Electron Arrangement and Stability
You'll be able to write electron configurations for elements 1–20.
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.
Printable Worksheets
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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 · How can outer-shell electron arrangements help explain why neon and argon are generally unreactive under ordinary conditions? What else can affect a reaction?
● 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 arrangements help predict patterns in chemical 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
A helium balloon will not burn when brought near a candle, while hydrogen can ignite and burn rapidly in air. That flame reaction is combustion of hydrogen molecules (H2) with oxygen; it is not caused by hydrogen atoms trying to gain an electron from the flame. 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 full outer shell is a useful introductory way to describe the relatively low tendency of noble gases to form bonds; it does not mean that every atom with a full shell can never react. Examples include helium (2), neon (2,8), and argon (2,8,8).
The valence electrons—those in the outermost occupied shell—are involved in chemical bonding. In many introductory examples, atoms with 1 or 7 valence electrons can form stable ions by losing or gaining electrons, while noble gases with full outer shells are generally unreactive under ordinary conditions. This is a model for common bonding patterns, not a complete rule for predicting whether or how quickly a reaction occurs; observed reactivity also depends on the substances, conditions, energy changes, and reaction pathways.
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 leads into senior 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.
Valence-electron arrangements help explain common bonding patterns and broad trends, but electron count alone does not determine how reactive a substance is or how fast it reacts. Group 1 and 2 metals commonly lose outer electrons to form positive ions; within a group, changing atomic size and the energy needed to remove electrons help explain trends. Group 16 and 17 non-metals commonly gain or share electrons in bonding. These patterns are useful starting points, while the observed reaction also depends on the other reactants, conditions, energy changes, and possible reaction pathways.
Atoms with full outer shells (Group 18 noble gases such as helium, neon, argon, and krypton) are generally unreactive under ordinary conditions. Their electron arrangements help explain this pattern, but a full shell is not an impenetrable shield and does not make reactivity impossible in every setting.
Sodium reacts vigorously with water, while magnesium reacts much more slowly with cold water (and more readily with steam or acids). Aluminium often appears unreactive in water because a thin, protective aluminium oxide layer limits contact between the metal and water. These observations cannot be explained by counting valence electrons alone: electron arrangement helps describe ion formation, while measured reactivity also depends on factors such as energy changes, reaction conditions, and surface layers.
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 (Li, Na, K, Rb, Cs; hydrogen is a non-metal) each have 1 valence electron and are generally reactive metals. Their reactivity trend down the group is associated with how readily the outer electron is removed. Group 17 halogens (F, Cl, Br, I) have 7 valence electrons and are reactive non-metals; their actual reactivity varies across the group. Group 18 noble gases (He, Ne, Ar, Kr) have full outer shells and are generally unreactive under ordinary conditions. 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 and commonly form positive ions by losing it. The halogens in Group 17 have valence electrons and commonly gain or share electrons in bonding. Group 18 elements have a full outer shell, so these gases are generally unreactive under ordinary conditions. Electron arrangements help explain common patterns but do not alone determine the rate of a reaction. When atoms form bonds, they may gain, lose, or share electrons to reach a stable, full shell.
At the start of this lesson, you heard about the contrast between relatively unreactive neon and highly reactive fluorine. Both are gases, but their chemical behaviour differs. Their electron arrangements help explain some of this contrast; other factors also affect reactivity.
Now that you've worked through the lesson, how has your thinking about electron shells changed? Can you explain how outer-shell electron arrangements help account for the bonding patterns of noble gases and halogens? What other factors can affect whether or how quickly a reaction occurs?
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?
Model answers (click to reveal)
SAQ 1 (2 marks)
Model answer: Sodium has 11 electrons, which fill the shells from the inside out, 2 in the first shell, 8 in the second shell, and 1 in the third shell, giving the electron configuration 2,8,1. Sodium’s 2,8,1 arrangement helps explain why it commonly forms a Na+ ion by losing its outer electron and reaching a 2,8 arrangement. This bonding model does not by itself explain how quickly sodium reacts with water; reaction rate depends on the particular reactants and conditions.
SAQ 2 (3 marks)
Model answer: Sodium (atomic number 11) has the electron configuration 2,8,1, so it has 1 electron in its outer shell. Chlorine (atomic number 17) has the configuration 2,8,7, so it has 7 electrons in its outer shell. Sodium needs to lose its single valence electron to reach the stable arrangement 2,8, while chlorine needs to gain just 1 electron to fill its outer shell to 2,8,8. These needs match perfectly: sodium transfers its outer electron to chlorine. Sodium then becomes a stable Na+ ion and chlorine becomes a stable Cl- ion, and both now have full outer shells like a noble gas. The electron transfer forms oppositely charged ions that attract, producing the ionic compound sodium chloride (NaCl). The electron-shell model describes the ions and bonding; it does not alone determine the reaction rate.
SAQ 3 (3 marks)
Model answer: The octet rule states that atoms tend to gain, lose, or share electrons so that they end up with 8 electrons in their outer shell, the same stable arrangement as the noble gases such as neon and argon. One exception is the first shell, which is full with only 2 electrons rather than 8. For example, helium (2) and a stable hydrogen arrangement both achieve stability with a full first shell of 2 electrons, not 8. The octet rule is a useful introductory model: many atoms gain, lose, or share electrons in bonding to reach an arrangement like a noble gas, while hydrogen and helium follow a two-electron first-shell pattern. The model describes common bonding outcomes; it does not alone predict whether or how quickly a reaction occurs, and noble gases are generally rather than absolutely unreactive.