Electron shell in the context of Closed shell

In chemistry and physics, valence electrons are electrons in the outermost shell of an atom, and that can participate in the formation of a chemical bond if the outermost shell is not closed. In a single covalent bond, a shared pair forms with both atoms in the bond each contributing one valence electron.

The presence of valence electrons can determine the element's chemical properties, such as its valence—whether it may bond with other elements and, if so, how readily and with how many. In this way, a given element's reactivity is highly dependent upon its electronic configuration. For a main-group element, a valence electron can exist only in the outermost electron shell; for a transition metal, a valence electron can also be in an inner shell.

In chemistry, a group (also known as a family) is a column of elements in the periodic table of the chemical elements. There are 18 numbered groups in the periodic table; the 14 f-block columns, between groups 2 and 3, are not numbered. The elements in a group have similar physical or chemical characteristics of the outermost electron shells of their atoms (i.e., the same core charge), because most chemical properties are dominated by the orbital location of the outermost electron.

The modern numbering system of "group 1" to "group 18" has been recommended by the International Union of Pure and Applied Chemistry (IUPAC) since 1988. The 1-18 system is based on each atom's s, p and d electrons beyond those in atoms of the preceding noble gas. Two older incompatible naming schemes can assign the same number to different groups depending on the system being used. The older schemes were used by the Chemical Abstract Service (CAS, more popular in the United States), and by IUPAC before 1988 (more popular in Europe). The system of eighteen groups is generally accepted by the chemistry community, but some dissent exists about membership of elements number 1 and 2 (hydrogen and helium). Similar variation on the inner transition metals continues to exist in textbooks, although the correct positioning has been known since 1948 and was twice endorsed by IUPAC in 1988 (together with the 1–18 numbering) and 2021.

In quantum mechanics, the azimuthal quantum number ℓ is a quantum number for an atomic orbital that determines its orbital angular momentum and describes aspects of the angular shape of the orbital. The azimuthal quantum number is the second of a set of quantum numbers that describe the unique quantum state of an electron (the others being the principal quantum number n, the magnetic quantum number m_ℓ, and the spin quantum number m_s).

For a given value of the principal quantum number n (electron shell), the possible values of ℓ are the integers from 0 to n − 1. For instance, the n = 1 shell has only orbitals with ${\displaystyle \ell =0}$, and the n = 2 shell has only orbitals with ${\displaystyle \ell =0}$, and ${\displaystyle \ell =1}$.

Electronegativity, symbolized as χ, is the tendency for an atom of a given chemical element to attract shared electrons (or electron density) when forming a chemical bond. An atom's electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus. The higher the associated electronegativity, the more an atom or a substituent group attracts electrons. Electronegativity serves as a simple way to quantitatively estimate the bond energy, and the sign and magnitude of a bond's chemical polarity, which characterizes a bond along the continuous scale from covalent to ionic bonding. The loosely defined term electropositivity is the opposite of electronegativity: it characterizes an element's tendency to donate valence electrons.

On the most basic level, electronegativity is determined by factors like the nuclear charge (the more protons an atom has, the more "pull" it will have on electrons) and the number and location of other electrons in the atomic shells (the more electrons an atom has, the farther from the nucleus the valence electrons will be, and as a result, the less positive charge they will experience—both because of their increased distance from the nucleus and because the other electrons in the lower energy core orbitals will act to shield the valence electrons from the positively charged nucleus).

In chemistry, the carbon–hydrogen bond (C−H bond) is a chemical bond between carbon and hydrogen atoms that can be found in many organic compounds. This bond is a covalent, single bond, meaning that carbon shares its outer valence electrons with up to four hydrogens. This completes both of their outer shells, making them stable.

