Atomic Structure and Electron Configuration¶
Syllabus mapping
CH-11-01 — Properties and Structure of Matter. Covers: using atomic number and mass number to model atoms; how mass and charge are spread through an atom; isotopes, including comparing proton and neutron numbers, calculating relative atomic mass, and why isotopes behave the same chemically but have different densities; Aboriginal Knowledge of "sickness Country" and its connection to radioactive material (explicit, examinable content integration point); naturally occurring radioisotopes and background radiation; alpha, beta and gamma radiation; balanced nuclear equations for decay and fission; how radiation type and half-life decide what a radioisotope is used for (including a secondary-source investigation); scientists' contributions to the model of the atom; energy levels, sublevels and orbitals; spdf electron configurations; linking electron configuration to position on the periodic table; flame tests and emission spectroscopy (including a laboratory experiment).
Wording above is paraphrased for teaching use — check the official syllabus before treating any line as verbatim NESA text.
You need to know: how to count the protons, neutrons and electrons in any atom or ion, what isotopes are and how to calculate relative atomic mass, the three types of nuclear radiation and how to balance nuclear equations, how to write an spdf electron configuration, and why heated metals give off coloured light.
This is a long page with a lot in it. Work through it one section at a time — each section builds on the one before.
Assumed knowledge check¶
Before this page makes sense, you should be comfortable with (all from Years 7–10 Science):
- Atoms have a nucleus (protons and neutrons) surrounded by electrons
- Writing electron arrangements in shells, e.g. sodium is 2, 8, 1
- Finding an element on the periodic table, and the basics of nuclear radiation and half-life
Core¶
What's inside an atom¶
| Particle | Where it is | Charge | Relative mass |
|---|---|---|---|
| Proton | Nucleus | +1 | 1 |
| Neutron | Nucleus | 0 | 1 |
| Electron | Around the nucleus, in energy levels | −1 | about \(\tfrac{1}{1836}\) (almost nothing) |
Two big ideas come straight out of this table:
- Almost all the mass is in the nucleus, because protons and neutrons are nearly 2000 times heavier than electrons.
- The nucleus is tiny but the atom is mostly empty space. If an atom were the size of a football stadium, the nucleus would be about the size of a pea in the centre. The electrons decide how big the atom is.
A neutral atom has the same number of protons and electrons, so the charges cancel.
Reading an element's numbers¶
- Atomic number (\(Z\)) = number of protons. This is what makes an element that element — every carbon atom has 6 protons.
- Mass number (\(A\)) = protons + neutrons.
- So, neutrons = \(A - Z\).
These are written as \(\ce{^{A}_{Z}X}\), where X is the element symbol.
Worked example
How many protons, neutrons and electrons are in (a) a sodium atom, \(\ce{^{23}_{11}Na}\), and (b) a sodium ion, \(\ce{Na+}\)?
(a) Protons \(= Z = 11\). Neutrons \(= A - Z = 23 - 11 = 12\). Electrons \(=\) protons \(= 11\) (it's neutral).
(b) The \(+\) means it has lost one electron. Protons and neutrons don't change: 11 protons, 12 neutrons, \(11 - 1 = 10\) electrons.
Ions: a positive ion (cation) has lost electrons; a negative ion (anion) has gained electrons. For an ion, electrons = protons − charge. So \(\ce{O^{2-}}\) (8 protons) has \(8 - (-2) = 10\) electrons.
Isotopes¶
Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons, so they have different mass numbers.
| Isotope | Protons | Neutrons |
|---|---|---|
| Carbon-12, \(\ce{^{12}_{6}C}\) | 6 | 6 |
| Carbon-13, \(\ce{^{13}_{6}C}\) | 6 | 7 |
| Carbon-14, \(\ce{^{14}_{6}C}\) | 6 | 8 |
Isotopes have the same chemical properties — they react the same way — because chemical reactions involve electrons, and isotopes have the same number of electrons.
