The Periodic Table Explained: How Chemistry Is Organized

The periodic table is the most information-dense single image in science. It hangs on the walls of chemistry classrooms worldwide, and it encodes how every known kind of atom behaves, from the hydrogen that fuels stars to synthetic elements that exist for a few milliseconds inside particle accelerators. Its shape is not a design choice. It is a map of how electrons fill the space around a nucleus, and once you can read that map you can predict the boiling point trend of a column, the charge an ion will carry, and why gold is yellow.

This article covers how the table was built, why it has the shape it does, the trends that make it predictive, and the strange places, such as element 118, where it stops being simple.

Mendeleev's 1869 gamble

In 1869, the Russian chemist Dmitri Mendeleev was writing a chemistry textbook and needed a logical way to present the 63 elements then known. He arranged them by atomic weight and noticed that elements with similar chemical behavior kept turning up at regular intervals. That recurrence is the periodicity the table is named for.

He was not the first to see a pattern. Alexandre-Emile Béguyer de Chancourtois arranged elements by atomic weight on a spiral in 1862, John Newlands compared the repeating properties to musical octaves in the 1860s, and Lothar Meyer published a table of 28 elements in 1864 and a fuller one in 1869 (history of the periodic table). What set Mendeleev apart was what he did when the pattern broke. Where an element did not fit, he left a gap and declared that an undiscovered element belonged there. Then he predicted what it would be like. For more on the person, see our profile of Dmitri Mendeleev.

Three of those predictions came true within 17 years.

Mendeleev's name	Actual element	Discovered	Predicted atomic mass	Observed atomic mass	Predicted density (g/cm³)	Observed density (g/cm³)
Eka-aluminium	Gallium	1875, Paul-Émile Lecoq de Boisbaudran	68	69.72	6.0	5.91
Eka-boron	Scandium	1879, Lars Fredrik Nilson	44	44.96	not listed	not listed
Eka-silicon	Germanium	1886	72	72.63	5.5	5.323

Values are from the list of Mendeleev's predicted elements. The predictions were close, not exact, and that is the right way to describe them. Gallium's observed density was within about 2% of the prediction, and germanium's atomic mass was within about 1% and its density within about 3%. The Swedish chemist Per Teodor Cleve recognized that Nilson's new element matched Mendeleev's eka-boron and told Mendeleev. A table built on pattern recognition had forecast the existence and rough properties of matter nobody had yet seen, which persuaded chemists that the pattern was real and not a coincidence.

From atomic mass to atomic number

Mendeleev ordered by atomic mass, and it worked most of the time. A few reversals bothered him. Tellurium (atomic mass about 127.6) comes before iodine (about 126.9) by chemical behavior even though it is slightly heavier. Putting the elements strictly in order of mass put them in the wrong columns.

The fix came from Henry Moseley, a young physicist at the University of Manchester. In 1913, he measured the X-ray spectra of many elements and found a systematic relationship between the wavelength of an element's characteristic X-rays and an integer that increased by one from element to element. In 1914 he concluded that this atomic number is the number of positive charges in the nucleus, which is the number of protons. Ordering by atomic number put tellurium (52) before iodine (53) without any special pleading, and it gave the first accurate count of how many elements remained to be found. Moseley pointed to four gaps, at 43, 61, 72, and 75, that were later filled by technetium, promethium, hafnium, and rhenium. He was killed at Gallipoli on August 10, 1915, aged 27 (Britannica and Manchester).

Why mass roughly tracks atomic number is a matter of nuclear structure. Adding a proton usually means adding about as many neutrons, so heavier atoms generally have more protons, and a mass-ordered table is nearly correct. It fails only where isotope mixes tilt the average mass, as with tellurium and iodine. The history of how physicists worked out what is inside an atom, from Dalton onward, is in the history of the atom, and Rutherford's gold-foil result that showed the nucleus exists is in The Rutherford Nuclear Atom.

Today's table is arranged by increasing atomic number. Each step to the right adds one proton to the nucleus, and with it one electron to the neutral atom, which changes how the atom behaves.