Carbon–hydrogen bonds have a bond length of about 1.09 Å (1.09 × 10 m) and a bond energy of about 413 kJ/mol (see table below). Using Pauling's scale—C (2.55) and H (2.2)—the electronegativity difference between these two atoms is 0.35. Because of this small difference in electronegativities, the C−H bond is generally regarded as being non-polar. In structural formulas of molecules, the hydrogen atoms are often omitted. Compound classes consisting solely of C−H bonds and C−C bonds are alkanes, alkenes, alkynes, and aromatic hydrocarbons. Collectively they are known as hydrocarbons.

In chemistry, a lone pair refers to a pair of valence electrons that are not shared with another atom in a covalent bond and is sometimes called an unshared pair or non-bonding pair. Lone pairs are found in the outermost electron shell of atoms. They can be identified by using a Lewis structure. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding. Thus, the number of electrons in lone pairs plus the number of electrons in bonds equals the number of valence electrons around an atom.

Lone pair is a concept used in valence shell electron pair repulsion theory (VSEPR theory) which explains the shapes of molecules. They are also referred to in the chemistry of Lewis acids and bases. However, not all non-bonding pairs of electrons are considered by chemists to be lone pairs. Examples are the transition metals where the non-bonding pairs do not influence molecular geometry and are said to be stereochemically inactive. In molecular orbital theory (fully delocalized canonical orbitals or localized in some form), the concept of a lone pair is less distinct, as the correspondence between an orbital and components of a Lewis structure is often not straightforward. Nevertheless, occupied non-bonding orbitals (or orbitals of mostly nonbonding character) are frequently identified as lone pairs.

In quantum mechanics, the principal quantum number (n) of an electron in an atom indicates which electron shell or energy level it is in. Its values are natural numbers (1, 2, 3, ...).

Hydrogen and Helium, at their lowest energies, have just one electron shell. Lithium through Neon (see periodic table) have two shells: two electrons in the first shell, and up to 8 in the second shell. Larger atoms have more shells.

In atomic physics, a magnetic quantum number is a quantum number used to distinguish quantum states of an electron or other particle according to its angular momentum along a given axis in space. The orbital magnetic quantum number (m_l or m) distinguishes the orbitals available within a given subshell of an atom. It specifies the component of the orbital angular momentum that lies along a given axis, conventionally called the z-axis, so it describes the orientation of the orbital in space. The spin magnetic quantum number m_s specifies the z-axis component of the spin angular momentum for a particle having spin quantum number s. For an electron, s is 1⁄2, and m_s is either +1⁄2 or −1⁄2, often called "spin-up" and "spin-down", or α and β. The term magnetic in the name refers to the magnetic dipole moment associated with each type of angular momentum, so states having different magnetic quantum numbers shift in energy in a magnetic field according to the Zeeman effect.

The four quantum numbers conventionally used to describe the quantum state of an electron in an atom are the principal quantum number n, the azimuthal (orbital) quantum number ${\displaystyle \ell }$, and the magnetic quantum numbers m_l and m_s. Electrons in a given subshell of an atom (such as s, p, d, or f) are defined by values of ${\displaystyle \ell }$ (0, 1, 2, or 3). The orbital magnetic quantum number takes integer values in the range from ${\displaystyle -\ell }$ to ${\displaystyle +\ell }$, including zero. Thus the s, p, d, and f subshells contain 1, 3, 5, and 7 orbitals each. Each of these orbitals can accommodate up to two electrons (with opposite spins), forming the basis of the periodic table.

In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals. For example, the electron configuration of the neon atom is 1s 2s 2p, meaning that the 1s, 2s, and 2p subshells are occupied by two, two, and six electrons, respectively.

Electronic configurations describe each electron as moving independently in an orbital, in an average field created by the nuclei and all the other electrons. Mathematically, configurations are described by Slater determinants or configuration state functions.

A period on the periodic table is a row of chemical elements. All elements in a row have the same number of electron shells. Each next element in a period has one more proton and is less metallic than its predecessor. Arranged this way, elements in the same group (column) have similar chemical and physical properties, reflecting the periodic law. For example, the halogens lie in the second-to-last group (group 17) and share similar properties, such as high reactivity and the tendency to gain one electron to arrive at a noble-gas electronic configuration. As of 2022, a total of 118 elements have been discovered and confirmed.

