Elements| Periodic Table of Elements, Properties & Descriptions

Elements| Periodic Table of Elements

Elements

Elements are the main constituents of all types of matter exist. Since the beginning of time, scientists have been discovering new chemical elements. Some elements, like gold, were first identified thousands of years ago because they may be found in nature in their elemental form.

In contrast, some elements have inherent instability and radioactivity. Despite being stable, the majority of the elements are extensively distributed in nature and included in a variety of compounds. Because of this, scientists were ignorant about them for a very long time. There were not many known elements at the start of the seventeenth century. Chemistry advancements in the eighteenth and early nineteenth centuries made it simpler to separate elements from their complexes. As a result, from 31 in 1800 to 63 in 1865, the number of known elements has more than doubled. The International Union of Pure and Applied Chemistry had approved 118 elements as of November 2016.

The first 94 elements are found naturally on Earth, whereas the final 24 are created artificially through nuclear processes. Nearly all of the elements are accessible industrially in varied quantities, with the exception of unstable radioactive elements (radionuclides) that disintegrate swiftly. Science is always researching new elements and developing new ways to create them.

118 elements* are currently recognized. Recent discoveries among these are synthetic, or man-made, elements, and efforts are still being undertaken to synthesize additional elements through artificial transmutation. As the number of known elements expanded, it became increasingly difficult to examine the chemistry of each one separately, and researchers started looking into the possibility of grouping these elements in useful ways.

The fundamental objective and prerequisite of classification, in the words of Huxley, is "the actual or ideal arrangement together of that which is like and separation of that which is unlike; the purpose of this arrangement being primarily to disclose the correlations, or laws of union, or properties, or circumstances, and secondarily to facilitate the operation of the mind in dearly receiving and then retaining in memory the characteristics of the objects in quest." In a nutshell, the main goal of categorization is to organise information about elements and their compounds in a way that will give us the most influence over their properties with the least amount of work. The best classification would group together components that are most similar to one another and separate others.

Description of Elements

The lightest chemical elements are hydrogen and helium, which were both produced during the first 20 minutes of the universe by the process known as Big Bang nucleosynthesis in a ratio of about 3:1 by mass (or 12:1 by number of atoms), along with minute amounts of the next two elements, lithium and beryllium. Nearly every other element in nature was created using different natural processes for nucleosynthesis. On Earth, cosmogenic activities like cosmic ray spallation or nucleogenic reactions naturally produce modest amounts of new atoms. On Earth, radiogenic daughter isotopes of continuous radioactive decay processes like alpha decay, beta decay, spontaneous fission, cluster decay, and other more uncommon forms of decay also naturally make new atoms.

The 94 naturally occurring elements have at least one stable isotope, with atomic numbers 1 through 82 having the most (except for technetium, element 43 and promethium, element 61, which have no stable isotopes). When there hasn't been any evidence of radioactive decay for an isotope, it is said to be stable.

All radioactive decay of the elements with atomic numbers 83 through 94 can be observed since they are so unstable. Some of these elements, most notably bismuth (atomic number 83), thorium (atomic number 90), and uranium (atomic number 92), have one or more isotopes with long enough half-lives to survive as leftovers from the frenzied stellar nucleosynthesis that created the heavy metals prior to the formation of our Solar System. Bismuth-209 (atomic number 83) has the longest known alpha decay half-life of any naturally occurring element at over 1.910¹⁹ years, more than a billion times longer than the current estimate of the age of the universe. It is frequently compared to the other 80 stable elements because of this.

The heaviest elements (those heavier than element 94, plutonium), which undergo radioactive decay, have half-lives that are so brief that they cannot be found in nature and must be created.

Currently, 118 elements are known. "Known" in this context refers to having been distinguished from other elements after being sufficiently observed, even from a small number of decay products. Most recently, the synthesis of element 118 (now known as oganesson) and element 117 (tennessine) were reported in October 2006 and April 2010, respectively. 94 of these 118 elements are naturally present on Earth. Technetium, atomic number 43; promethium, number 61; astatine, number 85; francium, number 87; neptunium, number 93; and plutonium, number 94—six of these—occur in extremely minuscule amounts. These 94 substances have been found throughout the cosmos, in the spectra of stars and supernovae, where new short-lived radioactive substances are being created. The first 94 elements have been found directly on Earth as either naturally occurring fission or transmutation products of uranium and thorium, or as primordial nuclides left over from the Solar System's creation.