Why the table has the shape it does

The table's staircase outline looks arbitrary until you notice that it records, almost literally, the order in which electrons fill the available spaces around a nucleus. Quantum mechanics says electrons occupy shells (numbered 1, 2, 3, and so on), and within each shell there are subshells labeled s, p, d, and f. A subshell holds a fixed number of electrons: 2 for s, 6 for p, 10 for d, and 14 for f. Wolfgang Pauli's exclusion principle, which forbids two electrons in an atom from sharing the same quantum state, is the reason the counts are finite. You can read about the physicist behind it in our profile of Wolfgang Pauli.

The order of filling follows the Madelung rule: subshells fill in order of increasing n + ℓ, where n is the shell number and ℓ is 0, 1, 2, and 3 for s, p, d, and f, and ties go to the lower n (Aufbau principle). That gives the sequence 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. Each stretch of that list ending in a p subshell is one row, or period, of the table.

Period	Subshells filled	Number of elements
1	1s	2
2	2s, 2p	8
3	3s, 3p	8
4	4s, 3d, 4p	18
5	5s, 4d, 5p	18
6	6s, 4f, 5d, 6p	32
7	7s, 5f, 6d, 7p	32 (118 reached)

The period lengths, 2, 8, 8, 18, 18, 32, and 32, are simply the sums of the subshell capacities in each stretch. Nobody chose them.

Periods, groups, and valence electrons

• Periods (rows) correspond to the outermost shell being filled. Period 1 has only hydrogen and helium, which fill the 1s subshell. Period 2 fills 2s and 2p and so has eight elements.
• Groups (columns) collect elements with the same number of valence electrons, the electrons in the outermost shell. Valence electrons do almost all the work of chemistry because they are the ones involved in bonding. Carbon forms four bonds because it has four. Every alkali metal in Group 1 has one, and they all behave alike.

The wide middle block, the transition metals, spans Groups 3 to 12 and corresponds to filling d subshells. The two rows set below the main table, the lanthanides and actinides, fill the f subshells. They are set apart only because inserting them in line would make the table inconveniently wide.

Where the rule has exceptions

The Madelung rule is a good guide, not a law. Chromium would be expected to have the configuration [Ar] 3d⁴ 4s², and copper [Ar] 3d⁹ 4s². The measured ones are [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹, because a half-filled or filled d subshell is slightly lower in energy and one electron moves over from the 4s. When a transition metal forms an ion, it loses its outer s electrons before its d electrons, so iron (Fe) gives up two 4s electrons to become Fe²⁺ and one more to become Fe³⁺, which is why transition metals so often show several oxidation states.

There is also an argument about where the table's third column ends. The textbook layout puts scandium and yttrium in Group 3 and then lanthanum and actinium beneath them, but many chemists argue that lutetium and lawrencium belong there instead. IUPAC, the body that sets chemical conventions, has run a project to decide between the two arrangements (IUPAC periodic table page). Neither is wrong in terms of the physics. It is a question of which layout best expresses the electron-filling pattern.

Element groups and their properties

The table's columns reveal chemistry's family resemblances.

Group	Name	Valence electrons	Key properties	Examples
1	Alkali metals	1	Soft, low density, react violently with water; form +1 ions	Li, Na, K, Cs
2	Alkaline earth metals	2	Reactive, though less than Group 1; form +2 ions; vital in biology	Mg, Ca, Ba
3 to 12	Transition metals	varies (d-block)	Variable oxidation states, often colored compounds, good catalysts, high melting points	Fe, Cu, Au, Ti
13	Boron group	3	A mix of metal and metalloid behavior; aluminum is the most abundant metal in Earth's crust	B, Al, Ga
14	Carbon group	4	Enormous variety: a nonmetal (C), metalloids (Si, Ge), and metals (Sn, Pb)	C, Si, Ge
17	Halogens	7	One electron short of a full shell; strongly electronegative; form -1 ions	F, Cl, Br, I
18	Noble gases	8 (2 for He)	Full outer shell; nearly unreactive; monatomic gases at room temperature	He, Ne, Ar, Xe
57 to 71	Lanthanides	varies (f-block)	Very similar chemistry across the row; used in strong magnets and phosphors	La, Nd, Eu
89 to 103	Actinides	varies (f-block)	All radioactive; uranium and thorium occur naturally in quantity	U, Pu, Am

Think of a column as a family and a row as a generation. Family members share the same outer-shell pattern, so they act alike, and each generation down the page adds another layer of inner electrons that changes the size but not the personality.

