essay about periodic table of elements

Breakdown of the Periodic Table Essay

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The Periodic Table is the tool for arranging elements based on the correlation between the periodic function of their atomic numbers and the properties of the elements in question (i.e., physical and chemical ones). Groups are the eight horizontal columns in the table. In every group, the atomic size and electropositivity levels drop, whereas the nuclear charge, ionization potential, and electronegativity rise, when viewed in descending order.

Although the Periodic Table is typically represented by its short form, the long one is sometimes used to demonstrate the correlations mentioned above in a more graphic manner.

The first group consists of seven elements, which are defined as alkali metals and include Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), and Francium (Fr). The electrons in the identified group are arranged on orbital 1s-7s. The elements in the group are typically labelled as metals. Although Hydrogen (H) is often attributed to the group, it is not viewed as a metal.

The second group contains the same number of elements as the first one. It incorporates (Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra). The 6d orbital is filled by the electrons of the elements belonging to the second group, which includes alkaline earth metals.

Groups 3-10

The third group, in its turn, contains electrons located at the 6d orbital. Unlike the first two groups, this one incorporates the so-called transition metals, also known as semi-metals, or metalloids. The electrons in the atoms of the elements located in the specified group. There are three electrons on the orbital of the elements correspondingly. Furthermore, groups 3-10 do not have specific names based on their properties and, therefore, are labelled based on the elements that they contain (e.g., the scandium family, the titanium family, etc.).

The eleventh group includes coinage metals. The orbital that corresponds to the group is 3d and 4s. The name of the group can be explained by the fact that the metals in it were used to make coins.

The group, including volatile metals, incorporates Zinc (Zn), Cadmium (Cd) and Mercury (Hg). Among the essential properties of the metals, the fact that they are soft needs to be brought up.

Also known as the Boron group, it includes Boron (B), Aluminium (Al), Gallium (Ga), Indium (In), and Thallium (Tl), as the title shows. The number of electrons per shell ranges from 2 to 32. The elements fill the full s orbital and partially the p one (with an electron on it).

Including crystallogens, the group in question belongs to the carbon family. The electrons fill the 2p orbital.

The Pnictogen family, or Group 15, is also referred to as the nitrogen family. The elements in the identified groups can form very strong covalent bonds. The elements have five electrons, in general (specifically, the electrons are located in the 2s and the 3p orbital).

Also labeled as the Oxygen family, the specified group includes Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), and Polonium (Po). It should be noted that Polonium is the only radioactive element in the group.

Group 17, or Halogens, is known for the unique property of its elements, which is creating compounds with any other element.

Last but not least, the group includes the so-called noble gases. The term was coined to indicate their low ability to produce a reaction with any other element under standard conditions.

Argon n.d., 2016, Web.

Beryllium n.d., 2016, Web.

Boron n.d., 2016, Web.

Cadmium n.d., 2016, Web.

Carbon n.d., 2016, Web.

Chlorine n.d., 2016, Web.

Copper n.d., 2016, Web.

Lithium n.d., 2016, Web.

Nitrogen n.d., 2016, Web.

Sulfur n.d., 2016, Web.

Titanium n.d., 2016, Web.

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The periodic table has gone through many changes since Dmitri Mendeleev drew up its original design in 1869, yet both the first table and the modern periodic table are important for the same reason: The periodic table organizes elements according to similar properties so you can tell the characteristics of an element just by looking at its location on the table.

Before all naturally occurring elements were discovered, the periodic table was used to predict the chemical and physical properties of elements in the gaps on the table. Today, the table can be used to predict properties of elements yet to be discovered, although these new elements are all highly radioactive and break down into more familiar elements almost instantly.

Now, the table is useful for modern students and scientists because it helps predict the types of chemical reactions that a particular element is likely to participate in. Rather than memorizing facts and figures for each element, students and scientists need only glance at the table to learn much about the reactivity of an element, whether it is likely to conduct electricity, whether it is hard or soft, and many other characteristics.

Elements in the same column as one other are known as groups and they share similar properties. For example, the elements in the first column (the alkali metals ) are all metals that usually carry a 1+ charge in reactions, react vigorously with water, and combine readily with nonmetals.

Elements in the same row as one other are known as periods and they share the same highest unexcited electron energy level.

Another useful feature of the periodic table is that most tables provide all the information you need to balance chemical reactions at a glance. The table tells each element's atomic number and usually its atomic weight. The typical charge of an element is indicated by its group.

Trends or Periodicity

The periodic table is organized according to trends in element properties.

As you move from left to right across a row of elements, the atomic radius (the size of an element's atoms) decreases, ionization energy (the energy required to remove an electron from an atom) increases, electron affinity (the amount of energy released when an atom forms a negative ion) generally increases, and electronegativity (an atom's tendency to attract a pair of electrons) increases.

As you move from top to bottom down a column of elements, the atomic radius increases, ionization energy decreases, electron affinity usually decreases, and electronegativity decreases.

To summarize, the periodic table is important because it is organized to provide a great deal of information about elements and how they relate to one another in one easy-to-use reference.

• The table can be used to predict the properties of elements, even those that have not yet been discovered.
• Columns (groups) and rows (periods) indicate elements that share similar characteristics.
• The table makes trends in element properties apparent and easy to understand.
• The table provides important information used to balance chemical equations .
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• What Are the Parts of the Periodic Table?
• Periodicity Definition in Chemistry
• What Is Periodicity on the Periodic Table?
• Periodic Table Study Guide - Introduction & History
• The Periodic Properties of the Elements
• 10 Periodic Table Facts
• Periodic Law Definition in Chemistry

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4.6: The Periodic Table of the Elements

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Learning Objectives

• To become familiar with the organization of the periodic table.

Rutherford’s nuclear model of the atom helped explain why atoms of different elements exhibit different chemical behavior. The identity of an element is defined by its atomic number (Z) , the number of protons in the nucleus of an atom of the element. The atomic number is therefore different for each element. The known elements are arranged in order of increasing Z in the periodic table ( Figure \(\PageIndex{1}\) ). The rationale for the peculiar format of the periodic table is explained later. E ach element is assigned a unique one-, two-, or three-letter symbol. The names of the elements are listed in the periodic table, along with their symbols, atomic numbers, and atomic masses. The chemistry of each element is determined by its number of protons and electrons. In a neutral atom, the number of electrons equals the number of protons.

The elements are arranged in a periodic table , which is probably the single most important learning aid in chemistry. It summarizes huge amounts of information about the elements in a way that facilitates the prediction of many of their properties and chemical reactions. The elements are arranged in seven horizontal rows, in order of increasing atomic number from left to right and top to bottom. The rows are called periods, and they are numbered from 1 to 7. The elements are stacked in such a way that elements with similar chemical properties form vertical columns, called groups, numbered from 1 to 18 (older periodic tables use a system based on roman numerals). Groups 1, 2, and 13–18 are the main group elements, listed as A in older tables. Groups 3–12 are in the middle of the periodic table and are the transition elements, listed as B in older tables. The two rows of 14 elements at the bottom of the periodic table are the lanthanides and the actinides, whose positions in the periodic table are indicated in group 3.

Metals, Nonmetals, and Semimetals

The heavy orange zigzag line running diagonally from the upper left to the lower right through groups 13–16 in Figure \(\PageIndex{1}\) divides the elements into metals (in blue, below and to the left of the line) and nonmetals (in bronze, above and to the right of the line). Gold-colored lements that lie along the diagonal line exhibit properties intermediate between metals and nonmetals; they are called semimetals.

The distinction between metals and nonmetals is one of the most fundamental in chemistry. Metals—such as copper or gold—are good conductors of electricity and heat; they can be pulled into wires because they are ductile; they can be hammered or pressed into thin sheets or foils because they are malleable; and most have a shiny appearance, so they are lustrous. The vast majority of the known elements are metals. Of the metals, only mercury is a liquid at room temperature and pressure; all the rest are solids.

Nonmetals, in contrast, are generally poor conductors of heat and electricity and are not lustrous. Nonmetals can be gases (such as chlorine), liquids (such as bromine), or solids (such as iodine) at room temperature and pressure. Most solid nonmetals are brittle, so they break into small pieces when hit with a hammer or pulled into a wire. As expected, semimetals exhibit properties intermediate between metals and nonmetals.

Example \(\PageIndex{1}\): Classifying Elements

Based on its position in the periodic table, do you expect selenium to be a metal, a nonmetal, or a semimetal?

Given : element

Asked for : classification

Find selenium in the periodic table shown in Figure \(\PageIndex{1}\) and then classify the element according to its location.

The atomic number of selenium is 34, which places it in period 4 and group 16. In Figure \(\PageIndex{1}\) , selenium lies above and to the right of the diagonal line marking the boundary between metals and nonmetals, so it should be a nonmetal. Note, however, that because selenium is close to the metal-nonmetal dividing line, it would not be surprising if selenium were similar to a semimetal in some of its properties.

Exercise \(\PageIndex{1}\)

Based on its location in the periodic table, do you expect indium to be a nonmetal, a metal, or a semimetal?

As previously noted, the periodic table is arranged so that elements with similar chemical behaviors are in the same group. Chemists often make general statements about the properties of the elements in a group using descriptive names with historical origins. For example, the elements of Group 1 are known as the alkali metals, Group 2 are the alkaline earth metals, Group 17 are the halogens, and Group 18 are the noble gases.

Group 1: The Alkali Metals

The alkali metals are lithium, sodium, potassium, rubidium, cesium, and francium. Hydrogen is unique in that it is generally placed in Group 1, but it is not a metal. The compounds of the alkali metals are common in nature and daily life. One example is table salt (sodium chloride); lithium compounds are used in greases, in batteries, and as drugs to treat patients who exhibit manic-depressive, or bipolar, behavior. Although lithium, rubidium, and cesium are relatively rare in nature, and francium is so unstable and highly radioactive that it exists in only trace amounts, sodium and potassium are the seventh and eighth most abundant elements in Earth’s crust, respectively.

