization Energies
Aasiyah Rahman@aasiyahrahman
Ever wondered why some electrons are harder to remove than others? Atomic structure and electron configuration explain how electrons are arranged around the nucleus and why ionisation energies change in predictable patterns. Understanding these concepts is crucial for A-level chemistry...
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of 5
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Atomic Structure Basics
Think of atoms like a block of flats - electrons live in different floors (shells) with specific rooms (orbitals). Electron configuration tells us exactly where each electron lives, and it's not random at all.
Shells are divided into sub-shells (s, p, d, f), and each sub-shell contains orbitals - regions where you're most likely to find an electron. Each orbital can hold a maximum of 2 electrons, but they must have opposite spins (think of them as spinning in different directions).
The Aufbau principle is your best friend here - electrons always fill the lowest energy orbitals first. Remember the capacities: s holds 2 electrons, p holds 6, and d holds 10. It's like filling up the cheapest seats at a concert before moving to the expensive ones!
Quick Tip: Draw the energy level diagram - it shows you exactly which orbitals fill in which order, and you'll use this constantly in exams.
2
of 5Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Electron Configuration Rules
Writing electron configurations is like giving someone directions to find each electron. Start by counting the total electrons (same as atomic number), then fill orbitals from lowest to highest energy using the pattern: 1s, 2s, 2p, 3s, 3p, 4s, 3d...
Hund's rule is the "bus seating rule" - electrons prefer to sit alone in orbitals of equal energy before pairing up. This reduces electrostatic repulsion between negatively charged electrons, making the atom more stable.
For ions, remember that 4s electrons are lost before 3d electrons even though 4s fills first. This catches many students out in exams! When an atom loses electrons to form a positive ion, the highest energy electrons go first.
Exam Alert: Watch out for chromium - it's an exception that fills 3d⁵ 4s¹ instead of the expected pattern because half-filled d orbitals are extra stable.
3
of 5Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
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Ionisation Energy Fundamentals
First ionisation energy is the energy needed to remove one electron from a gaseous atom - imagine trying to pull an electron away from the nucleus. Four main factors affect how difficult this is: distance from nucleus, nuclear charge, shielding, and electron repulsion.
Down a group, ionisation energy decreases because atoms get bigger and there's more shielding from inner electrons. Across a period, it generally increases because nuclear charge increases, pulling outer electrons more strongly.
However, there are two important dips across period 3. The Mg→Al dip happens because Al's outer electron is in a 3p orbital (further from nucleus than Mg's 3s). The P→S dip occurs because S has a paired electron in one p orbital, creating repulsion that makes it easier to remove.
Memory Trick: Think of shielding like sunglasses - more electron layers between the nucleus and outer electrons make the nuclear "pull" feel weaker.
4
of 5Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
- Join milions of students
Successive Ionisation Energies
Successive ionisation energies show the energy needed to remove each electron one by one. They always increase because you're removing electrons from an increasingly positive ion - imagine trying to take something away from someone who's gripping it tighter each time!
The pattern reveals which group an element belongs to. Group 1 elements show a huge jump after the 1st ionisation energy, Group 2 after the 2nd, and so on. This happens when you start removing electrons from the next shell down, which are much closer to the nucleus.
For aluminium, the first three ionisation energies are relatively similar (577, 1820, 2740 kJ mol⁻¹), but the fourth jumps to 11,600 kJ mol⁻¹. That's why aluminium forms Al³⁺ ions but never Al⁴⁺ - the energy cost is just too high.
Exam Strategy: Use ionisation energy data to identify unknown elements by looking for the big jump - it tells you how many outer electrons the atom has.
5
of 5Sign up to see the content. It's free!
- Access to all documents
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Identifying Elements from Data
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The huge increase after the 5th ionisation happens because you're now removing an electron from the 3p orbital, which is much closer to the nucleus and has less shielding. Transition metals can be tricky because they have both 3d and 4s electrons to lose before hitting the inner shell.
Pro Tip: Count how many "reasonable" ionisation energies there are before the massive jump - this tells you the group number and helps predict chemical behaviour.
We thought you’d never ask...
What is the Knowunity AI companion?
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.
Where can I download the Knowunity app?
You can download the app from Google Play Store and Apple App Store.
Is Knowunity really free of charge?
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The app is very easy to use and well designed. I have found everything I was looking for so far and have been able to learn a lot from the presentations! I will definitely use the app for a class assignment! And of course it also helps a lot as an inspiration.
Stefan SiOS user
This app is really great. There are so many study notes and help [...]. My problem subject is French, for example, and the app has so many options for help. Thanks to this app, I have improved my French. I would recommend it to anyone.
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Wow, I am really amazed. I just tried the app because I've seen it advertised many times and was absolutely stunned. This app is THE HELP you want for school and above all, it offers so many things, such as workouts and fact sheets, which have been VERY helpful to me personally.
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