The alkali metals illustrate the pattern. Drop sodium into water and it skids across the surface, fizzing and often igniting the hydrogen it releases. Drop caesium and the reaction is explosive. The violence increases down the group because each successive element has its single valence electron farther from the nucleus, shielded by more inner shells, and so it is lost more easily. The halogens sit at the other extreme. With seven valence electrons they are one electron short of a noble gas configuration, which makes them eager to take one. Fluorine has the highest electronegativity of any element, at 3.98 on the Pauling scale, so it pulls bonding electrons harder than anything else and forms compounds with nearly every other element. Our article on how atoms and molecules work covers how those bonds form.

Metals, nonmetals, and the diagonal in between

A jagged line running from boron down to astatine separates the metals on the left from the nonmetals on the right. Elements along that line, including silicon, germanium, and arsenic, are called metalloids because they have mixed properties. Silicon conducts electricity better than an insulator and worse than a metal, and a small amount of an added impurity changes its conductivity by orders of magnitude, which is the reason it underlies modern electronics.

Trends across the table

Three properties change in predictable ways across a period and down a group. Understanding them explains a large share of practical chemistry.

Atomic radius is the effective size of an atom. It increases down a group because each row adds a new shell, and it decreases from left to right across a period, because adding protons without adding a new shell pulls the whole electron cloud tighter toward the nucleus. Caesium atoms are the largest of any element you can handle in bulk, while helium atoms are the smallest.

Ionization energy is the energy needed to remove one electron from a gaseous atom, and it runs opposite to size. Small atoms with electrons close to a highly charged nucleus hold them tightly. The pattern is easy to see across Period 3.

Element	Na	Mg	Al	Si	P	S	Cl	Ar
First ionization energy (eV)	5.14	7.65	5.99	8.15	10.49	10.36	12.97	15.76

Values are from CRC Handbook data collected on the Wikipedia data page. The general rise from sodium to argon is the trend. The two dips, aluminum below magnesium and sulfur below phosphorus, are the exceptions that teach something. Aluminum's outer electron sits in a higher-energy p subshell, easier to remove than magnesium's paired s electrons, and sulfur's outer electron is one of a pair in a p orbital, where electron-electron repulsion makes it a little easier to pull off. The overall extremes are helium, at 24.59 eV, and caesium, at 3.89 eV.

Electronegativity is how strongly an atom pulls on shared electrons in a bond. It is highest at the upper right and lowest at the lower left.

Element	Cs	Na	Fe	Si	H	C	Cl	N	O	F
Pauling electronegativity	0.79	0.93	1.83	1.90	2.20	2.55	3.16	3.04	3.44	3.98

These values are collected on the electronegativity data page. The difference in electronegativity between two bonded atoms indicates the bond type. A small difference means the electrons are shared, a covalent bond. A large one means one atom takes the electrons nearly outright and the bond is ionic. Sodium (0.93) and chlorine (3.16) differ by 2.2, so table salt is ionic. Carbon (2.55) and hydrogen (2.20) differ by 0.35, so methane is covalent.

Why gold is gold: relativistic effects

Most metals are silvery. Gold is not, and the reason has nothing to do with ordinary chemistry. It requires special relativity.

Gold has 79 protons, so its nucleus carries a very large positive charge, and its innermost electrons are pulled into very tight, fast orbits. A rough estimate puts the speed of a 1s electron at about Z/137 of the speed of light, where 137 is roughly the inverse of the fine-structure constant. For gold that works out to about 58% of the speed of light (79 ÷ 137 ≈ 0.58). At that speed relativity matters: the electrons behave as if they were heavier, and their orbitals contract. The effect cascades outward to the other s orbitals, including gold's outermost 6s orbital, which shrinks and drops in energy. The 5d orbitals, which have little density near the nucleus, are pushed the other way, up in energy and outward in size.

The result is that the energy gap between gold's filled 5d states and its half-filled 6s state shrinks to the point where it absorbs blue light. A material that absorbs blue and reflects the rest looks yellow. In silver, the same transition needs ultraviolet light, which we cannot see, so visible light is reflected evenly and silver looks silvery. This account is the standard one in the literature, for example in Pekka Pyykkö's review of relativistic effects in chemistry (Annual Review of Physical Chemistry, 2012), though specialists note that the exact optical picture can also be framed in terms of band structure, with absorption setting in around 2.3 eV. The relativistic stabilization of the 6s orbital is also commonly credited with gold's low chemical reactivity, which is a large part of what makes it valuable. The economic side of that story is in our article on why gold is so expensive.