Group 2: The Alkaline Earth Metals

The alkaline earth metals are beryllium, magnesium, calcium, strontium, barium, and radium. Beryllium, strontium, and barium are rare, and radium is unstable and highly radioactive. In contrast, calcium and magnesium are the fifth and sixth most abundant elements on Earth, respectively; they are found in huge deposits of limestone and other minerals.

Group 17: The Halogens

The halogens are fluorine, chlorine, bromine, iodine, and astatine. The name halogen is derived from the Greek words for “salt forming,” which reflects that all the halogens react readily with metals to form compounds, such as sodium chloride and calcium chloride (used in some areas as road salt).

Compounds that contain the fluoride ion are added to toothpaste and the water supply to prevent dental cavities. Fluorine is also found in Teflon coatings on kitchen utensils. Although chlorofluorocarbon propellants and refrigerants are believed to lead to the depletion of Earth’s ozone layer and contain both fluorine and chlorine, the latter is responsible for the adverse effect on the ozone layer. Bromine and iodine are less abundant than chlorine, and astatine is so radioactive that it exists in only negligible amounts in nature.

Group 18: The Noble Gases

The noble gases are helium, neon, argon, krypton, xenon, and radon. Because the noble gases are composed of only single atoms, they are called monatomic. At room temperature and pressure, they are unreactive gases. Because of their lack of reactivity, for many years they were called inert gases or rare gases. However, the first chemical compounds containing the noble gases were prepared in 1962. Although the noble gases are relatively minor constituents of the atmosphere, natural gas contains substantial amounts of helium. Because of its low reactivity, argon is often used as an unreactive (inert) atmosphere for welding and in light bulbs. The red light emitted by neon in a gas discharge tube is used in neon lights.

The noble gases are unreactive at room temperature and pressure.

The periodic table is used as a predictive tool. It arranges of the elements in order of increasing atomic number. Elements that exhibit similar chemistry appear in vertical columns called groups (numbered 1–18 from left to right); the seven horizontal rows are called periods. Some of the groups have widely-used common names, including the alkali metals (Group 1) and the alkaline earth metals (Group 2) on the far left, and the halogens (Group 17) and the noble gases (Group 18) on the far right. The elements can be broadly divided into metals, nonmetals, and semimetals. Semimetals exhibit properties intermediate between those of metals and nonmetals. Metals are located on the left of the periodic table, and nonmetals are located on the upper right. They are separated by a diagonal band of semimetals. Metals are lustrous, good conductors of electricity, and readily shaped (they are ductile and malleable), whereas solid nonmetals are generally brittle and poor electrical conductors. Other important groupings of elements in the periodic table are the main group elements, the transition metals, the lanthanides, and the actinides.

Introduction to the Periodic Table of the Elements

Periodic Table with molecular model (OntheRunPhoto, iStockphoto)

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Learn about the elements of the periodic table and how they are organized.

Matter that is composed of only one type of atom is called an  element . People have known about some elements such as gold, silver, iron, mercury and tin for hundreds of years. Scientists discovered many others during the 18th and 19th centuries. Some  elements have even been discovered within your lifetime !

Naming the Elements

Many of the elements have their name or their symbol based in Latin or Greek. For example:

As well for famous scientists such as:

Other elements have been named for places, such as:

Did you know? Argentina got its name from an element. Its name is from the latin name of silver,  argentum . This is because its natives gave silver gifts to the first European conquerors.

The periodic table of the elements

The  periodic table of the elements  is a visual and logical way to organize all elements.  Dmitri Mendeleev,  a Russian scientist, is usually credited with creating the first periodic table in 1869. In Germany,  Lothar Meyer  also came up with an almost identical table later the same year. Some people attribute the first categorization to  Alexandre-Emile Béguyer de Chancourtois .

Each element on the periodic table is listed in a box with its atomic symbol and atomic number. The element’s full name and atomic mass is also sometimes indicated. The image below shows a typical entry for the element calcium.

 The number above the atomic symbol represents the atomic number. The  atomic number  is equal to the number of  protons  in an atom's nucleus and determines which element an atom is. For example, any atom that contains exactly 20 protons in its nucleus is an atom of calcium. The number of protons influences the chemical behaviour of an element.

The  atomic mass  of an element refers to the average mass of an element in  unified atomic mass units   (short form  u ). This is usually given at the bottom of an element’s entry in the periodic table. The  atomic mass  is a decimal number because it is an average of the masses of the various  isotopes  of an element. Isotopes of an element have the same number of protons, but different numbers of neutrons. For example, hydrogen has three naturally occurring isotopes. The most common isotope,  protium , has no neutrons! Deuterium has one neutron and tritium has two neutrons.

Scientists have also created other highly unstable isotopes of hydrogen. Those, with up to 7 neutrons, do not occur in nature.

Getting the average number of neutrons for an element is easy. You need to subtract the number of protons (atomic number) from the atomic mass. For example, calcium has an atomic number of 20 and an atomic mass of 40.078 u. By subtracting 20 from 40 (rounded), we determine that the average number of neutrons is 20.

Arrangement of Elements

Mendeleev organized the table so that the elements from low to high atomic mass. Then he organized rows and columns to highlight elements with similar chemical properties. The rows are known as  periods . All elements in the same period have the same number of  electron layers , also called shells. The columns are known as  groups  (or families).

Mendeleev also left gaps in his table. He believed that some elements had not yet been discovered. Those gaps were for undiscovered elements. Mendeleev predicted the missing elements would have the same properties as the other elements in the same column.

Discoveries led to many new elements during the 1900s. Mendeleïev’s table was rearranged to give us the format that we use today. But the principles of groups and periods still exist in our modern version.

The groups on the periodic table are  numbered from 1 to 18 . The elements within a group usually have similar properties. For example, the elements of group 18 are all gases that do not readily react with other elements. Groups can sometimes be given names. For example, the elements in group 18 are called ‘ Noble Gases’ .

Did you know? Noble gases get their  name  because the ability to stay calm and not react is considered noble in humans!

Except for hydrogen, elements on the left side of the periodic table are metals. Metals are solid at room temperature, except for mercury, which is liquid. They are also malleable and can be ductile. Ductile is another word for stretchy. Metals also have a shiny appearance. They are good conductors of heat and electricity.

Elements on the right side of the periodic table are non-metals.  Non-metals  can be solids, liquids or gases at room temperature. They are poor conductors of heat and electricity.

Elements that have some metallic as well as non-metallic properties are known as  metalloids  (e.g., silicon, arsenic, etc.).

We sometimes call the elements in groups 3 to 12 the  d-block . These elements are also known as  transition elements .

There are two rows placed at the bottom of the periodic table. These are the Lanthanide and Actinide series. The Lanthanide series includes fifteen metallic elements with atomic numbers 57 through 71. The Lanthanide elements, along with the two elements of group 3, Scandium and Yttrium, are often called the rare earth elements. You can find Lanthanides naturally on Earth. The Actinide series is much different. All of the Actinide elements are radioactive and some are not found in nature. Some of the elements with higher atomic numbers can only be created in laboratories.

Trends of the Periodic Table

In general, the size or radius of an atom increases as you move down a group. As you go down, elements have more layers of electrons, which results in bigger atoms. From left to right within a period, atomic size generally decreases. This is because more protons pull the electrons closer to the nucleus.

Electronegativity  is the chemical property that describes the tendency of an atom to form bonds with others. Electronegativity usually decreases down each group and increases from left to right across a period. The most electronegative elements are in the upper right hand side of the periodic table. This is because non-metals have high electronegativity. They tend to attract electrons from other atoms. While metals tend to lose electrons easily.

Assigning Charges to an Element

To become electrically-charged, stable atoms must lose or gain electrons. There are two types of electric charges. There are  positive charges  (+) and  negative charges  (-). Substances with the same charge will repel each other. While substances with opposite charges attract each other. This is similar to how magnets work!

The atoms of some elements can lose or gain electrons more easily than others. For example, the elements in group 1 (e.g., Li, Na, K) can easily lose one electron, which gives them a  +1 charge . The elements in group 17 (Cl, Br, I) can easily gain one electron, which gives them a  -1 charge . When an atom or molecule has a charge, it is said to be an  ion . When an atom or molecule has more electrons than protons, the result is a negative ion, called an  anion . When it has fewer electrons than protons, then the result is a positive ion, called a  cation .

Each group in the periodic table has a specific charge commonly associated with its ions. The exceptions to this are the transition elements, and the lanthanide and actinide series. Below are the charges for the different groups within the table.

You may have noticed that the elements in group 18 have a charge of 0. This is because these elements are very stable and they do not readily gain or lose electrons.

The photographic periodic table This page by periodictable.com has photographs of common things made from, or using, each of the element.

Introduction to the Particle Theory of Matter This article from Let’s Talk Science introduces the history of the theory that describes matter as different assemblies of smaller particles, or atoms.

The Genius of Mendeleev's Table This article from Let’s Talk Science presents how Dmitri Mendeleev created the first periodic table.

Introduction to the Atom This article from Let’s Talk Science presents the history of the different atomic models created.

The Newest Elements on the Periodic Table This article from Let’s Talk Science presents the most recent elements added to the periodic table.

The Periodic Table of Elements This page by the Thomas Jefferson National Accelerator Facility has an interactive periodic table as well as flash cards, games, etc. about the elements.

Helmenstine, A. (2019).  Introduction to the Periodic Table . Thought.co.

Webelements.com. (n.d.).  The periodic table of the elements .

Western Oregon University. (n.d.).  A Brief History Of The Development Of Periodic Table .  People.wou.edu

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How the periodic table went from a sketch to an enduring masterpiece.

150 years ago, Mendeleev perceived the relationships of the chemical elements

REVOLUTIONARY   Russian chemist Dmitrii Mendeleev (shown around 1880) was the first to publish a periodic table, which put the known elements into a logical order and left room for elements not yet discovered.