The remaining 24 heavier elements, which are all radioactive and have very short half-lives and are not currently found on Earth or in astronomical spectra, have all been created artificially. If any atoms of these elements were present during the formation of Earth, they are almost certainly to the point of certainty already decayed, and if present in novae, they were in quantities too small to have been noticed. In 1937, technetium was supposedly created for the first time artificially, however small amounts have subsequently been discovered in nature (and also the element may have been discovered naturally in 1925). With numerous additional radioactive naturally occurring rare elements, the process of artificial creation and subsequent natural discovery has been repeated.

There is a list of the elements that includes the element's name, atomic number, density, melting point, boiling temperature, symbol, and ionization energy. The stable and radioactive elements' nuclides are also available as a list of nuclides, arranged by the stability of the unstable ones' half-lives. The periodic chart, which groups together elements with similar chemical properties, is one of the most practical and undoubtedly the oldest way to present the elements (and usually also similar electronic structures).

Periodic Table of Elements

There have been attempts to categorize the elements based on various characteristics. The elements were divided into electropositive and electronegative groups as well as metals and nonmetals. They were categorised as monovalent, divalent, trivalent, etc. based on their valency. Since the same material frequently featured in multiple classes in these early classifications, they were all thought of as crude categories. Following the development of Dalton's atomic theory, chemists began to look for a correlation between the properties of the elements and their atomic masses, treating atomic mass as the essential characteristic of the element. On the basis of the atomic masses of the elements, various classifications were offered. The periodic table, which lists elements in a tabular format according to their properties, was the best formulation. A chemical family is made up of components from the same vertical column that share characteristics.

The periodic table is made up of elements arranged in vertical columns according to increasing atomic masses and comparable properties. The study of such a great number of elements has been made incredibly simple by the periodic classification (long or extended version of periodic table) of elements.

Periodic Table

Atomic Number

An element's atomic number, which corresponds to the number of protons in each atom, identifies the substance. For instance, the atomic number of Carbon is 6 because all carbon atoms have six protons in their atomic nucleus. The number of neutrons in a carbon atom can vary; atoms of the same element with differing numbers of neutrons are referred to as the element's isotopes.

In addition to determining the atomic nucleus's electric charge, the number of protons in an atom also controls how many electrons it has when it is not ionised. The atomic orbitals in which the electrons are positioned define the numerous chemical characteristics of the atom. The chemical characteristics of an element are typically not greatly influenced by the number of neutrons in its nucleus (except in the case of hydrogen and deuterium). As a result, despite the fact that carbon atoms, for instance, can have 6 or 8 neutrons, all carbon isotopes have almost identical chemical properties since they all have six protons and six electrons. For this reason, the atomic number—and not the mass number or atomic weight—is regarded as the distinguishing feature of a chemical element.

The symbol for atomic number is Z.

Isotopes

Isotopes are atoms of the same element that have variable numbers of neutrons despite having the same number of protons in their atomic nucleus. There are three primary isotopes of carbon, for instance. There are 6 protons in the nucleus of every carbon atom, however the neutrons can be 6, 7, or 8. The three carbon isotopes are referred to as carbon-12, carbon-13, and carbon-14, or ¹²C, ¹³C, and ¹⁴C, respectively, due to their respective mass numbers of 12, 13, and 14. In both daily life and chemistry, carbon is made up of a mixture of ¹²C (98.9%), ¹³C (1.1%), and roughly 1 atom per trillion of ¹⁴C.

Most naturally occurring elements (66 out of 94) contain several stable isotopes. The isotopes of a given element are essentially identical chemically, with the exception of hydrogen, whose isotopes range significantly from one another in relative mass and can therefore have an impact on chemical reactions.

There are radioactive isotopes of all the elements, albeit not all of these radioisotopes are found naturally. When a radioisotope emits an alpha or beta particle, it normally decays into another element. Stable isotopes are those that an element has that are not radioactive. The known stable isotopes are all produced naturally. After being created intentionally, the numerous radioisotopes that are absent from nature have been characterised. The elements without any stable isotopes are technetium (atomic number 43), promethium (atomic number 61), and all observed elements with atomic numbers more than 82. These elements are entirely made up of radioactive isotopes.

26 of the 80 elements that have at least one stable isotope each have just one. For the 80 stable elements, there are an average of 3.1 stable isotopes per element. There are 10 stable isotopes, which is the most for any one element (for tin, element 50).