Isotopes have different densities, because the extra neutrons add mass without changing the size of the atom much. For example, "heavy water" (made with the hydrogen isotope deuterium, \(\ce{^{2}_{1}H}\)) is about 11% denser than normal water — an ice cube of heavy water sinks in normal water.
Relative atomic mass¶
Most elements are a mix of isotopes. The relative atomic mass on the periodic table is the weighted average mass of all the isotopes, based on how common each one is. That's why it's usually not a whole number.
Worked example
Chlorine is 75% chlorine-35 and 25% chlorine-37. Calculate its relative atomic mass.
This matches the 35.45 on the periodic table. The answer is closer to 35 than 37 because chlorine-35 is much more common.
Radioactivity and background radiation¶
Some isotopes have an unstable nucleus. Over time the nucleus breaks down (decays) and gives off radiation. These isotopes are called radioisotopes.
Some radioisotopes occur naturally — for example uranium-238 and thorium-232 in rocks, radon-222 gas that seeps out of the ground, potassium-40 in soil and food (and in your body), and carbon-14 in all living things. Together with radiation from space, these make up background radiation: the low level of radiation that's around us all the time.
Aboriginal and Torres Strait Islander content — placeholder, not yet written
NESA lists relating Aboriginal Knowledge of "sickness Country" to the presence of radioactive material as explicit, examinable syllabus content, not an optional cultural add-on. This Knowledge belongs to specific Peoples and places, and needs to be taught from proper sources — ideally with input from your school's Aboriginal Education team, local Community, or a community-verified source — rather than a generic paraphrase.
Deliberately left unwritten here rather than filled with unverified content.
Alpha, beta and gamma radiation¶
| Alpha (\(\alpha\)) | Beta (\(\beta\)) | Gamma (\(\gamma\)) | |
|---|---|---|---|
| What it is | 2 protons + 2 neutrons (a helium nucleus) | A fast-moving electron, made when a neutron turns into a proton | High-energy electromagnetic radiation (not a particle) |
| Symbol in equations | \(\ce{^{4}_{2}He}\) | \(\ce{^{0}_{-1}e}\) | \(\gamma\) |
| Charge | +2 | −1 | 0 |
| Relative mass | 4 | about 0 | 0 |
| Stopped by | A sheet of paper or skin | A few millimetres of aluminium | Thick lead or concrete (reduces, never fully stops) |
| Ionising ability | Very high | Medium | Low |
A useful pattern: the heavier and more charged the radiation, the more it ionises things it hits — and so the less far it gets before using up its energy.
Balancing nuclear equations¶
In a nuclear equation, two things must balance:
- The mass numbers (top numbers) add up to the same total on both sides.
- The atomic numbers (bottom numbers) add up to the same total on both sides.
Worked example — alpha decay
Uranium-238 decays by alpha emission. Write the equation.
Check the top: \(238 = 234 + 4\) ✓ Check the bottom: \(92 = 90 + 2\) ✓
The new atomic number is 90, which is thorium on the periodic table. Alpha decay always takes 4 off the mass number and 2 off the atomic number.
Worked example — beta decay
Carbon-14 decays by beta emission. Write the equation.
Top: \(14 = 14 + 0\) ✓ Bottom: \(6 = 7 + (-1)\) ✓
Beta decay keeps the mass number the same and adds 1 to the atomic number (a neutron has turned into a proton).
Gamma radiation is often given off as well as alpha or beta. It has no mass or charge, so it doesn't change either number.
Nuclear fission is when a large nucleus is hit by a neutron and splits into two smaller nuclei, releasing more neutrons and a lot of energy:
Top: \(235 + 1 = 141 + 92 + 3(1) = 236\) ✓ Bottom: \(92 + 0 = 56 + 36 + 0 = 92\) ✓
Half-life, and choosing a radioisotope for a job¶
Half-life is the time it takes for half of the radioactive atoms in a sample to decay. After one half-life, half is left; after two, a quarter; after three, an eighth — and so on.