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By Tom Siegfried

January 8, 2019 at 12:29 pm

Every field of science has its favorite anniversary.

For physics, it’s Newton’s Principia of 1687, the book that introduced the laws of motion and gravity. Biology celebrates Darwin’s On the Origin of Species (1859) along with his birthday (1809). Astronomy fans commemorate 1543, when Copernicus placed the sun at the center of the solar system.

And for chemistry, no cause for celebration surpasses the origin of the periodic table of the elements, created 150 years ago this March by the Russian chemist Dmitrii Ivanovich Mendeleev.

Mendeleev’s table has become as familiar to chemistry students as spreadsheets are to accountants. It summarizes an entire science in 100 or so squares containing symbols and numbers. It enumerates the elements that compose all earthly substances, arranged so as to reveal patterns in their properties, guiding the pursuit of chemical research both in theory and in practice.

“The periodic table,” wrote the chemist Peter Atkins, “is arguably the most important concept in chemistry.”

Mendeleev’s table looked like an ad hoc chart, but he intended the table to express a deep scientific truth he had uncovered: the periodic law. His law revealed profound familial relationships among the known chemical elements — they exhibited similar properties at regular intervals (or periods) when arranged in order of their atomic weights — and enabled Mendeleev to predict the existence of elements that had not yet been discovered.

“Before the promulgation of this law the chemical elements were mere fragmentary, incidental facts in Nature,” Mendeleev declared. “The law of periodicity first enabled us to perceive undiscovered elements at a distance which formerly was inaccessible to chemical vision.”

Mendeleev’s table did more than foretell the existence of new elements. It validated the then-controversial belief in the reality of atoms. It hinted at the existence of subatomic structure and anticipated the mathematical apparatus underlying the rules governing matter that eventually revealed itself in quantum theory. His table finished the transformation of chemical science from the medieval magical mysticism of alchemy to the realm of modern scientific rigor. The periodic table symbolizes not merely the constituents of matter, but the logical cogency and principled rationality of all science.

An ordered vision

Mendeleev’s periodic table, published in 1869, was a vertical chart that organized 63 known elements by atomic weight. This arrangement placed elements with similar properties into horizontal rows. The title, translated from Russian, reads: “Draft of system of elements: based on their atomic masses and chemical characteristics.”

Laying the groundwork

Legend has it that Mendeleev conceived and created his table in a single day: February 17, 1869, on the Russian calendar (March 1 in most of the rest of the world). But that’s probably an exaggeration. Mendeleev had been thinking about grouping the elements for years, and other chemists had considered the notion of relationships among the elements several times in the preceding decades.

In fact, German chemist Johann Wolfgang Döbereiner noticed peculiarities in groupings of elements as early as 1817. In those days, chemists hadn’t yet fully grasped the nature of atoms, as described in the atomic theory proposed by English schoolteacher John Dalton in 1808. In his New System of Chemical Philosophy , Dalton explained chemical reactions by assuming that each elementary substance was made of a particular type of atom.

Chemical reactions, Dalton proposed, produced new substances when atoms were disconnected or joined. Any given element consisted entirely of one kind of atom, he reasoned, distinguished from other kinds by weight. Oxygen atoms weighed eight times as much as hydrogen atoms; carbon atoms were six times as heavy as hydrogen, Dalton believed. When elements combined to make new substances, the amounts that reacted could be calculated with knowledge of those atomic weights.

Dalton was wrong about some of the weights — oxygen is really 16 times the weight of hydrogen, and carbon is 12 times heavier than hydrogen. But his theory made the idea of atoms useful, inspiring a revolution in chemistry. Measuring atomic weights accurately became a prime preoccupation for chemists in the decades that followed.

When contemplating those weights, Döbereiner noted that certain sets of three elements (he called them triads) showed a peculiar relationship. Bromine, for example, had an atomic weight midway between the weights of chlorine and iodine, and all three elements exhibited similar chemical behavior. Lithium, sodium and potassium were also a triad.

Intrinsic order

Every element on this venerated table has its own story. All together, they capture the entire repertoire of known chemistry. Read up on the tales between the lines . 

Other chemists perceived links between atomic weights and chemical properties, but it was not until the 1860s that atomic weights had been well enough understood and measured for deeper insights to emerge. In England, the chemist John Newlands noticed that arranging the known elements in order of increasing atomic weight produced a recurrence of chemical properties every eighth element, a pattern he called the “law of octaves” in an 1865 paper. But Newlands’ pattern did not hold up very well after the first couple of octaves, leading a critic to suggest that he should try arranging the elements in alphabetical order instead. Clearly, the relationship of element properties and atomic weights was a bit more complicated, as Mendeleev soon realized.

Organizing the elements

Born in Tobolsk, in Siberia, in 1834 (his parents’ 17th child), Mendeleev lived a dispersed life, pursuing multiple interests and traveling a higgledy-piggledy path to prominence. During his higher education at a teaching institute in St. Petersburg, he nearly died from a serious illness. After graduation, he taught at middle schools (a requirement of his scholarship at the teaching institute), and while teaching math and science, he conducted research for his master’s degree.

He then worked as a tutor and lecturer (along with some popular science writing on the side) until earning a fellowship for an extended tour of research at Europe’s most prominent university chemistry laboratories.

When he returned to St. Petersburg, he had no job, so he wrote a masterful handbook on organic chemistry in hopes of winning a large cash prize. It was a long shot that paid off, with the lucrative Demidov Prize in 1862. He also found work as an editor, translator and consultant to various chemical industries. Eventually he returned to research, earning his Ph.D. in 1865 and then becoming a professor at the University of St. Petersburg.

Soon thereafter, Mendeleev found himself about to teach inorganic chemistry. In preparing to master that new (to him) field, he was unimpressed by the available textbooks. So he decided to write his own. Organizing the text required organizing the elements, so the question of how best to arrange them was on his mind.

By early 1869, Mendeleev had made enough progress to realize that some groups of similar elements showed a regular increase in atomic weights; other elements with roughly equal atomic weights shared common properties. It appeared that ordering the elements by their atomic weight was the key to categorizing them.

By Mendeleev’s own account, he structured his thinking by writing each of the 63 known elements’ properties on an individual note card. Then, by way of a sort of game of chemical solitaire, he found the pattern he was seeking. Arranging the cards in vertical columns from lower to higher atomic weights placed elements with similar properties in each horizontal row. Mendeleev’s periodic table was born. He sketched out his table on March 1, sent it to the printer and incorporated it into his soon-to-be-published textbook. He quickly prepared a paper to be presented to the Russian Chemical Society.

A handwritten draft of Mendeleev’s periodic table, in which he organized the elements by atomic weight to reveal the periodic law, showing how elements had similar properties at regular intervals, or periodicities.

“Elements arranged according to the size of their atomic weights show clear periodic properties,” Mendeleev declared in his paper. “All the comparisons which I have made … lead me to conclude that the size of the atomic weight determines the nature of the elements.”

Meanwhile, the German chemist Lothar Meyer had also been working on organizing the elements. He prepared a table similar to Mendeleev’s, perhaps even before Mendeleev did. But Mendeleev published first.

More important than beating Meyer to the publication punch, though, was Mendeleev’s use of his table to make bold predictions about undiscovered elements. In preparing his table, Mendeleev had noticed that some note cards were missing. He had to leave blank spaces to get the known elements to properly align. Within his lifetime, three of those blanks were filled with the previously unknown elements gallium, scandium and germanium.

Not only had Mendeleev predicted the existence of these elements, but he had also correctly described their properties in detail. Gallium, for instance, discovered in 1875, had an atomic weight (as measured then) of 69.9 and a density six times that of water. Mendeleev had predicted an element (he called it eka-aluminum) with just that density and an atomic weight of 68. His predictions for eka-silicon closely matched germanium (discovered in 1886) in atomic weight (72 predicted, 72.3 observed) and density (5.5 versus 5.469). He also correctly predicted the density of germanium’s compounds with oxygen and chlorine.

Mendeleev’s table had become an oracle. It was as if end-of-game Scrabble tiles spelled out the secrets of the universe. While others had glimpsed the periodic law’s power, Mendeleev was the master at exploiting it.

Mendeleev’s successful predictions earned him legendary status as a maestro of chemical wizardry. But today, historians dispute whether the discovery of the predicted elements cemented the acceptance of his periodic law. The law’s approval may have been more due to its power to explain established chemical relationships. In any case, Mendeleev’s prognosticative accuracy certainly attracted attention to the merits of his table.

Elements arranged according to the size of their atomic weights show clear periodic properties.

— Dmitrii Mendeleev, 1869

By the 1890s, chemists widely recognized his law as a landmark in chemical knowledge. In 1900, the future Nobel chemistry laureate William Ramsay called it “the greatest generalization which has as yet been made in chemistry.” And Mendeleev had done it without understanding in any deep way why it worked at all.

A mathematical map

In many instances in the history of science, grand predictions based on novel equations have turned out to be correct. Somehow math reveals some of nature’s secrets before experimenters find them. Antimatter is one example, the expansion of the universe another. In Mendeleev’s case, the predictions of new elements emerged without any creative mathematics. But in fact, Mendeleev had discovered a deep mathematical map of nature, for his table reflected the implications of quantum mechanics, the mathematical rules governing atomic architecture.

In his textbook, Mendeleev had noted that “internal differences of the matter that comprises the atoms” could be responsible for the elements’ periodically recurring properties. But he did not pursue that line of thought. In fact, over the years he waffled about how important atomic theory was for his table.

But others could read the table’s message. In 1888, German chemist Johannes Wislicenus declared that the periodicity of the elements’ properties when arranged by weight indicated that atoms are composed of regular arrangements of smaller particles. So in a sense, Mendeleev’s table did anticipate (and provide evidence for) the complex internal structure of atoms, at a time when nobody had any idea what an atom really looked like, or even whether it had any internal structure at all.

