Every row and block of the periodic table maps directly to an electron energy level, so you can read off most of an element’s electronic structure just from where it sits on the chart. The period (row number) tells you the highest principal energy level that element’s electrons occupy, while the block it belongs to (the s, p, d, or f block) tells you which sublevel is currently being filled. Once you understand that the table was built around this filling pattern, finding energy levels becomes a matter of reading coordinates rather than memorizing long lists.
What Rows Tell You About Principal Energy Levels
The periodic table has seven rows, called periods, numbered 1 through 7 from top to bottom. The period number equals the highest principal energy level that element uses for its outermost electrons. Hydrogen and helium sit in period 1, so their electrons occupy only the first energy level (n = 1). Sodium is in period 3, meaning its outermost electron lives in the third energy level. Cesium is in period 6, so its valence electron reaches all the way out to the sixth shell.
This is why rows get longer as you go down the table. The first energy level has room for only two electrons (one sublevel), so period 1 holds just two elements. The second energy level adds a second sublevel, so period 2 holds eight. By the time you reach periods 6 and 7, a third and fourth sublevel have become available, and each row stretches to 32 elements if you count the lanthanides and actinides that are traditionally pulled out below.
What Blocks Tell You About Sublevels
Within each principal energy level, electrons fill smaller compartments called sublevels, labeled s, p, d, and f. The periodic table is divided into four blocks that correspond directly to these sublevels, and the block an element sits in tells you which sublevel its last electron entered.
- s block: The first two columns on the left (groups 1 and 2), plus helium. These elements are filling an s sublevel, which holds up to two electrons.
- p block: The six columns on the right (groups 13 through 18). These elements are filling a p sublevel, which holds up to six electrons.
- d block: The ten columns in the middle (groups 3 through 12), commonly called the transition metals. These elements are filling a d sublevel, which holds up to ten electrons.
- f block: The two rows that are usually pulled out and placed below the main table, the lanthanides and actinides. These elements are filling an f sublevel, which holds up to fourteen electrons.
So if someone asks “what sublevel is selenium filling?” you find selenium on the table, see that it sits in the p block in period 4, and know immediately that its outermost electrons are going into a 4p sublevel. No memorization required beyond knowing where the blocks are.
Reading an Element’s Configuration from Its Position
Combining the period and the block gives you the specific sublevel designation. For the s block and p block, the sublevel’s principal number matches the period number. Calcium is in period 4, s block, so its outermost sublevel is 4s. Chlorine is in period 3, p block, so its outermost sublevel is 3p.
The d block has a twist: the principal energy number for the d sublevel is one less than the period number. Iron sits in period 4 of the d block, but the sublevel being filled is 3d, not 4d. This is because the 3d sublevel only begins filling after the 4s sublevel is already occupied. Likewise, the f block elements have an f sublevel number that is two less than the period. The lanthanides in period 6 are filling the 4f sublevel, not the 6f.
You can also count across the block to figure out how many electrons are in that sublevel. Sulfur is the fourth element across in the p block of period 3, so it has four electrons in its 3p sublevel. Nickel is the eighth element across in the d block of period 4, so it has eight electrons in its 3d sublevel. This counting method works cleanly for most elements, though a handful of transition metals break the pattern.
Why the Filling Order Zigzags
If electrons simply filled sublevels from lowest to highest number, the table would be a lot tidier. In reality, the order goes 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, and so on. The 4s sublevel fills before 3d. The 5s fills before 4d. This means that a higher-numbered shell can start receiving electrons before a lower-numbered shell is complete.
This filling sequence is described by what textbooks call the Madelung rule, which states that orbitals fill in order of increasing sum of their two key quantum numbers, and when two orbitals have the same sum, the one with the lower principal number fills first.1Chemical Physics Letters. Configuration irregularities: deviations from the Madelung rule and inversion of orbital energy levels The periodic table was designed with this pattern baked in. That is exactly why the d block sits between the s and p blocks and why the f block is placed even further back. The table’s layout literally traces the zigzag filling path so that reading left to right, top to bottom gives you the correct order of sublevel filling for most elements.
