Mendeleev and the Periodic Table
Ordering the Elements and Predicting the Ones Not Yet Found
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The book explains how Mendeleev's work built upon earlier attempts by scientists like Lothar Meyer. It covers the history of chemistry from ancient times through the development of atomic theory. Henry Moseley's later discovery that atomic number, not atomic weight, determined elemental order revolutionized the table.
This audiobook offers a clear explanation of how one scientist's insight led to our modern understanding of matter. Anyone interested in the story behind the periodic table will find this worth their time.
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Overview
Dmitri Ivanovich Mendeleev was a Russian chemist who created the periodic table of elements and formulated the periodic law. Born on 8 February 1834, or 27 January by the old style calendar, he worked to organize the known elements and correct their properties. His system allowed him to predict the existence and characteristics of three elements not yet discovered—germanium, gallium, and scandium. He also helped found the Russian Chemical Society in 1868 and contributed to discussions on trade and agriculture. The synthetic element mendelevium is named after him, honoring his lasting impact on science.
Early life
Dmitri Mendeleev was born in Verkhnie Aremzyani, Siberia, to Ivan Pavlovich Mendeleev, a school principal, and Maria Dmitrievna Mendeleeva, a descendant of Tobolsk merchants who founded the first Siberian printing house. His father’s family traced roots to Tver, and his grandfather, Pavel Maximovich Sokolov, was a Russian Orthodox priest who gave his children new surnames while attending theological seminary, with Ivan receiving the name Mendeleev after a local landlord. Maria came from a well-known merchant family; some later claims about her ancestry, including connections to ethnic minorities, are dismissed by biographers as unverified legends. Raised in the Orthodox faith, Mendeleev's mother encouraged him to search for truth both divine and scientific. He was the youngest of 17 siblings, with only 14 surviving baptism. After his father became blind and lost his position, and his mother’s glass factory burned down, Mendeleev attended the Gymnasium in Tobolsk. In 1849, his mother took him to Moscow to apply to the university there, but he was rejected. They went on to Saint Petersburg, where he entered the Main Pedagogical Institute in 1850. After graduating, he developed tuberculosis and moved to the Crimean Peninsula, where he worked as a science master at the 1st Simferopol Gymnasium. He returned to Saint Petersburg in 1857 with restored health. Between 1859 and 1861, he studied capillarity and the spectroscope in Heidelberg. In 1862, he published a textbook on organic chemistry, winning him the Demidov Prize from the Petersburg Academy of Sciences. On April 4, 1862, he became engaged to Feozva Nikitichna Leshcheva, and they married on April 27 at Nikolaev Engineering Institute's church in Saint Petersburg. He became a professor at the Saint Petersburg Technological Institute in 1864 and Saint Petersburg State University in 1865. In 1865, he earned his Doctor of Science degree for his dissertation "On the Combinations of Water with Alcohol," and achieved tenure in 1867, beginning to teach inorganic chemistry and transforming Saint Petersburg into an internationally recognized center for chemical research.
Periodic table
In 1863, fifty-six elements were known, with a new one discovered each year. Earlier scientists like John Newlands and Lothar Meyer had noticed patterns in elements, but Mendeleev was unaware of their work when he began classifying them. As a teacher, he wrote Principles of Chemistry, where he dreamed of a table where all elements fell into place. He presented his findings to the Russian Chemical Society on March 6, 1869, stating that elements arranged by atomic weight showed periodicity. He predicted several unknown elements, calling them ekasilicon, ekaaluminium, and ekaboron—later identified as germanium, gallium, and scandium. Mendeleev also corrected atomic weights, such as for uranium, and questioned the accepted values for tellurium and iodine. His periodic table not only organized known elements but foresaw new ones, which were later discovered, proving his system correct.
Later life
In 1876, Mendeleev met Anna Ivanovna Popova, and by 1881 he had proposed to her, threatening suicide if she refused. They married on 2 April 1882, just a month after his divorce from Leshcheva was finalized, though he remained technically a bigamist due to the Orthodox Church's seven-year rule. The controversy surrounding his divorce likely played a role in his rejection from the Russian Academy of Sciences, despite his international recognition. His daughter Lyubov from this marriage became the wife of poet Alexander Blok. Mendeleev received honors across Europe, including the Davy Medal from the Royal Society of London in 1882 and later the Copley Medal in 1905. He resigned from Saint Petersburg University in 1890, was elected a Foreign Member of the Royal Society in 1892, and became director of the Bureau of Weights and Measures in 1893, holding that post until his death. He studied petroleum's composition and helped found Russia's first oil refinery, remarking that burning it as fuel "would be akin to firing up a kitchen stove with bank notes." Mendeleev was nominated for the Nobel Prize in Chemistry three times—1905, 1906, and 1907. In 1906, the Nobel Committee recommended him, but the Swedish Academy rejected him by one vote after Svante Arrhenius and others opposed him, citing his critique of Arrhenius's dissociation theory.
Death
Dmitri Mendeleev died in 1907 at seventy-two years old in Saint Petersburg from influenza. His final words to his physician were: “Doctor, you have science, I have faith,” a line possibly attributed to Jules Verne. He was laid to rest at Literatorskie Mostki cemetery. The Russian chemist and science historian Lev Chugaev described him as a chemist of genius, first-class physicist, and prolific researcher in hydrodynamics, meteorology, geology, chemical technology, and related disciplines. In 1868, Mendeleev helped found the Russian Chemical Society and engaged with issues of protectionist trade and agriculture. He hypothesized two inert elements lighter than hydrogen, identifying one as an all-pervasive gas and the other as coronium. His studies included solutions and the expansion of liquids with heat, anticipating Thomas Andrews’ work on critical temperature. Mendeleev also introduced the metric system to Russia and invented pyrocollodion, a smokeless powder for the Navy, organizing its production in 1892. He investigated petroleum, concluding it originates deep within the earth, stating: “The capital fact to note is that petroleum was born in the depths of the earth, and it is only there that we must seek its origin.”