By the time of Mendeleev’s death in 1907, scientists knew that atoms had parts: electrons, which carried a negative electric charge, plus some positively charged component to make atoms electrically neutral. A key clue to how those parts were arranged came in 1911, when the physicist Ernest Rutherford, working at the University of Manchester in England, discovered the atomic nucleus. Shortly thereafter Henry Moseley, a physicist who had worked with Rutherford, demonstrated that the amount of positive charge in the nucleus (the number of protons it contained, or its “atomic number”) determined the correct order of the elements in the periodic table.

Atomic weight was closely related to Moseley’s atomic number — close enough that ordering elements by weight differs in only a few spots from ordering by number. Mendeleev had insisted that those weights were wrong and needed to be remeasured, and in some cases he was right. A few discrepancies remained, but Moseley’s atomic number set the table straight.

At about the same time, the Danish physicist Niels Bohr realized that quantum theory governed the arrangement of electrons surrounding the nucleus and that the outermost electrons determined an element’s chemical properties.

Similar arrangements of the outer electrons would recur periodically, explaining the patterns that Mendeleev’s table had originally revealed. Bohr created his own version of the table in 1922, based on experimental measurements of electron energies (along with some guidance from the periodic law).

Bohr’s table added elements discovered since 1869, but it was still, in essence, the periodic arrangement that Mendeleev had discovered. Without the slightest clue to quantum theory, Mendeleev had created a table reflecting the atomic architecture that quantum physics dictated.

In Danish physicist Niels Bohr’s 1922 version of the periodic table, adapted from a table by Danish chemist Julius Thomsen, elements with similar properties occupy horizontal rows connected by lines. The empty box on the right marks the expected occurrence of a group of elements that are chemically similar to the rare earth elements (numbers 58–70) in the preceding column.

Bohr’s new table was neither the first nor last variant on Mendeleev’s original design. Hundreds of versions of the periodic table have been devised and published. The modern form, a horizontal design in contrast with Mendeleev’s original vertical version, became widely popular only after World War II, largely due to the work of the American chemist Glenn Seaborg (a longtime member of the board of Science Service, the original publisher of Science News ).

Seaborg and collaborators had synthetically produced several new elements with atomic numbers beyond uranium, the last naturally occurring element in the table. Seaborg saw that these elements, the transuranics (plus the three elements preceding uranium) demanded a new row in the table, something Mendeleev had not foreseen. Seaborg’s table added the row for those elements beneath a similar row for the rare earth elements, whose proper place had never been quite clear, either. “It took a lot of guts to buck Mendeleev,” Seaborg, who died in 1999, said in a 1997 interview.

Seaborg’s contributions to chemistry earned him the honor of his own namesake element, seaborgium, number 106. It’s one of a handful of elements named to honor a famous scientist, a list that includes, of course, element 101, discovered by Seaborg and colleagues in 1955 and named mendelevium — for the chemist who above all others deserved a place at the periodic table.

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The Periodic Table and its Iconicity: an Essay

2019, Substantia

In this essay, we aim to provide an overview of the periodic table's origins and history, and of the elements which conspired to make it chemistry's most rec-ognisable icon. We pay attention to Mendeleev's role in the development of a system for organising the elements and chemical knowledge while facilitating the teaching of chemistry. We look at how the reception of the table in different chemical communities was dependent on the local scientific, cultural and political context, but argue that its eventual universal acceptance is due to its unique ability to accommodate possessed knowledge while enabling novel predictions. Furthermore, we argue that its capacity to unify apparently disconnected phenomena under a simple framework facilitates our understanding of periodicity, making the table an icon of aesthetic value, and an object of philosophical inquiry. Finally, we briefly explore the table's iconicity throughout its representations in pop art and science fiction.

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The Incredible History of The Periodic Table

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Periodic Table of Elements: History, Element Names, Symbols, Trick to Learn

Modern periodic table: this article will elaborate on the periodic table and its characteristics. read here to learn the key features, history, list of elements and easy tricks to remember the periodic table..

Key Features of the Periodic Table

• Periods: There are seven horizontal rows in the periodic table. Elements in the same period have the same number of electron shells.
• Groups: There are 18 vertical columns. Elements in the same group have similar chemical properties and the same number of electrons in their outermost shell.
• Element Symbols: Each element is represented by a unique one- or two-letter symbol, usually derived from its English or Latin name (e.g., H for Hydrogen, Au for Gold from "Aurum").
• Atomic Number: The atomic number, which indicates the number of protons in an atom, is displayed above the element symbol.
• Atomic Mass: The atomic mass, usually found below the element symbol, indicates the weighted average mass of an element's isotopes.

History of Periodic Table of Elements

• Johann Dobereiner, a German chemist, was the first to consider trends among element properties in the early 1800s.
• In 1829, Dobereiner identified similarities in the physical and chemical properties of several groups of three elements, called Triads.
• Dobereiner noted that the atomic weight of the middle element in each Triad was approximately halfway between the atomic weights of the other two, and its properties were also intermediate.
• Dobereiner's Law of Triads was dismissed as coincidental since it only applied to a few elements.
• In 1862, French geologist A.E.B. de Chancourtois arranged elements by increasing atomic weights and created a cylindrical table to show periodic property recurrence, but it gained little attention.
• John Alexander Newlands, an English chemist, proposed the Law of Octaves in 1865, observing that every eighth element had similar properties when arranged by increasing atomic weight.
• Newlands compared this to musical octaves but his Law of Octaves applied only up to calcium and was not widely accepted initially.
• Newlands was later awarded the Davy Medal in 1887 by the Royal Society, London.
• Dmitri Mendeleev (1834-1907) and Lothar Meyer (1830-1895) independently developed the Periodic Law in 1869, observing periodic similarities in element properties when arranged by atomic weight.
• Lothar Meyer plotted physical properties against atomic weight and found a repeating pattern, noting changes in the length of the pattern.
• By 1868, Lothar Meyer had a table resembling the Modern Periodic Table but published it after Mendeleev.
• Mendeleev is credited with publishing the Periodic Law, which states that the properties of elements are a periodic function of their atomic weights.
• Mendeleev's work led to the development of the Modern Periodic Table, recognising periodic trends and allowing predictions of undiscovered elements.
• Mendeleev arranged elements in a table with horizontal rows and vertical columns based on increasing atomic weights.
• Elements with similar properties were placed in the same vertical column or group.
• Mendeleev’s classification system was more elaborate than Lothar Meyer’s, considering a broader range of physical and chemical properties.
• He emphasised the importance of periodicity, using similarities in empirical formulas and properties of compounds to classify elements.
• Mendeleev recognised that some elements did not fit his classification scheme if strictly ordered by atomic weight.
• He chose to ignore the atomic weight order when necessary, believing that atomic measurements might be incorrect, to keep elements with similar properties together.

Elements of Periodic Table

This table includes all 118 known elements, providing their names, symbols, and atomic numbers.

Classification of Elements of Periodic Table

• Alkali Metals (Group 1): Highly reactive, especially with water (e.g., Lithium, Sodium, Potassium).
• Alkaline Earth Metals (Group 2): Less reactive than alkali metals, but still quite reactive (e.g., Beryllium, Magnesium, Calcium).
• Transition Metals (Groups 3-12): Includes elements like Iron, Copper, and Gold, known for their conductivity and malleability.
• Lanthanides: Rare earth elements, used in electronics and lasers (e.g., Lanthanum, Cerium).
• Actinides: Radioactive elements, some synthetic (e.g., Uranium, Plutonium).
• Hydrogen: Unique, not fitting into any single group.
• Halogens (Group 17): Very reactive nonmetals, form salts with metals (e.g., Fluorine, Chlorine, Bromine).
• Other Nonmetals: Essential for life, such as Carbon, Nitrogen, and Oxygen.
• Elements with properties intermediate between metals and nonmetals (e.g., Boron, Silicon, Germanium). These are often semiconductors.
• Noble Gases (Group 18)
• Inert gases, nonreactive due to having a complete valence electron shell (e.g., Helium, Neon, Argon).
• Atomic Radius: Decreases across a period and increases down a group.
• Ionization Energy: Increases across a period and decreases down a group.
• Electronegativity: Increases across a period and decreases down a group.
• States that the properties of elements are a periodic function of their atomic numbers. This law is the basis for the modern periodic table.

Tricks to Learn Periodic Table Through Easy Mnemonic Phases

Check here the easy ways to remember the periodic table for Indian students. Source: Quora

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Creating mnemonic phrases can help you remember the order of elements, especially for the first 20 elements or specific groups.

• Hydrogen (H), Helium (He), Lithium (Li), Beryllium (Be), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), Neon (Ne).
• Mnemonic: "Happy Henry Likes Beans Brown, Crusty, Not Over-Fried Nicely."
• Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), Francium (Fr).
• Mnemonic: "Little Naughty Kids Rub Cats Fur."
• Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At).
• Mnemonic: "Fat Cats Bring In Ants."

These phrases are general and are available on other free online sources. You may refer to these or create your own.

This article concludes here with all the relevant information taken from NCERT and other authorised sources on the periodic table of chemistry. For more related information, check out the official website of Jagran Josh.

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• How to calculate atomic mass from the periodic table? + Atomic mass is defined as the mass of an element that can be calculated by simply adding the number of protons and neutrons of that element. mass number = protons + neutrons.
• Name the first 20 elements. + The first 20 elements of the periodic table are as follows: H – Hydrogen, He – Helium, Li – Lithium, Be – Beryllium, B – Boron, C – Carbon, N – Nitrogen, O – Oxygen, F – Fluorine, Ne – Neon, Na – Sodium, Mg – Magnesium, Al – Aluminium, Si – Silicon. P – Phosphorus. S – Sulphur. Cl – Chlorine. Ar – Argon. K – Potassium. Ca – Calcium.
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Oliver Sacks: My Periodic Table

By Oliver Sacks

• July 24, 2015

I LOOK forward eagerly, almost greedily, to the weekly arrival of journals like Nature and Science, and turn at once to articles on the physical sciences — not, as perhaps I should, to articles on biology and medicine. It was the physical sciences that provided my first enchantment as a boy.