A common shortcut taught in introductory courses is the diagonal rule, where you draw diagonal arrows across a grid of sublevels. But honestly, you do not need the diagonal diagram if you have the periodic table in front of you. The table itself is the diagram. Start from hydrogen and read across and down: each block you enter tells you which sublevel is filling next.
Where the Simple Pattern Breaks
The Madelung rule works for the vast majority of elements, but roughly twenty elements have electron configurations that deviate from the predicted filling order.1Chemical Physics Letters. Configuration irregularities: deviations from the Madelung rule and inversion of orbital energy levels The two most commonly cited exceptions are chromium and copper. The predicted configuration for chromium would place four electrons in the 3d sublevel and two in 4s, but experiments show it actually has five in 3d and one in 4s. Copper does something similar, ending up with a full 3d sublevel and only one 4s electron instead of the predicted arrangement.
The common textbook explanation is that half-filled and fully filled sublevels carry extra stability. There is some truth to this, but the real picture involves subtle energy trade-offs between electron repulsion within sublevels and the slightly different energies of orbitals in multi-electron atoms. For practical purposes, the exceptions cluster in the d block and f block, where sublevels are very close in energy and small nudges tip the balance. If you are working with s-block or p-block elements, the straightforward period-and-block reading of the table is almost always correct.
Palladium is perhaps the most dramatic exception. It sits in period 5 of the d block, so you would expect a configuration involving 5s electrons, but palladium has no 5s electrons at all. Its 4d sublevel instead holds all ten electrons. This kind of irregularity matters in advanced chemistry and materials science, but for general energy-level reading, it is an edge case you are unlikely to encounter unless you are working with specific transition metals.
Why Heavier Elements Get Stranger
As you move toward the bottom of the periodic table, the simple rules start to wobble more. Elements with very high atomic numbers have inner-shell electrons moving at speeds that approach a meaningful fraction of the speed of light. At those speeds, relativistic effects alter the sizes and energies of orbitals in ways that the basic filling rules do not account for. The s and p orbitals contract and drop in energy, while the d and f orbitals expand and rise in energy.2Coordination Chemistry Reviews. Relativistic effects on the chemical bonding properties of the heavier elements and their compounds
Gold is a favorite example of this phenomenon. Its distinctive yellow color, which sets it apart from the silvery appearance of most metals, arises because relativistic contraction shifts the energy gap between certain electron levels into the visible range. Mercury being liquid at room temperature is another consequence of the same effect. Early attempts to calculate energy levels using simpler models found small but meaningful differences from experimental measurements when relativistic motion was not included, highlighting that these corrections matter even for basic energy-level work on heavier elements.3NIPES – Journal of Science and Technology Research. The Bohr Model of the Atom: A Critical Evaluation of its Impact and Limitations in Modern Physics
For elements in periods 1 through 4, none of this matters much. The simple reading of the periodic table gives you reliable energy-level information. The relativistic complications become progressively more significant in periods 5, 6, and especially 7, where you find the actinides and the superheavy elements. This is one reason the chemistry of lanthanides and actinides feels so different from the chemistry of lighter elements: the energy levels are not quite where the basic model predicts they should be.
Practical Tips for Using the Table as an Energy-Level Map
If you are a student trying to write out an element’s electron configuration, here is a step-by-step approach using nothing but the table itself. Start at hydrogen in the top-left corner. Move across each row left to right, then drop to the next row, just like reading a book. Each element you pass fills one more electron into the sublevel indicated by the block you are in. When you reach your target element, you have traced the entire filling order.
For example, to find the energy levels of bromine (element 35), you would trace: 1s fills with two electrons in period 1, then 2s fills with two, then 2p with six, then 3s with two, 3p with six, 4s with two, 3d with ten (since the d block in period 4 corresponds to 3d), and finally 4p with five electrons, stopping at bromine’s position five columns into the p block of period 4. That gives you the full configuration and every energy level involved.