Isotopic Mass and Atomic Mass of Elements

The quantity of nucleons (protons and neutrons) in an element's atomic nucleus, denoted by the letter A, is its mass number. The mass numbers of an element's many isotopes, which are often expressed as a superscript on the left side of the atomic symbol, serve as a means of identification (e.g. ²³⁸U). The mass number has "nucleons" as its units and is always a whole number. For instance, the atom of magnesium-24, whose mass number is 24, has 24 nucleons (12 protons and 12 neutrons).

The atomic mass of a single atom is a real number that indicates the mass of a specific isotope (or "nuclide") of the element, expressed in atomic mass units, as opposed to the mass number, which simply counts the total number of neutrons and protons and is therefore a natural (or whole) number (symbol: u). Since the mass of each proton and neutron is not exactly 1 u, the electrons' contribution to the atomic mass decreases as the number of neutrons increases, and (last but not least) because of the nuclear binding energy, the mass number of a given nuclide generally deviates slightly from its atomic mass. For instance, the atomic mass of chlorine-37 is 36.966 u while that of chlorine-35 is 34.969 u to five significant digits. However, each isotope's atomic mass in u is consistently within 1% of its simple mass number. Because u is determined to be 1/12 of the mass of a free neutral carbon-12 atom in its ground state, ¹²C is the only isotope whose atomic mass is precisely a natural number. As a result, ¹²C isotope has an exact mass of 12 by definition.

The average of the atomic masses of all the chemical element's isotopes as found in a specific environment, weighted by isotopic abundance, in relation to the atomic mass unit, is the standard atomic weight (often referred to as "atomic weight") of an element. It's possible that this number is a fraction that is not very close to a whole number. In contrast to a whole number, the relative atomic mass of chlorine, for instance, is 35.453 u, which is an average of around 76% chlorine-35 and 24% chlorine-37. Because a sample of an element naturally contains sizable amounts of more than one isotope, this averaging effect is always to blame when a relative atomic mass value deviates by more than 1% from a whole number.

Chemically Pure and Isotopically Pure Elements

Nuclear scientists and chemists have differing ideas of what a pure element is. In chemistry, a pure element is a substance whose atoms, or almost all of them, have the same atomic number, or number of protons.

A copper wire, for instance, is 99.99% chemically pure if 99.99% of its atoms contain 29 protons of copper. However, it is not isotopically pure since common copper has two stable isotopes with differing neutron ratios: 69% 63Cu and 31% 65Cu. A pure gold ingot would be both chemically and isotopically pure because typical gold only includes the isotope 197Au.

Allotropes

Chemically pure elements can be found in a variety of chemical configurations (spatial arrangements of atoms), known as allotropes, that have different properties because the atoms of those elements can link to one another chemically in more than one way. Graphite, which has layers of carbon atoms with a hexagonal structure stacked on top of each other, graphene, which is a single layer of extremely strong graphite, fullerenes, which have nearly spherical shapes, and carbon nanotubes, which are tubes with a hexagonal structure, are a few examples of materials that contain carbon (even these may differ from each other in electrical properties). Allotropy is the capacity of an element to exist in one of several structural configurations.

The definition of an element's standard state, sometimes referred to as the reference state, is its thermodynamically most stable condition under a pressure of 1 bar and a specific temperature (typically at 298.15K). An element is considered to have an enthalpy of production of zero in its standard state in thermochemistry. For instance, graphite is the reference state for carbon because it has a more stable structure than the other allotropes.

Properties of Elements

The elements can be widely categorized using a variety of descriptive criteria, such as general physical and chemical characteristics, states of matter under common circumstances, melting and boiling points, densities, crystal structures as solids, and origins.

General properties

The general physical and chemical characteristics of the chemical elements are described using a variety of terminology. A first distinction is made between metals, which conduct electricity easily, nonmetals, which do not, and a tiny group of substances called metalloids, which have intermediate qualities and frequently behave as semiconductors.

In colourful depictions of the periodic table, a more precise classification is frequently displayed. This approach adds additional terminology for specific sets of the more broadly regarded metals and nonmetals, limiting the terms "metal" and "nonmetal" to only certain of the more broadly defined metals and nonmetals. Actinides, alkali metals, alkaline earth metals, halogens, lanthanides, transition metals, post-transition metals, metalloids, reactive nonmetals, and noble gases are included in the version of this classification used in the periodic tables shown below. The transition metals, lanthanides, and actinides are special groups of the metals in this system, together with the alkali metals, alkaline earth metals, and transition metals. The noble gases and reactive nonmetals are also considered nonmetals in a more general sense. The halogens are sometimes not distinguished, with astatine being classified as a metalloid and the others as nonmetals.