Noble gases: full shells and the exceptions

The noble gases, Group 18, were the last major family to be found. Lord Rayleigh and William Ramsay isolated argon in 1894, and Ramsay went on to find krypton, neon, and xenon in 1898 by fractional distillation of liquefied air. Mendeleev at first resisted them, and in 1902 he added them to his table as a new group, Group 0 (history of the periodic table). They fit naturally: each has a full outer shell, so there is little reason for it to gain, lose, or share electrons.

"Unreactive" turned out to mean "very reluctant," not "never." In June 1962, Neil Bartlett at the University of British Columbia reported a solid made from xenon and platinum hexafluoride, the first true noble gas compound, and xenon tetrafluoride followed within months. Later work showed that Bartlett's product was more complicated than the simple formula first proposed, but the central point stood: a noble gas had been made to bond. Argon followed in 2000, when a Finnish team made argon fluorohydride (HArF) by photolyzing hydrogen fluoride in a solid argon matrix at about -265 °C. It is not stable on warming. Heavier noble gases react more readily than lighter ones, because their outer electrons are farther from the nucleus and held less tightly, and as of today helium, neon, and argon have no compounds that survive at room temperature.

Oganesson: a noble gas that probably is not

Element 118, oganesson, sits at the bottom of the noble gas column, and by position it should be the heaviest noble gas. It almost certainly does not behave like one.

The reason is the same relativity that colors gold, only far stronger. With 118 protons, the relativistic distortion of the electron orbitals is large enough that, in calculations, the neat shell structure that makes a noble gas inert begins to wash out. The valence and core electrons of oganesson are expected to be "smeared out" into something closer to a uniform gas of electrons than the shells of radon or xenon. The predictions that follow are surprising: oganesson should be significantly more reactive than the other noble gases, probably a solid at room temperature with a melting point around 325 K and a boiling point around 450 K, and possibly a semiconductor with a band gap of about 1.5 eV (Oganesson).

None of it can be tested directly. The element was first made in 2002 at the Joint Institute for Nuclear Research (JINR) in Dubna, Russia, by a team that included scientists from Lawrence Livermore National Laboratory, who fired calcium-48 ions at a californium-249 target. In total only three, possibly four, atoms of oganesson-294 have been detected. Its half-life is about 0.7 milliseconds, and the earlier estimate from the first three decays was 0.89 ms with a large uncertainty. There is no time to do chemistry in that interval. IUPAC recognized the discovery in December 2015, and on November 28, 2016, approved the names nihonium (113), moscovium (115), tennessine (117), and oganesson (118), the last honoring Yuri Oganessian, the pioneer of superheavy element research (IUPAC). The announcement completed the seventh row.

Synthetic elements and the frontier

Every element heavier than uranium (atomic number 92) is synthetic in the sense that it does not occur in nature in meaningful quantities. Elements with atomic number 104 and above are called superheavy. They are made by firing a beam of lighter nuclei at a thin target of a heavy element, at up to about a tenth of the speed of light, in the hope that two nuclei fuse instead of bouncing apart. The fused nucleus is highly excited and usually sheds neutrons or a gamma ray within about 10⁻¹⁶ seconds. A separator then sweeps whatever survives to a detector that records where it landed and the chain of decays that follows, and the decay chain identifies the element (superheavy elements). The odds are tiny. One of the reactions that produced hassium, for instance, had a cross-section of about 20 picobarns, which means one fusion in an enormous number of collisions.

Half-lives tend to fall as atomic number rises, though not smoothly. Except for rutherfordium and dubnium, all known superheavy isotopes have half-lives of minutes or less, and seaborgium's longest-lived known isotope lasts about 14 minutes. Flerovium (114), first detected at Dubna in December 1998, has a confirmed isotope, flerovium-289, with a half-life of about 1.9 to 2.1 seconds, and a heavier isotope that has not been confirmed with a reported half-life of about 19 seconds (flerovium). Experiments suggest flerovium is surprisingly volatile for a member of the carbon group and may be a gas at standard conditions.