Activities beyond chemistry
In the 1870s, Mendeleev began publishing widely beyond chemistry, examining Russian industry, agricultural productivity, and demographic trends. He sponsored studies of the Arctic Sea, assessed chemical fertilizers, and advocated for the merchant navy. His work focused heavily on Russia's petroleum sector, where he compared it to Pennsylvania's and helped the Ministry of Finance impose temporary tariffs in 1891 to support infant industries. In 1889, he was elected an honorary member of the Manchester Literary and Philosophical Society. The next year, 1890, he resigned his professorship at St. Petersburg University after a dispute with the Ministry of Education over student treatment. By 1892, he had become director of Russia's Central Bureau of Weights and Measures, where he standardized measurement prototypes and introduced the metric system. He also engaged in debates about spiritualism, dismissing its claims as superstition and criticizing its impact on scientific thought.
Vodka myth
A popular Russian story says that Dmitri Mendeleev set the standard for vodka at 40 percent. One vodka brand even claims: "In 1894, Dmitri Mendeleev, the greatest scientist in all Russia, received the decree to set the Imperial quality standard for Russian vodka and the 'Russian Standard' was born." But the government had already set that strength in 1843, when Mendeleev was just nine years old. He did, however, take charge of the Archive of Weights and Measures in Saint Petersburg in 1892, and the following year helped turn it into a government bureau. That office worked on standardizing trade tools and weights, not on product quality. In 1865, Mendeleev wrote a doctoral dissertation titled "A Discourse on the combination of alcohol and water," but it only discussed alcohol concentrations above 70 percent. He never wrote about vodka.
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Overview
The periodic table is an organized layout of the chemical elements in rows and columns, showing how their properties repeat in a predictable pattern. It's central to chemistry and used widely in physics and other sciences. The Russian chemist Dmitri Mendeleev created the first widely accepted version in 1869, arranging elements by atomic mass and using the pattern to predict properties of unknown elements. Later, scientists discovered that atomic numbers, not atomic mass, were the key to understanding the table. By 1945, Glenn T. Seaborg's work clarified the placement of actinides, helping shape a modern version of the table. Today, we know 118 elements, filling the first seven rows, though the heaviest ones still need full chemical testing. The table continues to evolve as science advances, with new elements possibly extending beyond what we currently know.
Structure
The periodic table arranges chemical elements by atomic number, the count of protons in each atom's nucleus—Z, from the German word for "number." Starting with hydrogen at 1 and ending with oganesson at 118, the table lists elements in order of increasing atomic number. Elements are grouped into columns, or groups, based on electron arrangement, especially in the outer shell, which governs chemical behavior. Oxygen, sulfur, and selenium all have four electrons in their outermost p-subshell, giving them similar traits. Though every element has multiple isotopes—atoms with the same number of protons but differing neutrons—these are grouped together under a single element. Naturally occurring elements usually exist as mixtures of isotopes, and their atomic weights reflect this average. Of the 118 known elements, 94 occur naturally on Earth; the other 24 were made only in laboratories. Some naturally found elements, like technetium and plutonium, were first synthesized before being discovered in nature. Only 83 of the naturally occurring elements are primordial, meaning they've existed since Earth formed, while others appear only as products of decay from primordial elements. Eighty of the natural elements have at least one stable isotope, with bismuth having an almost-stable one, and thorium and uranium containing isotopes that decay over timescales comparable to Earth's age. All 24 artificial elements are radioactive.
Group names and numbers
Under the current international system, the periodic table's groups are numbered 1 through 18 from left to right, with alkali metals on the far left and noble gases on the far right, though f-block groups are not included in this numbering. Groups can also be named after their first element—for example, the "scandium group" for group 3. Earlier, groups were labeled using Roman numerals. In the United States, those numerals were followed by an "A" or "B," depending on whether the group was in the s-, p-, or d-block. The Roman numerals corresponded to the final digit of today's system: group 4 elements were known as group IVB, and group 14 as IVA. In Europe, "A" was used for groups 1 through 7, and "B" for groups 11 through 17. Groups 8, 9, and 10 were once treated as one triple-sized group called group VIII in both systems. That changed in 1988 when the International Union of Pure and Applied Chemistry introduced the new 1-18 numbering system, officially deprecating the older group names.
Presentation forms
The periodic table is often shown with the f-block elements separated and placed below the main section, which cuts the number of columns from 32 down to 18. Both versions represent the same table, though the 32-column form, sometimes called the long form, displays all elements in their correct order, while the 18-column version, or medium-long form, saves space. The choice between them is a matter of presentation, not scientific difference. When discussing group 3, for example, either layout works equally well. Most tables show element symbols, and many include extra details like atomic numbers, names, blocks, standard atomic weights, melting and boiling points, densities, and states of matter. Some also use color-coding to help distinguish different types of elements.
