Patterns and Trends in the Periodic Table¶
Syllabus mapping
CH-11-01 — Properties and Structure of Matter. Covers: the difference between periods, groups and blocks; working out the charge of an ion from its electron configuration and atomic number; a secondary-source investigation to collect, tabulate and graph data on atomic radius, ionic radius, first ionisation energy and electronegativity; using that data to explain the trends; predicting the features of an atom from its position in the periodic table.
Wording above is paraphrased for teaching use — check the official syllabus before treating any line as verbatim NESA text.
You need to know: what periods, groups and blocks are, how to work out the charge an atom's ion will have, the four trends (atomic radius, ionic radius, first ionisation energy, electronegativity) across a period and down a group, and why each trend happens.
Assumed knowledge check¶
Before this page makes sense, you should be comfortable with:
- Writing electron configurations and blocks of the periodic table — from the previous page
- Plotting a line graph from a table of data — see Representing Data
Core¶
Periods, groups and blocks¶
- A period is a row. Elements in the same period have their outer electrons in the same energy level. Period 3 elements (sodium to argon) all have their outer electrons in \(n = 3\).
- A group is a column. Elements in the same group have the same number of outer (valence) electrons, arranged the same way — so they react in similar ways. Every group 1 element has one outer electron (… s¹).
- A block (s, p, d, f) tells you which type of sublevel the last electron goes into.
Some groups have names worth knowing:
| Group | Name | Outer electrons | Typical behaviour |
|---|---|---|---|
| 1 | Alkali metals | s¹ | Very reactive metals, form +1 ions |
| 2 | Alkaline earth metals | s² | Reactive metals, form +2 ions |
| 17 | Halogens | s² p⁵ | Reactive non-metals, form −1 ions |
| 18 | Noble gases | s² p⁶ (helium: s²) | Very unreactive — full outer energy level |
Working out an ion's charge¶
Atoms tend to gain or lose electrons until they have the same electron configuration as the nearest noble gas. Metals (left side) lose electrons; non-metals (right side) gain them.
Worked example
What ion will (a) magnesium (\(Z = 12\)) and (b) nitrogen (\(Z = 7\)) form?
(a) Magnesium: 1s² 2s² 2p⁶ 3s². The quickest way to a noble gas configuration (neon, 1s² 2s² 2p⁶) is to lose the 2 outer electrons.
Charge \(=\) protons − electrons \(= 12 - 10 = +2\), so \(\ce{Mg^{2+}}\).
(b) Nitrogen: 1s² 2s² 2p³. Losing 5 electrons is much harder than gaining 3 to reach neon's configuration.
Charge \(= 7 - 10 = -3\), so \(\ce{N^{3-}}\).
Shortcut by group: group 1 → +1, group 2 → +2, group 13 → +3, group 15 → −3, group 16 → −2, group 17 → −1. Group 14 elements usually don't form simple ions, and noble gases don't form ions at all.
The one idea behind every trend¶
Every trend on this page comes down to how strongly the nucleus pulls on the outer electrons. Two things decide that:
- How many protons are in the nucleus. More protons → stronger pull.
- How far away the outer electrons are, and how many inner electrons are in the way. Inner electrons shield (block) the outer electrons from the nucleus. More energy levels → outer electrons further away and more shielded → weaker pull.
Across a period (left to right): the number of protons goes up, but the outer electrons stay in the same energy level, with the same inner electrons shielding them. So the pull on the outer electrons gets stronger.
Down a group: each element has one more energy level than the one above. The outer electrons are further away and more shielded, so the pull gets weaker — even though there are more protons.