In a recent issue of Nature, there was a thrilling article by the Nobel Prize-winning physicist Frank Wilczek on a new way of calculating the slightly different masses of neutrons and protons. The new calculation confirms that neutrons are very slightly heavier than protons — the ratio of their masses being 939.56563 to 938.27231 — a trivial difference, one might think, but if it were otherwise the universe as we know it could never have developed. The ability to calculate this, Dr. Wilczek wrote, “encourages us to predict a future in which nuclear physics reaches the level of precision and versatility that atomic physics has already achieved” — a revolution that, alas, I will never see.

Francis Crick was convinced that “the hard problem” — understanding how the brain gives rise to consciousness — would be solved by 2030. “You will see it,” he often said to my neuroscientist friend Ralph, “and you may, too, Oliver, if you live to my age.” Crick lived to his late 80s, working and thinking about consciousness till the last. Ralph died prematurely, at age 52, and now I am terminally ill, at the age of 82. I have to say that I am not too exercised by “the hard problem” of consciousness — indeed, I do not see it as a problem at all; but I am sad that I will not see the new nuclear physics that Dr. Wilczek envisages, nor a thousand other breakthroughs in the physical and biological sciences.

A few weeks ago, in the country, far from the lights of the city, I saw the entire sky “powdered with stars” (in Milton’s words); such a sky, I imagined, could be seen only on high, dry plateaus like that of Atacama in Chile (where some of the world’s most powerful telescopes are). It was this celestial splendor that suddenly made me realize how little time, how little life, I had left. My sense of the heavens’ beauty, of eternity, was inseparably mixed for me with a sense of transience — and death.

I told my friends Kate and Allen, “I would like to see such a sky again when I am dying.”

“We’ll wheel you outside,” they said.

I have been comforted, since I wrote in February about having metastatic cancer , by the hundreds of letters I have received, the expressions of love and appreciation, and the sense that (despite everything) I may have lived a good and useful life. I remain very glad and grateful for all this — yet none of it hits me as did that night sky full of stars.

I have tended since early boyhood to deal with loss — losing people dear to me — by turning to the nonhuman. When I was sent away to a boarding school as a child of 6, at the outset of the Second World War, numbers became my friends; when I returned to London at 10, the elements and the periodic table became my companions. Times of stress throughout my life have led me to turn, or return, to the physical sciences, a world where there is no life, but also no death.

And now, at this juncture, when death is no longer an abstract concept, but a presence — an all-too-close, not-to-be-denied presence — I am again surrounding myself, as I did when I was a boy, with metals and minerals, little emblems of eternity. At one end of my writing table, I have element 81 in a charming box, sent to me by element-friends in England: It says, “Happy Thallium Birthday,”a souvenir of my 81st birthday last July; then, a realm devoted to lead, element 82, for my just celebrated 82nd birthday earlier this month. Here, too, is a little lead casket, containing element 90, thorium, crystalline thorium, as beautiful as diamonds, and, of course, radioactive — hence the lead casket.

At the start of the year, in the weeks after I learned that I had cancer, I felt pretty well, despite my liver being half-occupied by metastases. When the cancer in my liver was treated in February by the injection of tiny beads into the hepatic arteries — a procedure called embolization — I felt awful for a couple of weeks but then super well, charged with physical and mental energy. (The metastases had almost all been wiped out by the embolization.) I had been given not a remission, but an intermission, a time to deepen friendships, to see patients, to write, and to travel back to my homeland, England. People could scarcely believe at this time that I had a terminal condition, and I could easily forget it myself.

This sense of health and energy started to decline as May moved into June, but I was able to celebrate my 82nd birthday in style. (Auden used to say that one should always celebrate one’s birthday, no matter how one felt.) But now, I have some nausea and loss of appetite; chills in the day, sweats at night; and, above all, a pervasive tiredness, with sudden exhaustion if I overdo things. I continue to swim daily, but more slowly now, as I am beginning to feel a little short of breath. I could deny it before, but I know I am ill now. A CT scan on July 7 confirmed that the metastases had not only regrown in my liver but had now spread beyond it as well.

I started a new sort of treatment — immunotherapy — last week. It is not without its hazards, but I hope it will give me a few more good months. But before beginning this, I wanted to have a little fun: a trip to North Carolina to see the wonderful lemur research center at Duke University. Lemurs are close to the ancestral stock from which all primates arose, and I am happy to think that one of my own ancestors, 50 million years ago, was a little tree-dwelling creature not so dissimilar to the lemurs of today. I love their leaping vitality, their inquisitive nature.

NEXT to the circle of lead on my table is the land of bismuth: naturally occurring bismuth from Australia; little limousine-shaped ingots of bismuth from a mine in Bolivia; bismuth slowly cooled from a melt to form beautiful iridescent crystals terraced like a Hopi village; and, in a nod to Euclid and the beauty of geometry, a cylinder and a sphere made of bismuth.

Bismuth is element 83. I do not think I will see my 83rd birthday, but I feel there is something hopeful, something encouraging, about having “83” around. Moreover, I have a soft spot for bismuth, a modest gray metal, often unregarded, ignored, even by metal lovers. My feeling as a doctor for the mistreated or marginalized extends into the inorganic world and finds a parallel in my feeling for bismuth.

I almost certainly will not see my polonium (84th) birthday, nor would I want any polonium around, with its intense, murderous radioactivity. But then, at the other end of my table — my periodic table — I have a beautifully machined piece of beryllium (element 4) to remind me of my childhood, and of how long ago my soon-to-end life began.

Oliver Sacks is a professor of neurology at the New York University School of Medicine, and the author, most recently, of the memoir “On the Move.”

The Periodic Table Of Elements Essay Example

• Pages: 5 (1227 words)
• Published: February 9, 2017
• Type: Essay

The Periodic Table of Elements is used as a way of displaying all the known chemical elements; it is accepted and used all over the world. The periodic table’s layout is very well structured; it consists of vertical rows called groups and horizontal rows called periods. It is one of the most important resources in chemistry and the key to discovering new elements. The beauty of the Periodic Table is that a lot of information about any element can be gathered just by looking at its position within the table.

Effectively understanding the table is essential for chemistry. The physical and chemical properties of elements can also be predicted; even the prediction of how a certain element will react with another can be made with good accuracy, using the periodic table.

This is all because of its trends and amazing structure. In the Periodic Table, a group is a vertical row going from top to bottom. Groups contain elements that have similar outer shell electron configurations. This means that if you look down a group of elements, e. g. Group 2 and write each of the elements electronic configurations, they will all end in S2.

Each group is numbered according to the outer shell electron configuration of the atoms, of that particular group of elements. Usually groups of elements are numbered in Roman numerals, going from I-VIII. A period is a horizontal row going from left to right across the table. Each period is labelled from 1-7 and contains elements that have electrons which are in the same outer shell. The number of an elements outer shell will be equal to th

number of the period it’s in.

Positioning of Elements in the Periodic Table Elements have been strategically placed around the periodic table to have something in common with the elements around them… Atomic radius or width of an atom is half of the length between two nuclei of a diatomic molecule. The atomic radius increases as you move down a group of the periodic table because each element going down has an extra main layer of electrons, therefore expanding the radius. Atomic radius decreases going across a eriod because of the increase in valence or outer electrons.

This makes the electrostatic force at a higher level meaning the attraction is stronger for valance layers and they are drawn in closer together therefore making the atomic radius smaller. Melting points for metals usually decrease as you go down a group and the melting point for non metals usually increases going down a group. This is because when substances melt the attractive forces keeping their particles together are either broken or loosened enough for them to move freely.

The stronger the forces are holding the particles together the more energy is needed to break the attractive forces, therefore making it have a much higher melting point. An example of this is silicon because its atoms are held so strongly together by covalent bonds that it makes its melting point extremely high. Boiling point decreases going down a group and is quite similar to melting points because the boiling breaks or loosens the attractive forces at certain temperatures. Substances which are held together by stronger covalent bonds require a much higher temperature for their boiling

Electrical conductivity is the ability that an element has to conduct electricity. Metals are good conductors because of their free moving electrons. Ionisation energy relates to electronegativity and is the amount of energy that is required to take the outmost electron from an atom. Ionization energy decreases going down a group because the atoms are getting bigger and their hold on their electrons is a lot weaker as they are further out. Ionization energy increases going across a period because the atoms are getting smaller because their high levels of electronegativity.

Combining Power or (valency) is the number of chemical bonds; an element has, made by its atoms. Reactivity and Electronegativity Reactivity is the speed of response that a chemical substance has to certain stimulation. In chemistry, this would be how long it takes to have a chemical reaction with another substance. Reactivity decreases going left to right across a period and increases going down a group.

This is because as we move down a group atoms increase in size and there is a weaker force llowing the electrons to be taken much easier making it reactive. Electronegativity is the measure of attraction that an elements atoms have for electrons. The stronger the attraction for electrons, the atom has, the higher its electronegativity will be. The type of bond that one atom will form with another can be, either ionic or covalent; and this can be determined by the different electronegativities, the two atoms have. Electronegativity increases going left to right across a period and decreases going down a group.

History of the Periodic Table: Many chemists during the 19th

century started categorizing elements according to the similarities of their properties, both physical and chemical. Further studies, modifications and refinements of the Periodic Table; as well as the discovery of new elements; has resulted in our modern Periodic Table. Dmitri Mendeleev is considered the inventor or Father of the Periodic Table but one of the first people known to tackle the challenge of trying to organise all of the known elements at the time, was Johann Dobereiner.