Electron configurations
The periodic table organizes elements based on electron arrangements that explain repeating chemical properties. Each element's electrons occupy specific energy levels called shells divided into subshells—s, p, d, and f—that contain orbitals where electrons reside. An electron can only exist in certain energy states, and no two electrons in the same atom can have identical quantum numbers due to the Pauli exclusion principle. The 1s orbital holds up to two electrons, followed by 2s and three 2p orbitals filling up to eight electrons total. The third shell has one 3s, three 3p, and five 3d orbitals accommodating up to 18 electrons. The fourth shell includes 4s, 4p, 4d, and 4f orbitals holding up to 32 electrons. Electrons fill lowest energy orbitals first, with only outermost valence electrons participating in chemical reactions while core electrons remain bound. Elements are known with up to seven shells filled, though higher shells contain orbitals not typically occupied by known elements.
Order of subshell filling
The way electrons fill atomic orbitals follows the Aufbau principle, sometimes called the Madelung or Klechkovsky rule, named after Erwin Madelung and Vsevolod Klechkovsky. This rule outlines the order in which subshells are filled, with energies overlapping in a specific sequence. Electrons occupy orbitals based on increasing n + ℓ values, and when two orbitals share the same value, the one with the lower n is filled first. The s orbitals' energy shifts slightly due to quantum effects, which helps define where periods begin. Starting with hydrogen, each element's configuration builds upon the last: helium fills the first shell, lithium adds to the second, and so on. As elements grow heavier, shells fill in predictable ways, creating groups and periods that reflect recurring chemical properties. Exceptions like chromium and copper occur because electron repulsion alters energy levels slightly, but such deviations don't change overall chemistry.
Electron configuration table
The electron configuration table displays how electrons are arranged around the nucleus of each element in its neutral, gas-phase state. Main-group elements follow a predictable pattern, but transition and inner transition elements have twenty exceptions due to energy-level competition between subshells. For the last ten elements, from 109 to 118, there's no experimental data yet, so calculated configurations are used instead. Some subshells are shown as greyed out because they are completely filled.
Period 1
Hydrogen and helium, the first elements in the periodic table, are still debated in their proper placement. Hydrogen's single electron might place it in group 1 with the alkali metals, but its ability to gain an electron like halogens or form hydrides complicates that. While most place hydrogen in group 1, some arrangements put it in group 17 or separate it entirely—though chemist Eric Scerri argues that last option undermines the periodic law. Helium, on the other hand, is almost universally placed in group 18 with the noble gases due to its inertness, despite having only two outer electrons instead of eight and being an s-block element unlike other noble gases. A proposal to move it to group 2 was rejected by IUPAC in 1988, though some recent theoretical work suggests helium might share more traits with beryllium. The debate continues, rooted in whether chemical or electronic properties should decide placement.
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Overview
The periodic table organizes chemical elements by their atomic number, electron structure, and repeating properties. In its basic form, elements are listed in order of increasing atomic number, with rows and columns aligned to group elements that share similar behaviors. For instance, all elements in column 18 are noble gases, largely unreactive. This arrangement reflects more than two centuries of scientific progress, shaped by figures like Antoine-Laurent de Lavoisier, Johann Wolfgang Döbereiner, John Newlands, Julius Lothar Meyer, Dmitri Mendeleev, Glenn T. Seaborg, and others.
Aristotle
Around 330 BCE, the Greek philosopher Aristotle suggested that all matter is made up of combinations of basic roots, an idea first proposed by the Sicilian philosopher Empedocles. The Athenian philosopher Plato later referred to these roots as elements, naming them earth, water, air, and fire. Similar beliefs appeared in other ancient cultures, like Indian philosophy, where five elements together—earth, water, fire, air, and aether, called 'pañca bhūta'—were thought to make up everything.
First classification
The history of the periodic table is also a history of discovering chemical elements. In 1661, Robert Boyle defined elements as the simplest substances matter is made of. The first recorded discovery of a new element came from Hennig Brand, a German merchant, who in 1669 or later made a glowing white substance from human urine and called it "cold fire." It wasn't until 1680 that Boyle rediscovered phosphorus and shared his findings. This helped spark debate about what truly makes a substance an element. In 1718, Étienne Geoffroy created an Affinity Table using tabular grouping and chemical behavior. Then in 1789, Antoine Lavoisier published Elementary Treatise of Chemistry, listing elements like oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc, and sulfur—though he also included "light" and "caloric." He grouped them into metals and nonmetals. By the early 1800s, John Dalton proposed a method to determine atomic weights, which many chemists adopted. William Prout noted that atomic weights seemed related to hydrogen's weight. In 1817, Johann Döbereiner began classifying elements into groups of three—called triads—where the middle element's weight was the average of the other two. Later, Leopold Gmelin expanded this idea into a list of 55 elements grouped by traits. But early efforts were limited by inaccurate atomic weights; for example, carbon and oxygen were thought to be half their actual mass. It wasn't until the late 1850s that Stanislao Cannizzaro's work helped shift scientific consensus. At the 1860 Karlsruhe Congress, chemists agreed on a revised list of elements and atomic masses, setting the stage for a new classification system to emerge just two years later.
Comprehensive formalizations
In 1862, Alexandre-Émile Béguyer de Chancourtois created the "telluric helix," a three-dimensional chart showing elements arranged by atomic weight in a spiral, revealing similar properties aligned vertically. His work, published in Comptes rendus de l'Académie des Sciences, used geological terms and didn't include a chart initially. In 1863, he added a visual representation and expanded to include ions and compounds. British chemist John Newlands presented his classification in 1864, organizing 62 elements into eight groups based on recurring physical properties, likening the pattern to musical notes, though his table had no room for undiscovered elements and was dismissed by peers. Lothar Meyer independently noticed periodic trends in 1864, plotting atomic weights against volumes and identifying peaks for the most electropositive elements; he published an early version of a periodic table in 1864 and revised it in 1868, which appeared posthumously. His 1870 paper included a chart showing physical characteristics that helped determine element placements. In 1869, Dmitri Mendeleev arranged 63 elements by atomic weight, noting recurring chemical properties across columns, and even adjusted atomic weights when necessary—like reclassifying beryllium. He distributed broadsheets of his table and, by 1871, had formulated what he called the "law of periodicity."