A few things to keep in mind that trip people up. Helium sits on the far right of the table in period 1, visually placed above the noble gases in the p block, but it is actually an s-block element. Its two electrons are in 1s, not in any p sublevel. The placement is a concession to chemistry (helium behaves like a noble gas), not a guide to its energy levels. Second, the numbering for d-block sublevels being one behind the period number is not an exception or a mistake. It is built into the table’s structure because those sublevels genuinely fill in that offset order. The table is not wrong; the common expectation that the numbers should always match is wrong.
How Scientists Mapped Energy Levels in the First Place
The connection between the periodic table and energy levels was not obvious from the start. The original periodic tables, including the ones developed by Mendeleev, were organized by atomic weight and chemical behavior. It was not until Henry Moseley’s X-ray experiments, just over a century ago, that the table was reorganized around atomic number. Moseley showed that the frequencies of X-rays emitted by an element under bombardment were characteristic of that element and could be used to identify the charge on its nucleus.4Philosophical Transactions of the Royal Society A. Henry Moseley, X-ray spectroscopy and the periodic table That work anchored each element to a specific nuclear charge, which in turn determines the energy levels available to electrons.
The X-rays Moseley studied are themselves products of energy-level transitions. When a high-energy beam knocks an inner-shell electron out of an atom, an electron from a higher energy level drops down to fill the vacancy, releasing an X-ray photon in the process. The energy of that photon corresponds precisely to the gap between the two levels. By measuring those X-ray energies across the periodic table, Moseley was effectively measuring the spacing of inner energy levels element by element. His results lined up perfectly with atomic number, not atomic weight, which is why the modern table is ordered the way it is.
Later developments in electron spectroscopy made it possible to observe energy levels even more directly. Rather than watching photons come out when electrons jump between levels, researchers developed techniques that measure the energies of electrons themselves as they are ejected from atoms.5Philosophical Transactions of the Royal Society of London. Series A, Mathematical and Physical Sciences. Electron spectroscopy for chemical analysis (e.s.c.a.) This approach, developed in the mid-twentieth century, gives a direct readout of the binding energies of electrons in each sublevel, confirming the energy-level assignments the periodic table predicts and revealing finer details the table cannot show on its own.
Energy Levels Versus What the Table Cannot Tell You
The periodic table is remarkably good at encoding energy-level information, but it has limits. It tells you the order in which sublevels fill and which principal level an element’s outermost electrons occupy. It does not tell you the actual numerical energy values in electron volts, nor does it show you the fine structure within sublevels caused by electron spin and orbital interactions. For lighter elements this fine structure is small enough to ignore in most practical contexts, but for heavier elements, spin-orbit coupling splits sublevels into components with measurably different energies.2Coordination Chemistry Reviews. Relativistic effects on the chemical bonding properties of the heavier elements and their compounds
The table also does not capture how energy levels shift when atoms bond to other atoms. In a molecule or a solid, the energy levels of individual atoms broaden, overlap, and rearrange depending on the bonding environment. This is why a chemist looking at bonding in a complex molecule or a physicist studying semiconductor behavior will go well beyond what the periodic table shows. But for the foundational question of what energy levels a free atom has and how they fill, the table remains the most efficient reference tool available. Knowing how to read its rows, blocks, and the offset numbering for d and f sublevels gives you most of what you need without consulting any other source.
One misconception worth clearing up: the periodic table’s shape is not arbitrary, and the placement of blocks is not just a filing system for convenience. The width of each block (2 columns for s, 6 for p, 10 for d, 14 for f) directly reflects the maximum number of electrons each sublevel can hold, which in turn comes from the physics of angular momentum and how many distinct orientations each type of orbital can have in space. The table’s geometry is a physical consequence, not a design choice. Once you see the blocks as a direct expression of sublevel capacity, the entire structure clicks into place as a map of energy levels drawn to scale in terms of electron count, if not in terms of actual energy spacing.