He started in 1817 and in 1829 he managed to develop a way of classifying some of the elements into groups of three. He called these groups triads. The elements which were placed into a triad had similar chemical properties and neatly arranged physical properties; these groups became known as Dobereiner’s triads. Then in 1863 John Newlands devised a way of organising the elements into octaves, creating and publishing the first periodic table which was arranged in order of atomic weight.

Newland was ridiculed at the time until finally, five years later, Dmitri Mendeleev; a Russian chemist published a more advanced periodic table, independently, which organized elements by increasing atomic mass in 1869. It was this table that was the fundamentals for our modern Periodic Table today, which is now arranged by atomic number. There were many gaps in Mendeleev’s table when it was first published. These gaps were for undiscovered elements such as Re, Fr and all of the noble gases etc.

The reason Mendeleev’s method for the table was accepted and he is known as the father of the periodic table was because he was so confident in himself, that he was able

to predict the properties of 3 undiscovered elements. The undiscovered elements now known as Sc, Ga and Ge, had very close properties to what Mendeleev had predicted using his table, therefore his periodic table became widely accepted. Mendeleev was able to accurately predict the existence and physical properties of unknown elements by looking at many of the trends within his table.

He focussed mainly on the properties of other elements and placed titanium in a row with silicon and carbon, leaving a gap next to aluminium and boron. This gap he predicted was an unknown element, whose atomic mass would be in between 40 and 48. In 1878 an element called scandium was discovered, its properties were, as predicted by Mendeleev and its atomic mass was around 45. He could also predict, to a certain extent, how these undiscovered elements would react with other elements.

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An Essay on Periodic Tables

• First Online: 05 July 2021

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• Pekka Pyykkö 5  

Part of the book series: Perspectives on the History of Chemistry ((PHC))

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After a compact history of the PT, from Döbereiner’s triads to the theoretical predictions up to element 172, a number of particular issues is discussed: Why may Z  = 172 be a limit for stable electron shells? What are the expected stability limits of the nuclear isotopes? When are formally empty atomic orbitals used in molecular electronic structures? What is ‘Secondary Periodicity’? When do the elements (Ir, Pt, Au), at the end of a bond, simulate (N, O, I), respectively? Some new suggestions for alternative PTs are commented upon. As a local connection, Johan Gadolin’s 1794 analysis of the Ytterby mineral is mentioned.

This chapter has been reproduced with permission from Ref. [ 1 ]. https://doi.org/10.1515/pac-2019-0801 . Copyright 2019 IUPAC and De Gruyter.

A collection of invited papers based on presentations at “Mendeleev 150”: 4th International Conference on the Periodic Table endorsed by IUPAC, Saint Petersburg (Russia), 26–28 July 2019.

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Three related topics on the periodic tables of elements

The Discovery of the Elements in the Periodic Table

In praise of triads

For more information on triads, see Chapter 3 of this volume—Ed.

Other chapters in this volume treat in greater detail Beguyer de Chancourtois (5), Newlands (6), Meyer (8 and 9), Mendeleev (2), and the discovery of predicted elements (10)—Ed.

Chapter 12 of this volume recounts the discovery of the noble gas elements—Ed.

That is, Seaborg chose names for the actinides by analogy to the lanthanides’ names—Ed.

The very compact size of the 5g shell would make these elements ‘superlanthanides’.

See Chapter 16 of the present volume for further information on the Madelung rule and related topics—Ed.

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Pyykkö, P. (2021). An Essay on Periodic Tables. In: Giunta, C.J., Mainz, V.V., Girolami, G.S. (eds) 150 Years of the Periodic Table. Perspectives on the History of Chemistry. Springer, Cham. https://doi.org/10.1007/978-3-030-67910-1_17

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Potassium: a Vital Element in the Periodic Table and its Wide-Ranging Applications

This essay about potassium highlights its significance on the periodic table and its applications across various fields. It details the discovery of potassium by Sir Humphry Davy and explains its reactive nature and storage methods. The essay also explores potassium’s critical biological roles, particularly in regulating heart functions and maintaining fluid balance within the human body. Additionally, it discusses the element’s industrial uses, such as in fertilizers and chemical compounds like potassium permanganate, which are essential for agriculture and water treatment. The piece concludes by examining potential future applications of potassium in renewable energy storage, emphasizing ongoing research into potassium-ion batteries as a cost-effective alternative to lithium-ion technologies. This overview of potassium’s properties and uses underscores its versatility and vital role in both natural and human-engineered processes.

How it works

Potassium, an argentaceous metal denoted by the symbol ‘K’ on the periodic chart, bears substantial significance within the scientific realm and quotidian existence. Unearthed in 1807 by Sir Humphry Davy via the electrolysis of potash, potassium ranks as the seventh most plentiful element in the Earth’s lithosphere and assumes pivotal functions in myriad biological, industrial, and ecological mechanisms.

In its natural state, potassium eludes detection in its elemental guise due to its heightened reactivity with aqueous substances. Instead, it manifests in bonded states such as potash (from whence it derives its nomenclature) and assorted minerals.

This reactivity also underlies the rationale for encasing pure potassium within mineral oil or kerosene—its vigorous reaction with water precipitates the generation of potassium hydroxide and hydrogen gas, an exothermic phenomenon of considerable renown.

Biologically, potassium serves as a linchpin for vitality. It stands as a principal electrolyte within the human corpus, governing neural transmissions and muscular contractions, whilst preserving hydric equilibrium. Among its salient functions lies its pivotal role within the cardiovascular network, where it orchestrates cardiac rhythms. Sustaining a harmonious potassium equilibrium assumes paramount importance; an insufficiency may culminate in hypokalemia, engendering muscle enfeeblement, spasmodic convulsions, and lethargy, while an excess, termed hyperkalemia, poses commensurate hazards.

Moreover, potassium boasts significant industrial utility. Its deployment in fertilizers as potassium nitrate is indispensable for fostering vegetal proliferation by optimizing water utilization and fortifying resistance to aridity. This renders potassium pivotal within agrarian regimes, particularly in regions susceptible to water paucity. Additionally, potassium assumes an integral role within numerous pivotal compounds within the chemical industry, exemplified by potassium permanganate—a substance leveraged for disinfection and water purification.

Furthermore, potassium compounds wield considerable sway in environmental stewardship. Potassium carbonate finds application in glass and soap fabrication, whilst potassium hydroxide features prominently in biodiesel production. Intriguingly, the historical employment of potassium within soap, a tradition tracing its roots to antiquity, underscores its enduring significance within human civilization.

The future trajectory of potassium portends a trajectory as dynamic as its chemical constitution. Ongoing inquiry is exploring its efficacy within more efficient and ecologically benign fertilizers, alongside its prospective in next-generation energy storage modalities, exemplified by potassium-ion batteries, which proffer an alternative to conventional lithium-ion counterparts. These batteries loom as prospective game-changers within the realm of renewable energy storage, buoyed by potassium’s abundance and reduced cost.

In summation, potassium transcends its status as a mere constituent of the periodic schema. Its multifarious applications spanning diverse realms—from human health to industrial praxis and its potential within sustainable energy—underscore its versatility and indispensability. As the quest for knowledge regarding this intriguing element continues, the panorama of possibilities it presents is poised to burgeon, potentially ushering in novel innovations and ameliorations in technology and quality of life. This foray into the realm of potassium offers a glimpse into its intricacies and the substantial imprint it leaves upon our existence and milieu.

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An essay on periodic tables

After a compact history of the PT, from Döbereiner’s triads to the theoretical predictions up to element 172, a number of particular issues is discussed: Why may Z  = 172 be a limit for stable electron shells? What are the expected stability limits of the nuclear isotopes? When are formally empty atomic orbitals used in molecular electronic structures? What is ‘Secondary Periodicity’? When do the elements (Ir, Pt, Au), at the end of a bond, simulate (N, O, I), respectively? Some new suggestions for alternative PTs are commented upon. As a local connection, Johan Gadolin’s 1794 analysis of the Ytterby mineral is mentioned.

Historical introduction

A periodic table of elements (PT) arranges chemical elements as a function of their properties – how so? Any student might answer: by their nuclear charge, Z . Currently the elements with Z =1–118 have been found in nature or artificially produced. What did people use in the 19 th century, before they knew about nuclei or nuclear charges? Atomic weights, m . Except for a few anomalies, arrangement by m gave the same running order towards heavier elements, as by Z .

History in a nutshell . The periodic behaviour of chemical elements slowly became apparent in fragments, perhaps beginning with Döbereiner’s triads , such as (Ca, Sr, Ba) in 1817 or (Li, Na, K), (S, Se, Te) and (Cl, Br. I) in 1829. In each of these triads, the m of the second element is approximately equal to the average of the first and third.

In 1843, Gmelin had a table of 55 elements, with oxygen in the correct place. Following the spiralized ‘telluric screw’ (de Chancourtois 1862) and the Law of ‘octaves’ (Newlands 1863, 1865), Meyer constructed a square table of 28 elements (with gaps) in 1864. In 1869, Mendeleev wrote two articles (one in Russian and the other – a short summary – in German) explicitly predicting the existence of three missing elements with the atomic weights 45, 68 and 70. These were discovered in 1879, 1875 and 1886, and are now known as scandium (Sc; atomic weight 44.956), gallium (Ga; atomic weight 69.723) and germanium (Ge; atomic weight 72.640), respectively.

Mendeleev’s articles (1869) also overlapped with the writing of his textbook Fundamentals of Chemistry (see Kaji [ 1 ]).

In 1900 Ramsay suggested that the new (nearly) noble gases should form a separate group, which is currently referred to as Group 18. Similarly, in 1945, Seaborg proposed that the newly discovered actinides should form their own row below the lanthanides. He purposely chose the elemental names europium, americium (Eu, Am), gadolinium, curium (Gd, Cm), and terbium, berkelium (Tb, Bk) to emphasize the (4f, 5f) analogy by selecting for these actinides names of a continent, a celebrated scientist, and a town.