The four trends¶
| Property | What it means | Across a period → | Down a group ↓ |
|---|---|---|---|
| Atomic radius | The size of the atom | Decreases — stronger pull draws electrons in | Increases — more energy levels |
| First ionisation energy | The energy needed to remove one electron from each atom in a mole of gaseous atoms | Increases — stronger pull, harder to remove | Decreases — outer electron further away, easier to remove |
| Electronegativity | How strongly an atom attracts the electrons in a chemical bond | Increases | Decreases |
| Ionic radius | The size of an ion | Positive ions are smaller than their atoms; negative ions are larger than their atoms | Increases — more energy levels |
A shortcut that covers three of the four: the top-right of the periodic table (ignoring noble gases) has the smallest atoms, the highest ionisation energies and the highest electronegativity. Fluorine is the most electronegative element. The bottom-left is the opposite.
🖼️ Diagram needed — outline periodic table with four labelled arrows showing the direction each property increases (atomic radius: down and to the left; ionisation energy and electronegativity: up and to the right). Same outline style as the blocks diagram on the previous page. TODO — placeholder until sourced/drawn.
Why ions are a different size from their atoms:
- A positive ion has lost electrons — often its whole outer energy level. The same number of protons now pulls on fewer electrons, so it's smaller. (\(\ce{Na}\) 186 pm → \(\ce{Na+}\) 102 pm)
- A negative ion has gained electrons. The extra electrons repel each other and the same number of protons can't hold them as tightly, so it's larger. (\(\ce{Cl}\) 99 pm → \(\ce{Cl-}\) 181 pm)
Secondary-source investigation: graphing the trends¶
The syllabus asks you to collect, tabulate and plot real data. Here's period 3 to start with:
| Element | Na | Mg | Al | Si | P | S | Cl | Ar |
|---|---|---|---|---|---|---|---|---|
| Atomic number | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 |
| Atomic radius (pm) | 186 | 160 | 143 | 118 | 110 | 104 | 99 | — |
| First ionisation energy (kJ mol⁻¹) | 496 | 738 | 578 | 787 | 1012 | 1000 | 1251 | 1521 |
| Electronegativity | 0.93 | 1.31 | 1.61 | 1.90 | 2.19 | 2.58 | 3.16 | — |
And down group 1:
| Element | Li | Na | K | Rb | Cs |
|---|---|---|---|---|---|
| Atomic radius (pm) | 152 | 186 | 227 | 248 | 265 |
| First ionisation energy (kJ mol⁻¹) | 520 | 496 | 419 | 403 | 376 |
| Electronegativity | 0.98 | 0.93 | 0.82 | 0.82 | 0.79 |
Values vary slightly between data sources (different sources measure radius in different ways) — always state where your data came from. Noble gases don't usually have an electronegativity value because they rarely form bonds.
Your task: plot each property against atomic number (one graph per property), describe the trend in words, and explain it using the "how strongly does the nucleus pull" idea above. Then repeat for a group 17 (halogen) data set you find yourself. See Representing Data and Trends and Relationships for how to do this well.
When you plot period 3's ionisation energies you'll notice the trend isn't perfectly smooth — there are small dips at aluminium and sulfur. That's real, not a mistake in the data. The explanation is in the Advanced box below.
Predicting features from position¶
Worked example
Potassium (group 1, period 4) and bromine (group 17, period 4) are in the same period. Predict which has (a) the larger atomic radius, (b) the higher first ionisation energy, and (c) what ion each forms.
Same period, so they have the same number of energy levels — but bromine has many more protons (35 vs 19).
(a) Bromine's stronger pull draws its electrons in, so potassium is larger.
(b) Bromine's outer electrons are held more tightly, so bromine has the higher first ionisation energy.
(c) Potassium (group 1) loses one electron → \(\ce{K+}\). Bromine (group 17) gains one electron → \(\ce{Br-}\).
Advanced
The dips in period 3 ionisation energy.
- Mg → Al drops (738 → 578). Aluminium's outer electron is the first one in a 3p sublevel, which is slightly higher in energy (and a bit more shielded by the 3s electrons) than magnesium's 3s electrons — so it's easier to remove.