States of Matter

Another fundamental difference between the elements that is frequently made is their state of matter (phase), which can be either solid, liquid, or gas at a chosen standard temperature and pressure (STP). At normal temperatures and air pressure, the majority of the elements are solids, but some are gases. At 0 degrees Celsius (32 degrees Fahrenheit) with normal atmospheric pressure, only bromine and mercury are liquids; caesium and gallium are solids at that temperature but melt at 28.4 degrees Celsius (83.2 degrees Fahrenheit) and 29.8 degrees Celsius (85.6 degrees Fahrenheit), respectively.

Melting and boiling points

The melting and boiling points of various elements, usually stated in degrees Celsius at an atmospheric pressure, are frequently used to describe them. While one or both of these measures are known for the majority of elements, they are still unknown for some radioactive elements that are only found in very small amounts. In typical presentations, helium has merely a boiling point and not a melting point because it remains a liquid even at absolute zero at atmospheric pressure.

Densities

When describing an element, the density at a chosen standard temperature and pressure (STP) is usually utilised. The unit of measurement for density is grammes per cubic centimetre (g/cm3). Since several elements are gases at temperatures that are routinely encountered, their densities are typically reported for these gaseous states; when liquefied or solidified, the gaseous states of these elements have densities that are comparable to those of the other elements.

In summary presentations, one representative allotrope is often chosen when an element has allotropes with varied densities; densities for each allotrope can be described in more detail. As an illustration, the densities of the three well-known allotropes of carbon, amorphous carbon, graphite, and diamond, are 1.8-2.1, 2.267, and 3.515 g/cm3, respectively.

Crystal structures

Eight different types of crystal structures can be found in the elements that have been examined thus far as solid samples: cubic, body-centered cubic, face-centered cubic, hexagonal, monoclinic, orthorhombic, rhombohedral, and tetragonal. The available samples for several transuranic elements generated artificially have been too tiny to identify their crystal structures.

Occurrence and origin on Earth

The first 94 atomic numbers are thought to be naturally occurring chemical elements, while chemical elements with atomic numbers over 94 have only been created artificially as the synthetic byproducts of man-made nuclear reactions. Chemical elements can also be categorised by their origin on Earth.

Of the 94 elements that exist in nature, 83 are thought to be primordial and are either stable or mildly radioactive. The other 11 naturally occurring elements are categorised as transitory elements because their half lifetimes are too short for them to have existed at the start of the Solar System. The five most prevalent decay products of thorium and uranium, polonium, radon, radium, actinium, and protactinium, are among these 11 transitory elements. Technetium, promethium, astatine, francium, neptunium, and plutonium are the other six transitory elements, and they are only sporadic byproducts of uncommon nuclear reaction processes involving uranium or other heavy elements.

With the exception of 43 (technetium) and 61, no radioactive decay has been seen for elements with atomic numbers 1 through 82. (promethium). However, it is projected that some elements' observationally stable isotopes—like those of tungsten and lead—are mildly radioactive and have extremely long half-lives. For instance, the predicted half-lives for the observationally stable lead isotopes range from 10³⁵ to 10¹⁸⁹ years. The radioactive decay of elements having atomic numbers 43, 61, and 83 through 94 can be easily observed. Three of these elements, thorium (element 90), bismuth (element 83), and uranium (element 92), all have one or more isotopes with long enough half-lives to remain as byproducts of the frenzied star nucleosynthesis that created the heavy elements prior to the Solar System's birth. For instance, bismuth-209 has the longest known alpha decay half-life of any naturally occurring element, clocking in at more than 1.9x10¹⁹ years, more than a billion times longer than the universe's age as it is now calculated to be. Since the 24 heaviest elements (those heavier than element 94, plutonium), undergo radioactive decay with brief half-lives and cannot be created as daughters of longer-lived elements, it is not known that any of these elements exist in nature.

Conclusion.

Pure elements, bound compounds, or a combination of atoms and compounds make up all matter in the cosmos. The Periodic Table of Elements classifies all elements according to the number of protons in their atom's nucleus. There are elements in everything on the table.

We can conclude by saying that the periodic table is significant because of the way it is set up to present a wealth of knowledge about the elements and their interactions. Even for elements that have not yet been identified, the properties of those elements may be predicted using the periodic table.