That is the rationale for the island of stability, the idea that nuclear shell effects should make certain combinations of protons and neutrons, around 300 nucleons in total, resist spontaneous fission far better than their neighbors. Theorists place its center in the neighborhood of elements 114 to 126, though later discoveries suggest the island may lie farther away than first thought. An isotope lasting minutes or hours would let chemists study a superheavy element's behavior for the first time. Nature builds heavy elements by a different route, inside stars and stellar explosions, which our article on nucleosynthesis describes.

Elements 119 and 120 would start an eighth period. Element 119 would sit below francium as an alkali metal, and element 120 below radium as an alkaline earth metal. Neither has been reported as synthesized as of this article's October 2026 update. In 2024, Lawrence Berkeley National Laboratory reported that a beam of titanium-50 fired at a plutonium-244 target had produced two atoms of livermorium (116), a proof of concept for using titanium instead of the usual calcium-48 beam to reach element 120. A search for element 120 with titanium-50 on a californium-249 target, which is expected to be considerably harder than the livermorium run, was planned for late 2025 (Unbinilium).

The chemistry of life

The table lists 118 elements, and life on Earth makes serious use of about a dozen.

By mass, the human body is roughly 65% oxygen, 18% carbon, 10% hydrogen, and 3% nitrogen, which together make up about 96% of it (abundance of the chemical elements). These four are all clustered in the upper right of the table, and they share the properties that make them suited to building molecules: they form strong, directional covalent bonds, they can link into chains and rings, and they are small enough to pack densely.

The rest of the body's mass is dominated by a short list.

• Calcium and phosphorus in bone, with phosphorus also in the backbone of DNA.
• Potassium and sodium, whose gradients across cell membranes carry nerve signals.
• Sulfur, which forms disulfide bridges that hold protein shapes together.
• Chlorine, which helps regulate fluid balance.
• Magnesium, which is needed by many enzymes and sits at the center of chlorophyll in plants.
• Iron, which carries oxygen in hemoglobin.
• Zinc, copper, manganese, iodine, and selenium, needed in trace amounts.

The pattern of the elements that life uses looks like the pattern of what is abundant on the surface of the Earth, which makes sense. Oxygen, silicon, aluminum, and iron are the four most common elements in the crust by mass, at about 46%, 28%, 8.3%, and 5.6%. Over the universe as a whole, hydrogen is about 74% of ordinary matter and helium about 24%, which our article on the most abundant elements in the universe explains in more detail. The tool that first let astronomers read what stars are made of is described in how the spectroscope revealed what stars are made of.

Why carbon and not silicon?

Silicon sits directly below carbon, with the same four valence electrons, and science fiction has long imagined silicon-based life. The standard chemical objection is that carbon and silicon behave very differently once oxygen is around. Carbon dioxide is a gas that dissolves in water and can be carried through a cell and in and out of an organism. Silicon dioxide, the same formula pattern, is quartz, a hard solid that does not dissolve. A metabolism that depends on a waste product that can neither leave nor reenter easily is hard to imagine. Silicon-oxygen bonds are also very strong, which makes them slow to break and rebuild. None of this proves silicon-based life is impossible under every condition, only that on a water-rich planet like ours carbon has major advantages.

Common misconceptions

"The table was organized by weight and that was a mistake." Mendeleev's ordering by atomic weight worked in nearly every case and failed in a few. The fix, ordering by atomic number, did not overturn his table. It explained why it worked and why it sometimes did not.

"Noble gases never form compounds." Xenon and krypton do, with fluorine and oxygen, and even argon has been coaxed into a compound at cryogenic temperatures. "Inert" is better read as "very unwilling."

"Gold is yellow because of its atomic structure, just like copper is reddish." The relativistic account for gold is specific to heavy atoms. Copper's color has a related but different origin, and you cannot read either from the table alone.

"The table is finished." Elements 1 through 118 are all confirmed and named, but nothing in the physics says the list stops. It is the limit of what laboratories can make today, set by how quickly superheavy nuclei decay and how rarely they form.

FAQs

Q: Why is the periodic table organized by atomic number rather than atomic mass? A: Atomic number, the number of protons, identifies an element and sets its electron configuration, which determines its chemistry. Atomic mass occasionally gives the wrong order, as with tellurium and iodine, where the heavier element belongs first by chemical behavior. Henry Moseley showed in 1913 and 1914, using X-ray spectra, that atomic number is the real ordering principle.