- P → S drops (1012 → 1000). Phosphorus has three 3p electrons, one in each 3p orbital. Sulfur's fourth 3p electron has to pair up in an orbital that already has an electron. The two electrons repel each other, making that one slightly easier to remove.
Core charge. A simple way to put a number on "how strongly the nucleus pulls": core charge = protons − inner (non-valence) electrons. For sodium it's \(11 - 10 = +1\); for chlorine it's \(17 - 10 = +7\). Same shielding, much bigger pull — which is exactly why chlorine is smaller and more electronegative than sodium.
Isoelectronic ions. \(\ce{O^{2-}}\), \(\ce{F-}\), \(\ce{Na+}\) and \(\ce{Mg^{2+}}\) all have 10 electrons (the same configuration as neon). The only difference is the number of protons (8, 9, 11, 12) — so the size decreases in that order, because more protons pull the same 10 electrons in more tightly.
Successive ionisation energies. Removing electrons one after another, each one takes more energy than the last. There's a huge jump when you start removing electrons from a full inner energy level. For magnesium, the 1st and 2nd ionisation energies are fairly small but the 3rd is about five times bigger — which tells you magnesium has 2 outer electrons and is in group 2.
Extension
Beyond the syllabus for this course — won't appear in the HSC, included for interest / depth study inspiration.
Do noble gases really never bond? For a long time they were called the "inert gases". In 1962, Neil Bartlett made the first noble gas compound by reacting xenon with platinum hexafluoride, a powerful fluorine compound; xenon fluorides followed soon after. Xenon is big enough (and its outer electrons far enough from the nucleus) to be pulled into a bond by fluorine, the most electronegative element — exactly what the trends on this page would predict for the heaviest noble gases. Helium, neon and argon still have no stable compounds under normal conditions.
📎 Depth study idea
Mendeleev arranged his 1869 periodic table by patterns in properties — and left gaps for elements that hadn't been discovered yet, predicting their properties from their neighbours. A depth study could take one of his predictions (e.g. "eka-silicon", later discovered as germanium) and compare his predicted values with the real measured ones, or test how well the trends on this page predict the properties of an element you look up only after making your prediction.
Video/visual resources¶
- 🎥 Khan Academy — TODO: source a periodic trends explainer that explains trends using nuclear charge and shielding (not just "memorise the arrows").
- 🖥️ TODO: source an interactive periodic table that can colour-code elements by atomic radius, ionisation energy and electronegativity (e.g. ptable.com) — confirm link.
Check yourself¶
-
What is the difference between a period and a group?
Answer
A period is a row: elements in the same period have their outer electrons in the same energy level. A group is a column: elements in the same group have the same number of outer electrons, so they react in similar ways.
-
What ion does (a) aluminium and (b) sulfur form? Show how you worked it out.
Answer
(a) Aluminium (\(Z = 13\)): 1s² 2s² 2p⁶ 3s² 3p¹. Loses its 3 outer electrons → \(\ce{Al^{3+}}\) (13 protons, 10 electrons).
(b) Sulfur (\(Z = 16\)): 1s² 2s² 2p⁶ 3s² 3p⁴. Gains 2 electrons to fill the 3p sublevel → \(\ce{S^{2-}}\) (16 protons, 18 electrons).
-
Explain why atomic radius increases down group 1.
Answer
Each element down the group has one more energy level. Its outer electron is further from the nucleus and more shielded by inner electrons, so it's held less tightly and the atom is bigger.
-
Explain why first ionisation energy generally increases across period 3.
Answer
Across the period the number of protons increases, but the outer electrons stay in the same energy level with the same shielding. The nucleus pulls the outer electrons more strongly, so more energy is needed to remove one.
-
Which is larger, a chlorine atom or a chloride ion? Explain.
Answer
The chloride ion. It has one extra electron but the same 17 protons. The electrons repel each other more and the nucleus can't hold them as tightly, so the ion spreads out more.