Electron Arrangement and Stability
In 1898, William Ramsay at University College London isolated neon from air, just 18 parts per million of the atmosphere, yet it glows orange-red at 18,000 volts and never reacts with anything.
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Q1 · You've probably heard that atoms have protons, neutrons, and electrons, where do you think the electrons are, and what do you think determines how an atom behaves chemically?
Q2 · Why do you think some atoms readily react with others while atoms like neon or argon seem almost completely unreactive?
● Know
- How electrons are arranged in shells around the nucleus
- The octet rule and its exceptions
- Why noble gases are stable and unreactive
● Understand
- How valence electron count determines reactivity
- Why atoms gain, lose, or share electrons to achieve stability
- How electron configuration relates to position in the periodic table
● Can do
- Write electron configurations for elements 1–20
- Determine the number of valence electrons from a diagram
- Predict relative reactivity based on electron arrangement
If you hold a helium balloon near a candle flame, nothing happens, the gas inside is completely unaffected, but if you filled the same balloon with hydrogen and brought it near a flame, it would explode, because hydrogen needs one more electron to reach a stable arrangement and will react violently to get it. Electrons in an atom occupy energy shells (also called energy levels) surrounding the nucleus. The first shell holds a maximum of 2 electrons; the second holds up to 8; the third holds up to 8 for the first 18 elements. Electrons fill the lowest available shell first, then the next. A shell that is completely full represents a state of maximum stability, the atom has no tendency to gain, lose, or share electrons. These are the noble gas configurations: helium (2), neon (2,8), argon (2,8,8).
The valence electronsthose in the outermost (highest energy) shell, are the ones involved in chemical bonding. An atom with 1 valence electron (like sodium) desperately wants to lose it to achieve a full shell. An atom with 7 valence electrons (like chlorine) desperately wants to gain one. An atom with 8 valence electrons (like argon) has nothing to gain or lose, it is chemically inert. The filling rules are the engine that drives all of bonding theory.
Magnesium (atomic number 12) has electron configuration 2,8,2. Its two outer electrons are in the third shell. To achieve the neon configuration (2,8), magnesium loses both outer electrons, forming Mg²⁺, a stable, positively charged ion used in refractory materials that must withstand temperatures above 2000 °C.
Australian secondary students study electron shell diagrams from Year 9, consistent with the NESA science curriculum that underpins HSC Chemistry. Shell diagrams are also the foundation of every drug design tool used by CSIRO's pharmaceutical research teams to predict how new molecules will react.
The number of valence electrons is a powerful predictor of chemical behaviour. Atoms with 1–2 outer electrons (Group 1 and 2 metals: lithium, sodium, potassium, magnesium) readily lose those electrons, they are reactive metals. The fewer electrons they need to lose, the easier it is: potassium (1 outer electron) is more reactive than calcium (2 outer electrons). Atoms with 6–7 outer electrons (Group 16 and 17 non-metals: oxygen, sulfur, fluorine, chlorine) readily gain electrons, they are reactive non-metals. Fluorine (7 outer electrons, needs just 1 more) is the most reactive of all non-metals.
Atoms with full outer shells (Group 18 noble gases: helium, neon, argon, krypton) have no tendency to react at all, they are inert. This explains why neon gas lights contain a gas that never reacts with its container, despite being bombarded with electricity. The electron arrangement is what makes neon chemically stable even under high voltage, the full shell is an impenetrable shield.
Sodium has 1 valence electron and reacts violently with water. Magnesium has 2 valence electrons and reacts slowly with water but rapidly with acid. Aluminium has 3 valence electrons and reacts even more slowly, each additional valence electron to lose makes the metal less reactive with water.
Fluorine gas was first isolated in 1886 and is so reactive it attacks glass reaction vessels. Today, fluorine chemistry underpins the manufacture of Teflon (polytetrafluoroethylene, PTFE), the non-stick coating used in every Australian kitchen. CSL Limited in Melbourne also uses fluorine chemistry to sterilise pharmaceutical manufacturing equipment.
The periodic table arranges elements so that their electron configurations create predictable patterns in reactivity. Group 1 (alkali metals: H, Li, Na, K, Rb, Cs) all have 1 valence electron, they are the most reactive metals. Group 17 (halogens: F, Cl, Br, I) all have 7 valence electrons, most reactive non-metals. Group 18 (noble gases: He, Ne, Ar, Kr) all have full shells, completely inert. The period number (row number) tells you how many electron shells the atom has: Period 1 = 1 shell; Period 2 = 2 shells; Period 3 = 3 shells.
This pattern has direct engineering consequences. Engineers looking for a stable, inert gas to fill light bulbs (so the tungsten filament doesn't oxidise) choose argon, the cheapest noble gas, sitting in Group 18, Period 3. Engineers designing reactive ion thrusters for satellites choose xenon, Group 18, Period 5. In both cases, the decision flows directly from reading the reactivity information embedded in the periodic table's structure.
Sodium (Group 1, Period 3): 3 shells, 1 outer electron → very reactive metal. Chlorine (Group 17, Period 3): 3 shells, 7 outer electrons → very reactive non-metal. When they react, sodium transfers its electron to chlorine, both achieve noble gas configurations, forming the stable ionic compound NaCl (table salt).
Argon gas (Group 18) is used in MIG welding, the shield gas that prevents atmospheric oxygen from reacting with the hot weld pool. Australian steel fabricators in NSW consume millions of litres of argon per year; its inertness, predicted from its Group 18 position, is the entire reason it is chosen over cheaper nitrogen.
In the periodic table, the period number tells you the number of electron an atom has. Elements in Group 1, the alkali metals, each have just valence electron, which makes them very reactive. The halogens in Group 17 have valence electrons and are the most reactive non-metals. Group 18 elements have a full outer shell, so these gases are almost completely unreactive. When atoms react, they gain, lose, or share electrons to reach a stable, full shell.
At the start of this lesson, you heard about the dramatic contrast between neon (which glows in signs and never reacts) and fluorine (which attacks glass and explodes on contact with water). Both are gases, yet they couldn't behave more differently, all because neon's outer electron shell is full while fluorine's is not.
Now that you've worked through the lesson, how has your thinking about electron shells changed? Can you now explain in your own words why a full outer shell leads to stability, and why fluorine is so desperately reactive compared to neon?
Q1. Write the electron configuration for sodium (Na, atomic number 11) and explain why sodium is highly reactive.
Q2. Using electron configuration diagrams, explain why chlorine (atomic number 17) reacts readily with sodium (atomic number 11).
Q3. Explain the octet rule and describe one exception to this rule. Why do atoms react to achieve a full outer shell?