Priority dispute and recognition
In 1881, Mendeleev made a claim about who deserved credit for creating the periodic table. He said that the true creator of a scientific idea is someone who sees not just its philosophical side, but its real, practical value. That person presents it in a way that makes others believe. Only then does the idea, like matter itself, become lasting and unchangeable. This was Mendeleev’s view on how recognition should be given—based on clarity, impact, and understanding of the work's true significance.
Mendeleev's predictions and inability to incorporate the rare-earth metals
In 1870, Mendeleev began predicting properties of elements not yet discovered, calling them eka-boron, eka-aluminium, and eka-silicium, and he later expanded these ideas in 1871. He noted that some elements like the rare-earth metals—yttrium, cerium, lanthanum, erbium, and didymium—did not fit neatly into his periodic table, puzzling him because they didn't follow expected patterns of atomic weight and valency. Though he tried to adjust their weights and valencies, including suggesting that didymium was pentavalent, Mendeleev ultimately gave up on incorporating them fully by late 1871. His work inspired others like Bohuslav Brauner, who later helped place the rare-earth metals in the table. In addition to predicting scandium, gallium, and germanium, Mendeleev left space for eighteen undiscovered elements, though only half were eventually found. The Indian chemist Prafulla Chandra Ray, working with Thomas Henry Holland, searched for new elements in hopes of filling gaps in Mendeleev's system.
Priority of discovery
Mendeleev and Meyer both created periodic tables for their own teaching needs, but Mendeleev pushed his work far beyond the classroom, making bold predictions about unknown elements, while Meyer avoided such speculation altogether. Neither proposal was immediately accepted, and many chemists found them too abstract to be useful. Mendeleev promoted his vision actively, especially in his textbook Foundations of Chemistry, where he mocked Meyer’s reluctance to predict. Meyer, on the other hand, explicitly warned students not to rely on his table for predictions, as it appeared in his Modern Theories. While Mendeleev’s approach was confrontational and forward-looking, Meyer’s was cautious and pedagogical. Though both tables were based on atomic weights, their purposes and outcomes differed greatly. Mendeleev’s confidence even earned him criticism from contemporaries who called his boldness “remarkable,” though they remained skeptical.
Recognition of Mendeleev's table
The periodic table gained recognition for describing and organizing elements, though not everyone agreed. In 1881, Mendeleev and Meyer disputed priority in Chemical News, with articles from both men and critiques of periodicity. The next year, the Royal Society awarded the Davy Medal jointly to Mendeleev and Meyer, though the rationale didn't mention Mendeleev's predictions. His eka-aluminium, later named gallium, was discovered in 1875; eka-boron and eka-silicium became scandium and germanium in 1879 and 1886. Mendeleev advised French chemist Paul-Émile Lecoq de Boisbaudran to recheck gallium's density, initially met with skepticism but later confirmed. He also corresponded with German chemist Clemens Winkler, who credited a third scientist, Hieronymous Theodor Richter, for the idea first. Some chemists remained unconvinced, but by 1890, Mendeleev's table was widely accepted as fundamental chemical knowledge. In 1889, Mendeleev remarked at the Faraday Lecture in London that he had not expected to live long enough to see his predictions validated by the Chemical Society of Great Britain.
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Overview
In 1869, Dmitri Mendeleev created a periodic table of the chemical elements by organizing them from lightest to heaviest based on recurring properties. As he arranged the elements, he noticed gaps in the pattern and boldly predicted that unknown elements would eventually fill those spaces. He gave these yet-to-be-discovered elements names like eka-boron, eka-aluminium, eka-silicon, and eka-manganese, estimating their atomic masses to be 44, 68, 72, and 100 respectively.
Prefixes
Mendeleev named the elements he predicted using prefixes from Sanskrit: eka-, dvi-, or tri-, depending on how many spaces down from a known element each new one would fall in his table. For instance, germanium was referred to as eka-silicon until 1886, and rhenium was called dvi-manganese before 1926. Other scientists also used the eka- prefix; francium was known as eka-caesium and astatine as eka-iodine prior to their discoveries. The International Union of Pure and Applied Chemistry now assigns provisional names based on atomic number instead of position in the periodic table.
Original predictions
Mendeleev forecast the existence of four elements lighter than rare-earth metals, calling them eka-boron, eka-aluminium, eka-manganese, and eka-silicon. His predictions were strikingly accurate: scandium, gallium, technetium, and germanium were later found to occupy the places he had assigned. He wrote these names as экаборъ, экаалюминій, экамарганецъ, and экасилицій, following pre-1917 Russian spelling. Scandium oxide was isolated in 1879 by Lars Fredrick Nilson; Per Teodor Cleve recognized the match and told Mendeleev that same year. Mendeleev had guessed an atomic mass of 44 for eka-boron, while scandium weighs 44.955907. In 1871, he predicted eka-aluminium, which was discovered in 1875 by Paul Emile Lecoq de Boisbaudran. Technetium was isolated in 1937 by Carlo Perrier and Emilio Segrè, well after Mendeleev's death, though he had estimated its atomic mass as 100. Germanium was isolated in 1886 and offered the clearest confirmation of his theory among all his predictions.