Accelerator experiments have now completed the 6d series Rf–Cn and the 7p series Nh–Og, resulting in the currently accepted PT (highlighted in yellow in Fig. 1 ).

The present Periodic Table (yellow) and possible assignments of the future elements E119–E172 (white). Picture reproduced from Haba [ 2 ]. Table reproduced from Pyykkö [ 3 ], [ 4 ]. Note the p-orbital spin-orbit-induced anomalies at E139–140 and E167–168, and the 9s-orbital-induced location of E165–166.

The theoretical predictions (shown in white in Fig. 1 ) by the present author [ 3 ], [ 4 ] support the idea of two 8 s elements (E119 and 120), an overlapping ‘grey’ zone of additional shells (8p, 7d, and 6f) at E121–124, and then a systematic sequence of increasing 5g occupation numbers for the elements E125 onwards. Nominally, all of the elements E121–138 are assigned to a ‘5g’ series. As a parallel case, recall that we regard Th as an actinide, although the free thorium atom has no 5f orbital occupation. Accordingly, a name pre-f was recently introduced [ 5 ].

As a further data point supporting Fig. 1 , Indelicato et al. [ 6a ] find an 8s 2 8p 2 5g 18 configuration for the cations E143 3+ to E148 8+ . The atomic calculations on the superheavy elements (SHE) range from the Dirac-Slater (DS) ones by Fricke et al. [ 6b ], and the Dirac-Fock (DF) ones by Desclaux [ 6c ] to massive multiconfiguration DF (MCDF) approaches, such as [ 6a ], and sophisticated many-body approaches [ 6d ], [ 6e ]. As an ultimate example, see [ 35 ] below. As to molecules, theoretical calculations on hypothetical octahedral hexafluorides, MF 6 , support the onset of the expected 5g occupation from E125 to at least E129. Note that the metal atom, M, delivers six electrons to the six fluorides. The other valence electrons go to the 5g shell. These molecular relativistic density-functional calculations were reported by Dognon and Pyykkö [ 7 ] [1]

For more comprehensive treatises on the history of the PT, see Gordin [ 8 ], Kaji [ 1 ] or Scerri [ 9 ].

Technical details

Why must z be ≤172.

The current border of the PT at Z =118 is set by the nuclear instability of the existing isotopes and by their small nucleosynthetic cross sections. For discussing chemical properties, even if the nuclei existed, there may also be limits, arising from the chemical reactivity of the vacuum in strong Coulomb fields, due to quantum electrodynamics, QED.

For any elements beyond E172, or so, the lowest or 1s shell would dive to the lower, positron-like continuum of the Dirac equation. It is not yet fully understood what would physically happen. Another way to study the question is to consider heavy-atom collisions [ 10 ]. For one, point-like nucleus this diving would already take place at Z =137. For the earlier literature on this question, see [ 3 ], p. 162.

Relativity vs. QED

It is well-known that the (Dirac) relativistic effects contract and stabilize the ns and np* (=np 1/2 ) shells while the ensuing indirect relativistic effects expand and destabilize the d and f shells. As previously discussed [ 5 ], the QED effects, dominated by the vacuum fluctuations (the zero-point oscillations of the electromagnetic field), cancel about −1per cent of the previous effect, for the heavier elements. The other lowest-order contribution, of opposite sign, is vacuum polarization. One could say that the Dirac-Fock-Breit Hamiltonian is ‘101 percent correct’ (cp. [ 6a ]). The QED effects can be seen in accurate quantitative comparisons but have so far not led to qualitative chemical changes.

Which orbitals to use in chemistry?

The chemical behaviour of the elements in Fig. 1 is mostly driven by the orbitals, occupied in the atomic ground state, and given in the right-hand marginal. Sometimes also other orbitals, which are unoccupied in the atomic ground state but energetically accessible for bond formation, can participate. Thus we can have the predicted [ 11 ] pre-s Og − anion, the pre-p Be, Mg; Zn, making bonds with their ns+np orbitals, the pre-d Ca, Sr, Ba; Cs and the pre-f Th.

A recent example was the synthesis of [Ba(CO) 8 ], which fulfils an 18-electron rule by using the originally empty 5d shell of the central barium atom [ 12 ]. The atomic ground state does not always explain the molecular outcome. Moreover, remember that electron correlation can make the concept of electron configurations diffuse.

‘Secondary periodicity’

Biron [ 13 ] pointed out in 1915 that every second period has specific properties. Taking the Group 15 as an example, along the series (N, P, As, Sb, Bi) (Biron, p. 971), the dominant oxidation states are III, V, III, V, III, respectively. The anomaly at As can be attributed to partial screening by the filled 3d shell. The anomaly at Bi is due to both an analogous partial screening by the 4f shell and to relativistic effects [ 14 ].

Another vertical anomaly is the small radius of every atomic shell (1s, 2p, 3d, 4f, 5g) with a new orbital angular momentum quantum number, l . The author [ 15 ] used the name ‘primogenic repulsion’ for its effect on the higher shells. In Russian literature the term ‘kainosymmetric’ is often used for these ‘first-born’ shells [ 16 ].

The inert-pair effect

Sidgwick [ 17a ] called attention in 1933 to a decrease of the main oxidation state by two units for 6 th -period elements, take Pb(II) as an example. A first explanation would be the relativistic 6s stabilization. Closer studies involve the hybridization of the metal (6s, 6p) orbitals with the ligand np orbitals [ 17b ].

‘False friends’

The gold atom is almost as electronegative as iodine; we can see its outermost shell as either a 6s 1 electron, or a 6s −1 hole. A wide chemistry of the auride ion, Au − , is known [ 18 ]. For a comparison of aurides with other ‘halides’, see also [ 19 ].

Going one step left from gold, solid Cs 2 Pt and other Pt (−II) compounds were studied in the group of Jansen [ 20 ]. In molecules, in addition to σ bonding, also analogous 2pπ and 5dπ bonding was identified between OCO and PtCO, respectively, leading to multiple bonding [ 21 ].

In the uranyl-like isoelectronic series, OUN + and OUIr + were found to have similar triple bonds [ 22 ] and the latter species was later produced in mass-spectroscopy [ 23 ]. These later chemical analogies were initially unexpected.

Nuclear stability

The chemical predictions quoted here are based on theoretical, relativistic quantum chemical calculations using established electronic Hamiltonians. The nuclei are simply assumed to exist, with a realistic, finite nuclear size. The synthesis of heavier nuclei, up to E118 (Oganesson), is demanding. The most recent nuclear syntheses were completed in a friendly collaboration between laboratories in Oak Ridge and Dubna. The lifetimes of these nuclides are short; for example, the present Og isotopes have lifetimes below a millisecond. The most challenging production bottleneck is, however, not the short lifetime but the small nucleosynthetic cross-section for these elements. If the experiment runs for a year and yields less than a handful of desired product nuclei, this creates an obvious problem, even with nearly 4π detection (all scattering directions seen) and an almost noiseless apparatus. The current situation on superheavy elements is discussed by Giuliani et al. [ 24 ].

A quantum chemist can always assume a finite nucleus of realistic size for any Z <173 and do ab initio calculations for that theoretical model, whether or not such nuclei or their compounds are ever made. One even could claim that some general conclusions could possibly be drawn, such as the possible existence of a 5g series or the vast spin-orbit effects.

In defence of the current PT: what are the choices?

Figure 1 illustrates a recent choice of IUPAC PT layout (highlighted in yellow). Note the placement of all lanthanides and actinides in Group 3. One could, however argue certain points:

H is now in Group 1, because it is often manifested as H +1 , or neutral. If an emphasis was placed on hydrides (H −1 ), one could argue for having hydrogen also in Group 17.

The noble gas He is now in Group 18 with the other noble gases. Apart from spectroscopic species, like the astrophysical, diatomic HeH + , or high-pressure compounds, like Na 2 He [ 25 ], helium is still a very noble gas. As stressed by the authors, this compound is a sodium electride containing He(0). If one would like to emphasize heliums 1s 2 electronic structure, one could also have it in Group 2. Its chemical behaviour is, however, not that of an alkaline earth. Note the easy ns-np-(n-1)d hybridization among heavier Group-2 elements, lacking for helium.

How long f-element rows? One now has a 15-element lanthanide (Ln) row from La to Lu. All of these elements are (mostly) trivalent. Their ionic radii or ionization potentials exhibit a systematic relationship along the series. It is entirely plausible to count from 4f 0 to 4f 14 , and to leave a hole in Group 3 of Period 6. Moreover this completely avoids the heated argument on which end should one cut off – La or Lu. A clear advantage is then having all these, mostly trivalent, rare earths in the single Group 3, corresponding to three valence electrons.

Impressive experiments and computations [ 26 ] have recently verified that a free Lr atom has a 7s 2 7p 1 configuration, different from the 6s 2 5d 1 for a Lu atom. Computations for a handful of molecules , however, find a complete analogy between Lu and Lr [ 27 ], see Fig. 2 . That said, if our PT is to be driven by chemistry, there are no reasons to change Fig. 1 .

The 32-column option? Some desirable properties of a PT could be:

The column (or ‘Group’) number g (or g-10 ) is the maximum number of valence electrons. The oxidation state is then the same number or, counting holes in the spirit of Abegg’s contravalence [ 29 ], the negative number, g-18.

The valence atomic orbitals along a row are constant in a block and equal to those in the right-hand margin.

The nuclear charge, Z , increases systematically towards right.

The shape should be typographically convenient.

The valence molecular orbitals (MO) of theoretical LuCO and LrCO molecules [ 27 ]. Note the similarity. The LuCO is experimentally known, see [ 28 ]. As discussed there, the main Ln bonding orbitals are the 6s and 5d. The MO:s from left to right are the M–C pi bond (a donation from M to the CO pi* MO), a sigma lone pair, and the M–C sigma bond. The similarity between Lu and Lr is not a mistake, but the message.