Q: What are valence electrons and why do they matter? A: Valence electrons are the electrons in an atom's outermost shell, and they take part in chemical bonding. Elements in the same group have the same number, which is why they behave alike. Sodium and potassium each have one valence electron, and both react vigorously with water to give hydrogen gas and a metal hydroxide.

Q: Why did Mendeleev's table work when it was ordered by mass? A: Atomic mass generally rises with atomic number, because heavier atoms usually have more protons. Ordering by mass therefore gave nearly the right sequence, and it failed only where isotope mixes pushed an average mass out of line, as with tellurium (about 127.6) and iodine (about 126.9). His real insight was to leave gaps for missing elements and predict their properties.

Q: What makes transition metals different from main-group elements? A: Transition metals fill d subshells, and the energies of those d electrons are close to those of the outer s electrons. That allows several oxidation states, such as Fe²⁺ and Fe³⁺, and it lets d electrons absorb visible light, so many transition metal compounds are colored. It also makes many of them good catalysts and allows them to bind other molecules as complexes.

Q: Can noble gases form chemical compounds? A: Yes. Neil Bartlett's xenon-platinum hexafluoride in June 1962 was the first, and xenon tetrafluoride followed that year. Argon fluorohydride, reported in 2000, forms only at very low temperatures in a frozen matrix and breaks down on warming. The heavier the noble gas, the easier it is to react. Helium and neon have no known compounds that are stable under ordinary conditions.

Q: How are synthetic elements made? A: Accelerators fire a beam of ions, often calcium-48, at a thin target of a heavy element such as americium or californium. Rarely, two nuclei fuse into a superheavy nucleus that survives for a fraction of a second to a few minutes. Detectors record its arrival and the chain of decays that follows, and that chain identifies the element. Producing a handful of atoms can take weeks of beam time.

Q: Why is carbon the basis of life rather than silicon? A: Carbon forms four strong, versatile bonds and builds chains and rings with hydrogen, nitrogen, and oxygen. Its oxide, carbon dioxide, is a gas that dissolves in water, so the element can be cycled through living systems. Silicon dioxide is solid quartz, which cannot. In a water-based biochemistry, those differences give carbon a large practical advantage.

The takeaway

The periodic table is less a list of ingredients than a compression of quantum mechanics into a grid. Electrons fill subshells in a fixed order, the order sets the row lengths, and the number of outer electrons sets the column, so a single glance at an element's position tells you whether it will lose electrons, gain them, or share them. The patterns hold well enough that Mendeleev could predict three unseen elements, and they hold well enough that chemists today predict the behavior of atoms they will probably never hold.

The places where the pattern strains are as informative as the places where it holds. Gold's color is a relativistic effect, and oganesson, the last entry, is probably not a noble gas at all. The grid still works at the bottom of the table, but the explanation shifts from simple shell counting to physics that Mendeleev could not have imagined.

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Sources and further reading

• Mendeleev, D. I. (1869). On the relationship of the properties of the elements to their atomic weights. Zeitschrift für Chemie, 12, 405-406.
• Moseley, H. G. J. (1913). The high-frequency spectra of the elements. Philosophical Magazine, 26(156), 1024-1034.
• Pyykkö, P. (2012). Relativistic effects in chemistry: More common than you thought. Annual Review of Physical Chemistry, 63, 45-64.
• Oganessian, Y. T., et al. (2006). Synthesis of the isotopes of elements 118 and 116 in the ²⁴⁹Cf and ²⁴⁵Cm + ⁴⁸Ca fusion reactions. Physical Review C, 74, 044602.
• IUPAC (2016). IUPAC announces the names of the elements 113, 115, 117, and 118, November 30, 2016.
• IUPAC. Periodic table of elements (release dated May 4, 2022).
• Khriachtchev, L., et al. (2000). A stable argon compound. Nature, 406, 874-876.
• History of the periodic table, Mendeleev's predicted elements, Aufbau principle, Oganesson, Flerovium, Superheavy element, and Abundance of the chemical elements, Wikipedia.
• National Institute of Standards and Technology (NIST). Atomic Spectra Database, authoritative source for electron configurations and energy levels.
• Emsley, J. (2011). Nature's Building Blocks: An A-Z Guide to the Elements (new edition). Oxford University Press.

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See also: What Is Dark Matter? · How Nuclear Fusion Works · How DNA Encodes Life