Other predictions
In 1871, Mendeleev predicted an element between thorium and uranium. William Crookes isolated radioactive material from uranium in 1900 that was later found to be a mixture of 234Th and 234mPa. Protactinium-234m, named "brevium," was identified in Germany in 1913, though the name protactinium wasn't used until 1918 when protactinium-231 was discovered. With Glenn T. Seaborg's actinide concept accepted in 1945, thorium, uranium, and protactinium were grouped together as actinides, shifting protactinium's place so it no longer occupied the position of eka-tantalum under 73—instead, that spot belongs to dubnium. Mendeleev had also implied a heavier analog of titanium and zirconium in his 1869 table, but in 1871 he placed lanthanum there. That prediction was confirmed in 1923 with the discovery of hafnium. Some other predictions failed because he did not recognize the lanthanides in the sixth row. In 1902, Bohuslav Brauner reorganized the elements, placing the lanthanides in a special series and renaming some of Mendeleev's elements accordingly.
Later predictions
In 1902, after accepting helium and argon as real elements, Mendeleev placed them in Group 0 of his periodic table. Even though he had doubts about atomic theory, he speculated that an undiscovered, lighter member of this chemically inert group could help explain radioactivity. He associated such an element with coronium, a spectral line from the Sun’s corona initially misidentified at 531.68 nm before being corrected to 530.3 nm and later found to originate from iron. Mendeleev calculated the lightest of these gases as having a velocity of 2,500,000 meters per second and described it as nearly massless, moving through matter with little chemical interaction. In his 1904 booklet A Chemical Conception of the Ether, he proposed even lighter elements, treating them as components of an interstellar atmosphere generated by violent internal processes within stars.
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Overview
Chemistry has its roots in ancient times, with civilizations already using techniques that would later become fundamental to the field by 1000 BC. They discovered fire, extracted metals from ores, made pottery and glazes, fermented beer and wine, used plants for medicine and perfume, turned fat into soap, crafted glass, and created alloys like bronze. Though alchemy, a early form of chemical study, never fully explained matter or its changes, the alchemists’ experiments and record-keeping helped pave the way for modern chemistry. The history of chemistry also connects closely with thermodynamics, especially through the contributions of Willard Gibbs.
Fire
Fire was likely the first chemical reaction humans learned to control, though for thousands of years it was seen as a mysterious force that could change one substance into another—like turning wood to ash or water to steam—while giving off heat and light. It shaped early life in many ways, from basic needs like cooking and warming homes, to more complex tasks such as firing pottery and bricks, and melting metals to create tools. Fire led to the discovery of glass and helped advance the practice of purifying metals, which eventually gave rise to metallurgy.
Paint
At Blombos Cave in South Africa, archaeologists discovered a workshop dating back 100,000 years where early humans processed ochre, showing they understood basic mineral techniques. On the cave walls, paintings reveal early humans mixing animal blood with other liquids, pointing to a rudimentary grasp of chemistry. These findings suggest that even tens of thousands of years ago, people were experimenting with materials and their properties.
Early metallurgy
Humans first worked with metal around 40,000 BC, using gold that occurred naturally in places like Spanish caves. By 4600 BC, the Varna culture in Bulgaria was making tools from gold. Other metals like silver, copper, tin, and meteoric iron were also found in their native state, which meant early peoples could work them without complex smelting. One of the most prized items in ancient Egypt was a dagger made from meteoric iron around 3000 BC—so rare and special that people called it a “dagger from Heaven.” Gold was already valued in Egypt by 2900 BC, and as metalworking advanced, people began to search for ways to purify metals.
Tin, lead, and copper smelting
Certain metals like tin and lead could be extracted from their ores just by heating the rocks in a fire, a process called smelting. This early form of metallurgy first appeared around the 6th and 5th millennia BC, with evidence found in sites across Serbia, including Vinča, Majdanpek, Jarmovac, and Pločnik. The earliest copper smelting is shown at Belovode, where a copper axe from 5500 BC was discovered. Signs of early metal use also appear in places like Palmela in Portugal, Los Millares in Spain, and Stonehenge in the United Kingdom. Still, the true beginnings of this practice remain unclear, as new findings continue to emerge.
Bronze
The earliest metals used by humans were simple elements or natural alloys. By combining copper and tin, people created bronze, marking the start of the Bronze Age around 3500 BC. This advancement in metalworking allowed for stronger tools and weapons, and techniques spread across cultures. Copper and tin ores were often mixed with arsenic, and because these ores are rare, tin bronzes didn't appear in western Asia until after 3000 BC. The history of metallurgy then moved forward as armies sought better alloys for armor and weapons. In China, bronze artifacts like sword blades and arrowheads were once thought to have been protected by chromium plating, but recent studies showed the chromium came from the lacquer used on them, not from intentional coating. These same artifacts, however, remained corrosion-free for over 2,000 years due to their burial environment. Chromium wasn't used elsewhere until French chemist Louis Nicolas Vauquelin experimented with it in the late 1790s. Ancient India also made important contributions to metallurgy and alchemy.
Ferrous metallurgy
The process of turning iron ore into usable metal is harder than it is for copper or tin. Still, iron was more common and easier to find locally, which made it a better choice for tools once people figured out how to work it. The Hittites are thought to have invented iron working around 1200 BC, marking the start of the Iron Age. The secret of working iron helped the Philistines succeed. In ancient China, during the Warring States period, they developed cast iron techniques and even built blast furnaces and cupola furnaces for making weapons and armor. They also created hydraulic trip hammers and double-acting piston bellows. Ferrous metallurgy, or iron working, shaped many cultures across history, including those in the Middle East, Egypt, Nubia, Anatolia, China, India, Japan, and medieval Europe.