What a radioisotope is used for depends on two things: the type of radiation (how far it needs to travel) and the half-life (how long it needs to last).
| Radioisotope | Radiation | Half-life | Use | Why it fits |
|---|---|---|---|---|
| Technetium-99m | Gamma | About 6 hours | Medical imaging — injected, then detected outside the body | Gamma gets out of the body to the camera; short half-life means it's gone within a day or two |
| Iodine-131 | Beta (and gamma) | About 8 days | Treating thyroid cancer | The thyroid absorbs iodine; beta damages nearby cancer cells without travelling far |
| Cobalt-60 | Gamma | About 5 years | Radiotherapy; sterilising medical equipment and food | Gamma passes through packaging; lasts long enough to be practical |
| Americium-241 | Alpha | About 430 years | Smoke detectors | Alpha is easily blocked by smoke particles; long half-life means it lasts the life of the detector |
| Carbon-14 | Beta | About 5700 years | Dating once-living things (up to about 50 000 years old) | Half-life is on the right timescale for archaeology |
Secondary-source investigation. Choose one radioisotope used in medicine, industry or environmental monitoring. Find out its radiation type and half-life, and use them to justify why it's suitable for that use (and why a different isotope wouldn't be). The "Why it fits" column above shows the kind of reasoning you need. See Communicating for how to reference your sources.
How our model of the atom developed¶
| Scientist | Roughly when | What they contributed |
|---|---|---|
| John Dalton | 1803 | Matter is made of tiny, indivisible atoms; each element has its own kind of atom |
| J.J. Thomson | 1897 | Discovered the electron — so atoms can be split. Proposed the "plum pudding" model (electrons dotted through a ball of positive charge) |
| Ernest Rutherford | 1911 | Gold foil experiment: most alpha particles went straight through, a few bounced back. So the atom is mostly empty space with a tiny, dense, positive nucleus |
| Niels Bohr | 1913 | Electrons orbit the nucleus in fixed energy levels |
| Erwin Schrödinger | 1926 | Quantum model: electrons are found in regions called orbitals, not neat circular paths |
| James Chadwick | 1932 | Discovered the neutron |
The point isn't to memorise dates — it's that each model was changed when new evidence didn't fit the old one. That's how science works.
Energy levels, sublevels and orbitals¶
In Year 10 you wrote sodium as 2, 8, 1 — two electrons in the first shell, eight in the second, one in the third. Those shells are energy levels (\(n = 1, 2, 3, \dots\)). Electrons further from the nucleus have more energy.
Each energy level is split into sublevels, named s, p, d and f. Each sublevel is made of orbitals, and each orbital holds a maximum of 2 electrons.
| Sublevel | Number of orbitals | Maximum electrons |
|---|---|---|
| s | 1 | 2 |
| p | 3 | 6 |
| d | 5 | 10 |
| f | 7 | 14 |
| Energy level | Sublevels it contains | Maximum electrons |
|---|---|---|
| \(n = 1\) | 1s | 2 |
| \(n = 2\) | 2s, 2p | 8 |
| \(n = 3\) | 3s, 3p, 3d | 18 |
| \(n = 4\) | 4s, 4p, 4d, 4f | 32 |
Writing spdf electron configurations¶
Electrons fill sublevels from lowest energy to highest. The filling order is:
Watch out: 4s fills before 3d. This is the one that catches everyone.
🖼️ Diagram needed — the "diagonal rule" filling-order chart (sublevels in rows 1s / 2s 2p / 3s 3p 3d / 4s 4p 4d 4f, with diagonal arrows showing the filling order), plus a simple energy-level ladder showing 4s sitting just below 3d. TODO — placeholder until sourced/drawn.
The superscript tells you how many electrons are in that sublevel:
Worked example
Write the electron configuration of (a) sodium (\(Z = 11\)) and (b) iron (\(Z = 26\)).
(a) Sodium — 11 electrons. Fill in order until you've placed 11:
1s² (2 so far) → 2s² (4) → 2p⁶ (10) → 3s¹ (11) ✓
Sodium: 1s² 2s² 2p⁶ 3s¹. Compare with Year 10's 2, 8, 1 — it's the same thing in more detail.