Of these properties, Fig. 1 fulfils 1, 2 and 4 but violates 3 for certain superheavy elements when relativistic effects so require. Conversely, the 32-column ‘long-form’ PT favoured by Scerri [ 30 ] (see his Fig. 1 ) violates (1) by having very many potentially Group-3 columns and may also violate (4). It does satisfy (3).

The ‘Madelung rule’

The order of filling shells in a neutral atom is approximately the one given in Fig. 3 . It was extended up to Z =172 in [ 4 ]. This mnemonic device is usually called the Madelung or ( n+l , n ) rule [ 31 ], although its shape appears to be first presented by Janet [ 32a ] or actually Sommerfeld [ 32c ].

The ‘Madelung rule’ for filling atomic orbitals up to Z =172, corresponding to Fig. 1 (reproduced from Pyykkö [ 4 ]).

Models for reproducing the PT

How to explain this approximate order of level filling in neutral, or nearly neutral atoms? For an electron in a Coulomb field, each new n introduces a new l max and this degeneracy of all levels with the same n was discussed by Fock [ 33 ] using momentum-space wave functions in a four-dimensional space. Ostrovsky [ 34 ] used coordinate space.

Concerning the physics of many-electron atoms, the Dirac-Fock-Breit (DFB) Hamiltonian, supplemented with some estimate of leading quantum electrodynamic (QED) effects, gives an excellent description. As an example, Pašteka et al. [ 35 ] calculated the ionization potential and electron affinity of a gold atom with milli-electronvolt (meV) accuracy. The bottleneck rather was in the handling of electron correlation. Coupled-cluster methods with up to pentuple excitations were used. Thus no surprises are expected here. Atoms follow Physics.

A much simpler task is the question of the filling order of one-electron levels in some effective-potential model for a many-electron atom, such as a Thomas-Fermi one. This has been tried since Fermi [ 36 ] or Goeppert Mayer [ 37 ], and various versions have been included in textbooks, such as Sommerfeld [ 38 ], Gombás [ 39 ], or Landau and Lifshitz [ 40 ]. Some examples on these studies are [ 41 ], [ 42 ]. The predictions for filling new shells are quite similar, see Table 1 . In a screened-Coulomb potential, the attraction must be sufficient to balance the centrifugal potential for l . This type of reasoning was used by Goeppert Mayer [ 37 ] to discuss the Z where 4f and 5f states are first occupied. For reviews see Ostrovsky below. The T–F treatment yields the Z values for a new l at

The first nuclear charges, Z , where the atomic orbitals l are occupied for a Schrödinger equation in a Thomas-Fermi potential (from Fermi [ 36 ], Goeppert Mayer [ 37 ], Iwanenko and Larin [ 41 ], Landau and Lifshitz [ 39 ] [present eq. (1)], Essén [ 43 ]) and from Dirac-Fock-level relativistic calculations [ 3 ].

with the pre-coefficient chosen by Landau and Lifshitz [ 40 ]. The results in Table 1 are in a surprisingly close agreement with more exact results. Note that the T–F potential is just one of the various screened-Coulomb potentials, which do the job. The literature on which T–F potential at which Z starts a given l is broad, see Table 1 . Essén [ 43 ] gives for filling the first state of l the nuclear charge

The doubling of periods with two periods of lengths 8, 18 and 32 each has been noticed by many authors, Löwdin [ 44 ], Demkov and Ostrovsky [ 45 ], Odabasi [ 46 ], Ostrovsky [ 47 ], Essén [ 43 ], Katriel [ 48 ], Kitagawara and Barut [ 49 ], Novaro [ 50 ] and Kibler [ 51 ] have discussed possible underlying dynamical symmetries. An alternative is that there is no deep symmetry reason and that one only needs the quoted ‘nuts and bolts’ of the DFB Hamiltonian.

For later reviews on the doubling question, see Ostrovsky [ 52 ].

Concluding , of the existing literature on the PT, we would still like to remind the reader of the articles by Schwarz [ 53 ], [ 54 ], [ 55 ] and of a factor-analytical search for chemical similarities by Leal and Restrepo [ 56 ].

Did our department contribute to the story?

The new Chemical Laboratory building of the old Kungliga Åbo Akademi (our direct legal predecessor before the removal from Turku to Helsinki in 1828) was inaugurated on 13 April 1764 under Professor Pehr Gadd. The 250 th anniversary was celebrated in 2014 by having business as usual.

Gadd’s successor, Johan Gadolin (1760–1852) [ 57 ] published the first chemical analysis of a black mineral from Ytterby (near Stockholm) in 1794 (German version published in 1796). The paper was translated to modern chemical terminology by Pyykkö and Orama [ 58 ].

Starting from this mineral, which was essentially FeBe 2 Y 2 Si 2 O 10 , he obtained a new ‘earth’, i.e. oxide, essentially Y 2 O 3 , possibly with traces of other rare earths. The mineral was later named gadolinite . Although one cannot necessarily claim that Gadolin found a specific element, he did contribute to finding an entire family of 16 elements. Organizing these kept chemists busy for another century. In 1886 one of them was given the name gadolinium ( Gd ).

Circumstantial evidence from the parallel case of samarium (Sm) by the same authors suggests that Lecoq de Boisbaudran and Marignac considered both the mineral and the eponym, or the person behind its name. For more details on Gd, see [ 59 ] and supplements to it. In 1994 a bicentennial conference, 2-ICFE, was held in Helsinki, with a conference excursion to Ytterby, see Fig. 4 .

In 1994 an International Conference on f-Elements was held in Helsinki to celebrate the bicentennial of Johan Gadolin’s 1794 analysis [ 57 ] [2] .

Article note

A collection of invited papers based on presentations at “Mendeleev 150”: 4 th International Conference on the Periodic Table (Mendeleev 150), held at ITMO University in Saint Petersburg, Russian Federation, 26–28 July 2019.

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The universe's biggest explosions made elements we are composed of, but there's another mystery source out there

A fter its "birth" in the Big Bang, the universe consisted mainly of hydrogen and a few helium atoms. These are the lightest elements in the periodic table. More-or-less all elements heavier than helium were produced in the 13.8 billion years between the Big Bang and the present day.

Stars have produced many of these heavier elements through the process of nuclear fusion. However, this only makes elements as heavy as iron. The creation of any heavier elements would consume energy instead of releasing it.

In order to explain the presence of these heavier elements today, it's necessary to find phenomena that can produce them. One type of event that fits the bill is a gamma-ray burst (GRB) —the most powerful class of explosion in the universe. These can erupt with a quintillion (10 followed by 18 zeros) times the luminosity of our sun, and are thought to be caused by several types of event.

GRBs can be subdivided into two categories: long bursts and short bursts. Long GRBs are associated with the deaths of massive and fast-rotating stars. According to this theory, the fast rotation beams material ejected during the collapse of a massive star into narrow jets that move at extremely fast speeds.

The short bursts last only a few seconds. They are thought to be caused by the collision of two neutron stars—compact and dense "dead" stars. In August 2017, an important event helped support this theory. Ligo and Virgo , two gravitational wave detectors in the US, discovered a signal that seemed to be coming from two neutron stars moving in for a collision.

A few seconds later, a short gamma-ray burst, known as GRB 100817A, was detected coming from the same direction in the sky. For a few weeks, virtually every telescope on the planet was pointing at this event in an unprecedented effort to study its aftermath.

The observations revealed a kilonova at the location of GRB 170817A. A kilonova is a fainter cousin of a supernova explosion. More interestingly, there was evidence that many heavy elements were produced during the explosion . The authors of a study in Nature that analyzed the explosion showed that this kilonova seemed to produce two different categories of debris, or ejecta. One was composed primarily of light elements, while another consisted of heavy elements.

We've already mentioned that nuclear fission can only feasibly produce elements as heavy as iron in the periodic table. But there's another process which could explain how the kilonova was able to produce even heavier ones.

Rapid neutron-capture process , or r-process, is where the nuclei (or cores) of heavier elements such as iron capture many neutron particles in a short time. They then rapidly grow in mass, yielding much heavier elements. For r-process to work, however, you need the right conditions: high density, high temperature, and a large number of available free neutrons. Gamma ray bursts happen to provide these necessary conditions.

However, mergers of two neutron stars, like the one that caused the kilonova GRB 170817A, are very rare events. In fact, they may be so rare as to make them an unlikely source for the abundant heavy elements we have in the universe. But what of long GRBs?

A recent study investigated one long gamma ray burst in particular, GRB 221009. This has been dubbed the BOAT —the brightest of all time. This GRB was picked up as a pulse of intense radiation sweeping through the solar system on October 9 2022.

The BOAT sparked a similar astronomical observation campaign as the kilonova. This GRB was 10 times more energetic then the previous record holder, and so close to us that its influence on the Earth's atmosphere was measurable on the ground and comparable to a major solar storm.

Among the telescopes studying the aftermath of the BOAT was the James Webb Space Telescope (JWST). It observed the GRB about six months after it exploded, so as not to be blinded by the afterglow of the initial burst. The data JWST collected showed that, despite the event's extraordinary brightness, it was caused by a merely average supernova explosion .

In fact, previous observations of other long GRBs indicated that there is no correlation between the brightness of the GRB and the size of the supernova explosion associated with it. The BOAT seems no exception.

The JWST team also inferred the number of heavy elements produced during the BOAT explosion. They found no indication of elements produced by the r-process. This is surprising as, theoretically, the brightness of a long GRB is thought to be associated with the conditions in its core, most likely a black hole. For very bright events –- especially one as extreme as the BOAT –- the conditions should be right for the r-process to occur.

These findings suggest that gamma ray bursts may not be the hoped-for crucial source of the universe's heavy elements. Instead, there must be a source or sources still out there.

This article is republished from The Conversation under a Creative Commons license. Read the original article .

Provided by The Conversation

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