Classical antiquity and atomism
Ancient philosophers tried to understand why substances behave the way they do—why they have different colors, densities, or smells, why they exist as solids, liquids, or gases, and how they react when exposed to fire or water. These questions led them to propose early theories about the nature of matter. Such ideas can be found in nearly every ancient civilization, each attempting to identify a small set of basic building blocks that make up everything in nature. The classical elements—air, water, earth, and fire—appeared across cultures, from ancient Greece to India, Mayan civilizations, and China, even without direct influence between them. These thinkers also included more abstract concepts like light, thought, aether, and heaven as part of their explanations for the natural world.
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Overview
Henry Gwyn Jeffreys Moseley was an English physicist who made a key discovery about atomic structure. His work provided the first real experimental proof supporting Niels Bohr’s theory of the atom, beyond what that theory had been designed to explain for hydrogen. Moseley showed that the positive charges in an atom's nucleus match its place in the periodic table, building on earlier ideas from Ernest Rutherford and Antonius van den Broek. He developed what became known as Moseley's law using X-ray spectra. When World War I began, he left his research at Oxford to serve in the British Army. Assigned to the Gallipoli campaign in Turkey, he died there on 10 August 1915, at age 27, shot during the Battle of Gallipoli. Many believe he would have been awarded the Nobel Prize in Physics in 1916.
Education and early life
Henry G. J. Moseley, known as Harry, was born in Weymouth in 1887 to a family deeply connected to science. His father, Henry Nottidge Moseley, was a biologist and professor at Oxford, who'd been part of the Challenger Expedition. His mother, Amabel Gwyn Jeffreys, was a chess champion and daughter of a Welsh biologist. Moseley excelled at school, winning a King’s scholarship to Eton, where he took chemistry and physics prizes in 1906. He then studied at Trinity College, Oxford, earning his BA in 1910. After graduating, he became a physics demonstrator at the University of Manchester, working under Ernest Rutherford. He joined the Manchester Literary and Philosophical Society in 1911 and later shifted to research work. In 1913, he returned to Oxford, where he was given lab space but no support.
Career and research
In 1912, experimenting with beta particles, Henry Moseley showed that high potentials could be generated from radium, inventing the first atomic battery, though he couldn’t reach the 1 megavolt needed to stop the particles. The following year, he used X-ray spectroscopy and Bragg’s diffraction law to measure X-ray spectra of metals, discovering a mathematical link between wavelengths and atomic numbers—what became known as Moseley’s law. This proved atomic numbers were not arbitrary but rooted in physics, correcting earlier chemical placements like those of cobalt and nickel. His work also revealed gaps at atomic numbers 43, 61, 72, and 75—spaces later filled by technetium, promethium, hafnium, and rhenium. Moseley predicted these elements and confirmed the exact number of lanthanide elements, settling long-standing chemical confusion about rare earths.
Contribution to understanding of the atom
Before Henry Moseley and his law, atomic numbers were seen as a vague way to order elements, increasing with atomic weight but not clearly defined. Moseley proved they had a real physical meaning, showing each element’s nuclear charge was exactly one unit higher than the last. He redefined atomic numbers from a simple label into an exact sequence that made the periodic table precise. This work supported the new idea of the atom, proposed by Ernest Rutherford and Antonius van den Broek in 1911, where the atomic number equals the positive charges in the nucleus—what we now know as protons. As noted by Bohr, Moseley’s findings gave strong experimental support for this model. Simple adjustments to formulas from Rydberg and Bohr later explained Moseley’s law theoretically.
Use of X-ray spectrometer
Moseley worked with early X-ray spectrometers, learning techniques from William Henry Bragg and William Lawrence Bragg at the University of Leeds, and developing methods himself. These instruments used a glass-bulb electron tube, inside which electrons struck metallic samples, knocking inner-shell electrons out of atoms and causing X-ray emissions. The resulting X-rays were directed through an opening in the tube’s shielding, then diffracted by a salt crystal. A photographic plate recorded the angles of diffraction, allowing scientists to calculate wavelengths using Bragg's law. Some of these techniques came from visible light spectroscopy, but Moseley adapted them for X-rays, even modifying his equipment to work in a vacuum chamber when dealing with softer X-rays that couldn’t pass through air or paper.
Death and aftermath
Henry Moseley, who died at twenty-seven in August 1915 during the Battle of Gallipoli, might have revolutionized atomic science had he lived. Though only twenty-six when he published his groundbreaking "Moseley law," his work was already seen as fundamental—Niels Bohr said the great change in atomic theory came from him. Robert Millikan called his research one of the dozen most brilliant in history, and Isaac Asimov wrote that his death was among the war's costliest losses to mankind. Moseley had been nominated for the 1915 Nobel Prize in chemistry, but the prize went instead to Richard Willstätter. The 1917 physics prize went to C.G. Barkla, whose work built on Moseley's. A memorial plaque was placed at Eton, and a scholarship established by his will helped fund physicist P.M.S. Blackett, who later led the Royal Society. The Institute of Physics now honors him with its Henry Moseley Medal and Prize.
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Overview
Atomic theory is the idea that all matter is made up of tiny, indivisible particles called atoms. At first, people thought of atoms as the smallest possible pieces of matter, something too small to see. Later, chemists noticed that elements combined in simple whole-number ratios, so they redefined atoms as the basic units of chemical elements. Then physicists discovered that atoms actually have internal structure and can be broken apart. This theory is one of the most important in science, so much so that physicist Richard Feynman called it the single most productive concept in physics.