(b) Iron — 26 electrons.
1s² (2) → 2s² (4) → 2p⁶ (10) → 3s² (12) → 3p⁶ (18) → 4s² (20) → 3d⁶ (26) ✓
Iron: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
Check: the superscripts always add up to the number of electrons. \(2+2+6+2+6+2+6 = 26\) ✓
Ions: add or remove electrons from the configuration. For the main-group elements:
- \(\ce{Na+}\) (10 electrons): 1s² 2s² 2p⁶ — the last electron (3s¹) is the one removed.
- \(\ce{Cl-}\) (18 electrons): 1s² 2s² 2p⁶ 3s² 3p⁶ — one electron added to chlorine's 3p⁵.
Notice both ions end up with the same configuration as a noble gas (neon and argon). That's very common, and it's why these ions are stable.
Electron configuration and the periodic table¶
The periodic table is laid out according to electron configuration:
- The period (row) number = the highest energy level that has electrons in it. Sodium (… 3s¹) is in period 3.
- The block = the type of sublevel the last electron goes into. Sodium's last electron is in an s sublevel, so it's in the s-block.
🖼️ Diagram needed — outline periodic table shaded into s-block (groups 1–2, plus helium), p-block (groups 13–18), d-block (groups 3–12) and f-block (lanthanides/actinides underneath). TODO — placeholder until sourced/drawn.
Flame tests and emission spectroscopy¶
When a metal compound is heated in a flame:
- Electrons absorb energy and jump up to a higher energy level (they become "excited").
- They can't stay there — they fall back down to lower energy levels.
- As they fall, they release the extra energy as light.
- The energy gaps are different for every element, so each element gives off its own colours.
| Metal ion | Flame colour |
|---|---|
| Lithium | Crimson red |
| Sodium | Bright yellow-orange |
| Potassium | Lilac (pale purple) |
| Calcium | Orange-red (brick red) |
| Strontium | Red |
| Barium | Pale green |
| Copper | Blue-green |
If you pass the light through a prism or a spectroscope, it splits into a set of sharp coloured lines — the element's emission spectrum. No two elements have the same pattern, so it works like a fingerprint. A flame test is the simple, by-eye version of emission spectroscopy.
🖼️ Diagram needed — (1) an energy-level diagram with an electron jumping up (absorbs energy) then falling back down (emits a photon of light); (2) emission line spectra for hydrogen, sodium and one other element, side by side against a black background. TODO — placeholder until sourced/drawn.
Real-world use: atomic emission spectroscopy is used in industry and medicine to find out which metals are in a mixture and roughly how much — for example, checking for heavy metals like lead in drinking water, testing the composition of metal alloys, and measuring sodium and potassium levels in blood samples.
Practical¶
Suggested prac: flame tests. Dip a clean nichrome wire loop (or a wooden splint soaked in the solution) into solutions of different metal salts and hold it in the edge of a hot Bunsen flame. Record the colour each one produces, then test an unknown solution and identify the metal ion.
Tips: clean the wire between samples by dipping it in dilute hydrochloric acid and heating until there's no colour left — sodium contamination gives a strong yellow that hides weaker colours. Looking at the flame through a handheld spectroscope turns this into a basic emission spectroscopy experiment.
Advanced
Two exceptions: chromium and copper. These don't follow the filling order exactly. A half-full or completely full d sublevel is extra stable, so one 4s electron moves into 3d:
- Chromium (\(Z = 24\)): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵ (not 4s² 3d⁴)
- Copper (\(Z = 29\)): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰ (not 4s² 3d⁹)
Transition metal ions lose their 4s electrons first. Even though 4s fills before 3d, it empties first too. So \(\ce{Fe^{2+}}\) is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ (both 4s electrons gone), and \(\ce{Fe^{3+}}\) is … 3d⁵.
Working backwards from relative atomic mass. If an element has two isotopes and you know the relative atomic mass, you can find the percentages. Let the percentage of the first isotope be \(x\), so the second is \(100 - x\). For boron (boron-10 and boron-11, relative atomic mass 10.81):
So boron is about 19% boron-10 and 81% boron-11.