Philosophical atomism
The concept that matter consists of tiny, indivisible particles is ancient, appearing across many cultures. The term "atom" comes from the Greek "atomos," meaning "indivisible." These early thoughts were rooted in philosophy rather than science. Later, modern atomic theory emerged, not based on those old ideas.
Pre-atomic chemistry
In the late 1600s, Robert Boyle introduced the idea of a chemical element as something distinct from a compound. Then, in the late 1700s, chemistry began to advance without relying on atomic theory. Antoine Lavoisier played a key role, showing that compounds are made up of elements in fixed ratios and redefining an element as a substance that cannot be broken down further by experiment. His work disproved ancient beliefs about fire, earth, air, and water as the basic elements of matter. Lavoisier also established the law of conservation of mass, stating that matter is neither created nor destroyed during chemical reactions. Later, in 1797, Joseph Proust confirmed the law of definite proportions, which says that compounds always contain the same ratios of elements by weight, no matter their source or amount.
Dalton's chemical atomism
In 1804, John Dalton examined chemical data and observed that elements combine in small whole-number ratios, a pattern now known as the law of multiple proportions. Studying nitrogen oxides—nitrous oxide, nitrous gas, and nitric acid—he found oxygen amounts of 80, 160, and 320 grams per 140 grams of nitrogen, forming a 1:2:4 ratio. This supported his belief that elements unite by fixed, indivisible units he called atoms. Dalton shared his ideas with fellow chemist Thomas Thomson, who published the first full explanation in 1807. His own work appeared in 1808 as A New System of Chemical Philosophy, where he described atoms as ultimate particles and used terms like "compound atom" for combinations. Water, for example, he viewed as a "binary atom," HO. But his method relied on relative weights, so he gave oxygen an atomic weight of 7 based on water's composition. Later, Joseph-Louis Gay-Lussac and Amedeo Avogadro improved understanding of gases. In 1811, Avogadro suggested that equal volumes of gases at the same temperature and pressure contain equal numbers of particles. This insight made it possible to determine atomic weights more accurately—Avogadro calculated oxygen's weight as 15.074.
Opposition to atomic theory
Dalton's atomic theory sparked wide interest but faced strong resistance due to confusion over terms like "atom" and "molecule." Dalton said an atom was the smallest part of any chemical substance, not just elements, leading to debates—some scientists disliked the word "molecule," others preferred it exclusively. Jöns Jacob Berzelius used "organic atoms" for complex particles, while Jean-Baptiste Dumas distinguished between "physical atoms" and "chemical atoms." Modern definitions weren't settled until the latter half of the 19th century, especially after the Karlsruhe Congress in 1860, where Stanislao Cannizzaro helped clarify the distinction. Philosophical objections arose since atoms couldn't be seen directly—some scientists like Ernst Mach and Wilhelm Ostwald rejected atomic theory as unscientific. Others, such as Marcellin Berthelot, preferred equivalent weights to explain chemical reactions without needing atoms. These opposing views were eventually overcome by later advances, including the periodic table and discoveries about molecular structure.
Prout's hypothesis
In 1815, William Prout proposed that all atomic weights were whole-number multiples of a basic unit, which he believed was hydrogen, and suggested that matter was made up of combinations of this primitive element he called a protyle. Berzelius, who was the leading expert on atomic weight at the time, disagreed, pointing out that careful measurements showed the atomic weights didn't align with whole numbers. Because of this, Prout's hypothesis was rejected in favor of Dalton’s theory, which held that each element had a unique atomic weight and all atoms of that element were identical. Still, the idea lingered, and it would later be partly supported by Francis Aston in 1912.
Vortex theory
From the 1860s through about 1890, a theory originally suggested by William Thomson and later developed by him and J.J. Thomson described atoms as stable whirlpools in a smooth, continuous fluid. They imagined matter as rotating formations in this frictionless medium, using smoke rings as visual examples. The math behind the idea came from vortex hydrodynamics work by Hermann von Helmholtz, though he didn’t support the atomic model himself. This theory emerged at the same time as ideas about the luminiferous aether, sharing some concepts but not being the same. Though it influenced mathematical thinking—like inspiring knot theory—its own supporters eventually realized the vortices weren’t stable and couldn’t explain things like magnetism or gravity.
Isomerism
In 1827, Friedrich Wöhler discovered that silver fulminate and silver cyanate were made of the same elements in identical proportions. This puzzled scientists until Jöns Jacob Berzelius coined the term isomerism in 1830 to describe such substances. Then, in 1860, Louis Pasteur proposed that differences in molecular structure might explain why isomers had distinct properties. A major step came in 1874 when Jacobus Henricus van 't Hoff suggested carbon atoms arrange themselves in a tetrahedral fashion. Using this idea, he could predict how many isomers a compound might have. For instance, he determined pentane (C5H12) could exist in three forms—exactly the number later confirmed by experiment. Isomerism remained difficult to explain using older theories such as radical theory or the theory of types.
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Overview
A chemical element is a type of atom defined by its atomic number, which is the number of protons in its nucleus—for instance, oxygen atoms each contain 8 protons. While atoms of the same element can have different numbers of neutrons, creating isotopes, elements can also transform through nuclear reactions that change their atomic numbers. The term is also used to describe pure substances made entirely of one element, such as oxygen gas. The concept of elements evolved over time, with important progress made by chemists like Dmitri Mendeleev, who in 1869 created a system organizing the known elements by increasing atomic number into rows and columns that showed repeating properties. This arrangement allowed scientists to predict characteristics of elements not yet discovered. By November 2016, the International Union of Pure and Applied Chemistry had recognized 118 elements in total: 94 occur naturally on Earth, while the others are synthetic.
Description
A chemical element refers to a pure substance made of just one kind of atom, like hydrogen or oxygen, or to that type of atom itself when part of a compound, such as oxygen in water. Water contains oxygen and hydrogen atoms, but not the separate molecular forms of those elements. Some languages distinguish between these meanings—French uses élément chimique for the atom type and corps simple for the substance—but English hasn't widely adopted those terms. Elements can be organized by name, symbol, or properties, whether as atoms or substances. These properties include atomic number, weight, isotopes, and behaviors like electronegativity or conductivity. The periodic table arranges elements by similar chemical traits and electronic structure. Naturally occurring elements go up to atomic number 94; the rest have only been created in laboratories through nuclear reactions.
Occurrence
The lightest elements, hydrogen and helium, formed in the first 20 minutes of the universe through Big Bang nucleosynthesis, with a mass ratio of about 3:1, along with tiny amounts of lithium and beryllium. Nearly all other naturally occurring elements come from processes like nucleogenic reactions, cosmogenic methods such as cosmic ray spallation, or radioactive decay on Earth. There are now 118 known elements, with the most recent additions being tennessine in 2010 and oganesson in 2006. Of these, 94 occur naturally on Earth, while six—technetium, promethium, astatine, francium, neptunium, and plutonium—are found only in trace amounts or in space, like stars and neutron star mergers. Elements combine to form molecules; some, like hydrogen, exist as diatomic molecules, while others form compounds or mixtures. Only a few elements, including gold and iron group metals, are found uncombined in nature. Air is mostly nitrogen and oxygen, but also includes carbon dioxide, water, and inert argon.
Atomic nucleus properties
The standard model describes an atom as having a dense nucleus made up of protons and neutrons, with electrons orbiting around them. Protons carry a positive charge, while neutrons are neutral, and electrons are negatively charged. Even though the protons repel each other, the nucleus stays together thanks to the strong nuclear force, which only works over very short distances. The balance between neutrons and protons determines whether a nucleus is stable or not.
Nuclide
A nuclide is a type of atom defined by its protons, neutrons, and nuclear energy state; while most atomic nuclei contain both protons and neutrons bound together by the strong force, the ratio of neutrons to protons must increase with atomic number to maintain stability. For example, lead needs about three neutrons for every two protons. The number of protons in an atom's nucleus determines its identity as an element—carbon always has six protons, which also defines its chemical behavior. Isotopes are atoms of the same element with different numbers of neutrons; carbon, for instance, can have six, seven, or eight neutrons, making carbon-12, carbon-13, and carbon-14. Though isotopes differ in mass, they share nearly identical chemical properties because their electron configurations remain the same. It is the number of protons—not mass or weight—that identifies an element.
Stability
Most elements have radioactive isotopes, but many don't occur naturally because their half-lives are too short. Isotopes can decay through alpha, beta, or fission, and those with even numbers of protons or neutrons tend to be more stable due to pairing effects. Of the 94 naturally occurring elements, 54 have multiple stable isotopes, while only 26 are monoisotopic, mostly having odd atomic numbers except for beryllium-9. The element with the most stable isotopes is tin, at ten. Elements from 1 to 82 all have at least one stable isotope, though some—like lead—are theoretically unstable but appear observationally stable. Beyond atomic number 82, isotopes become increasingly unstable; the heaviest elements are radioactive and must be synthesized, with only five discovered in Przybylski's star. The lightest radioactive isotope is tritium, while bismuth-209 has the longest known alpha decay half-life, over 10 billion times the age of the universe. Carbon-14 and potassium-40 contribute most of the radiation exposure people receive annually.
Isotopic mass and atomic mass
The mass number of an element equals the total count of protons and neutrons in its nucleus, written as a superscript before the chemical symbol, like 238U for uranium-238. This number is always whole. In contrast, atomic mass measures the actual mass of a single atom of a specific isotope, typically given in daltons or atomic mass units. The relative atomic mass is this value divided by the atomic mass constant, making it dimensionless. For instance, chlorine-35 has an atomic mass of 34.969 Da, and chlorine-37 weighs 36.966 Da, with their relative masses close to but not exactly their mass numbers. The only isotope whose atomic mass equals a whole number is carbon-12, which is defined as exactly 12 Da. This difference between mass number and atomic mass comes from nuclear binding energy. Fusion processes in stars convert hydrogen into helium, releasing energy, while heavier elements like iron are formed through the alpha process, eventually stopping at iron-52 because further fusion absorbs energy instead of producing it. Fission releases energy as nuclei with higher binding energy form. The standard atomic weight of an element is an average of its isotopes' masses, weighted by how common each isotope occurs naturally, which often results in a fractional value like chlorine's 35.453 u.
Chemically pure and isotopically pure
In chemistry, a pure element means a substance made up of atoms with the same number of protons, like copper, which has 29 protons. A copper wire can be 99.99% chemically pure if nearly all its atoms are copper. But it’s not isotopically pure because natural copper contains two stable isotopes: 63Cu and 65Cu, with different numbers of neutrons. Gold, however, is both chemically and isotopically pure since it consists only of one isotope, 197Au. Nuclear scientists define purity differently, focusing on whether an element has just one isotope, not just the same atomic number.
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