How much is left after several half-lives? After \(n\) half-lives, the fraction remaining is \(\left(\tfrac{1}{2}\right)^n\). A \(40\ \text{mg}\) sample of iodine-131 (half-life 8 days) after 24 days (3 half-lives) has \(40 \times \tfrac{1}{8} = 5\ \text{mg}\) left.
Extension
Beyond the syllabus for this course — won't appear in the HSC, included for interest / depth study inspiration.
Why does 4s fill before 3d? In a simple picture, 3d "should" be lower in energy than 4s because it's in a lower shell. But 4s electrons spend some of their time very close to the nucleus (the s orbital "penetrates" inside the inner shells), so they feel a stronger pull than you'd expect. For potassium and calcium this makes 4s slightly lower in energy than 3d. Once there are electrons in 3d, the energies shift again — which is why transition metal ions lose their 4s electrons first. The full explanation needs quantum mechanics (the Schrödinger equation), which is where orbital shapes come from too.
📎 Depth study idea
You can build a working spectroscope from a cardboard tube and a piece of an old CD or DVD (the tiny grooves act as a diffraction grating), and photograph spectra with a phone camera. A depth study could compare the spectra of different light sources (fluorescent tubes, LEDs, sodium street lights, sunlight) and identify elements from their lines — the same technique astronomers use to work out what stars are made of.
Video/visual resources¶
- 🎥 Khan Academy — TODO: source an electron configuration explainer that covers the 4s-before-3d filling order explicitly.
- 🎥 TODO: source a short flame test video showing several metals side by side — useful for students who miss the prac, and for colours like potassium's lilac that are hard to see in a bright lab.
- 🖥️ PhET Simulation — Build an Atom — add protons, neutrons and electrons and watch the element, mass number and charge change
- 🖥️ PhET Simulation — Isotopes and Atomic Mass — build isotopes, then mix them to see how relative atomic mass is a weighted average
- 🖥️ PhET Simulation — Models of the Hydrogen Atom — compares atomic models from Dalton to Schrödinger, and shows electrons jumping between energy levels and emitting light
Check yourself¶
-
How many protons, neutrons and electrons are in \(\ce{^{27}_{13}Al^{3+}}\)?
Answer
Protons \(= 13\). Neutrons \(= 27 - 13 = 14\). Electrons \(= 13 - 3 = 10\) (it has lost three).
-
Explain why carbon-12 and carbon-14 react in exactly the same way.
Answer
Chemical reactions depend on electrons. Both isotopes have 6 protons and so 6 electrons, arranged the same way. They only differ in the number of neutrons, which are in the nucleus and don't take part in reactions.
-
Copper is 69% copper-63 and 31% copper-65. Calculate its relative atomic mass.
Answer
\(\dfrac{(63 \times 69) + (65 \times 31)}{100} = \dfrac{4347 + 2015}{100} = 63.62 \approx 63.6\)
-
Radium-226 (\(Z = 88\)) decays by alpha emission. Write the balanced nuclear equation and name the new element.
Answer
\(\ce{^{226}_{88}Ra -> ^{222}_{86}Rn + ^{4}_{2}He}\). The new element (\(Z = 86\)) is radon.
Check: \(226 = 222 + 4\) ✓ and \(88 = 86 + 2\) ✓
-
Write the electron configuration of (a) sulfur (\(Z = 16\)) and (b) calcium (\(Z = 20\)). Which block is each in?
Answer
(a) Sulfur: 1s² 2s² 2p⁶ 3s² 3p⁴ — last electron in a p sublevel, so p-block.
(b) Calcium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² — last electron in an s sublevel, so s-block.
-
Explain why a sodium compound gives a yellow-orange colour in a flame.
Answer
The heat gives sodium's electrons energy, so they jump to higher energy levels. When they fall back down, they release that energy as light. The energy gaps in sodium are the size that matches yellow-orange light, so that's the colour we see.