Modern quantum mechanics explains these periodic trends in properties in terms of electron shells. As atomic number increases, shells fill with electrons in approximately the order shown in the ordering rule diagram. The filling of each shell corresponds to a row in the table.

Moscovium is a synthetic chemical element; it has symbol Mc and atomic number 115. It was first synthesized in 2003 by a joint team of Russian and American scientists at the Joint Institute for Nuclear Research (JINR) in Dubna, Russia. In December 2015, it was recognized as one of four new elements by the Joint Working Party of international scientific bodies IUPAC and IUPAP. On 28 November 2016, it was officially named after the Moscow Oblast, in which the JINR is situated.

Moscovium is an extremely radioactive element: its most stable known isotope, moscovium-290, has a half-life of only 0.65 seconds. In the periodic table, it is a p-block transactinide element. It is a member of the 7th period and is placed in group 15 as the heaviest pnictogen. Moscovium is calculated to have some properties similar to its lighter homologues, nitrogen, phosphorus, arsenic, antimony, and bismuth, and to be a post-transition metal, although it should also show several major differences from them. In particular, moscovium should also have significant similarities to thallium, as both have one rather loosely bound electron outside a quasi-closed shell. Chemical experimentation on single atoms has confirmed theoretical expectations that moscovium is less reactive than its lighter homologue bismuth. Over a hundred atoms of moscovium have been observed to date, all of which have been shown to have mass numbers from 286 to 290.

A lanthanide (/ˈlænθənaɪd/) is any of the 14 metallic chemical elements with atomic numbers 57–70, from lanthanum through ytterbium. The IUPAC recommended standard name of this series is lanthanoid (/ˈlænθənɔɪd/).In the periodic table, they fill the 4f orbitals. Lutetium (element 71) is also sometimes considered a lanthanide, despite being a d-block element and a transition metal.

The informal chemical symbol Ln is used in general discussions of lanthanide chemistry to refer to any lanthanide. All but one of the lanthanides are f-block elements, corresponding to the filling of the 4f electron shell. Lutetium is a d-block element (thus also a transition metal), and on this basis its inclusion has been questioned; however, like its congeners scandium and yttrium in group 3, it behaves similarly to the other 14. The term rare-earth element or rare-earth metal is often used to include the stable group 3 elements Sc, Y, and Lu in addition to the 4f elements. All lanthanide elements form trivalent cations, Ln, whose chemistry is largely determined by the ionic radius, which decreases steadily from lanthanum (La) to lutetium (Lu).

Energy-dispersive X-ray spectroscopy (EDS, EDX, EDXS or XEDS), sometimes called energy dispersive X-ray analysis (EDXA or EDAX) or energy dispersive X-ray microanalysis (EDXMA), is an analytical technique used for the elemental analysis or chemical characterization of a sample. It relies on an interaction of some source of X-ray excitation and a sample. Its characterization capabilities are due in large part to the fundamental principle that each element has a unique atomic structure allowing a unique set of peaks on its electromagnetic emission spectrum (which is the main principle of spectroscopy). The peak positions are predicted by the Moseley's law with accuracy much better than experimental resolution of a typical EDX instrument.

To stimulate the emission of characteristic X-rays from a specimen a beam of electrons or X-ray is focused into the sample being studied. At rest, an atom within the sample contains ground state (or unexcited) electrons in discrete energy levels or electron shells bound to the nucleus. The incident beam may excite an electron in an inner shell, ejecting it from the shell while creating an electron hole where the electron was. An electron from an outer, higher-energy shell then fills the hole, and the difference in energy between the higher-energy shell and the lower energy shell may be released in the form of an X-ray. The number and energy of the X-rays emitted from a specimen can be measured by an energy-dispersive spectrometer. As the energies of the X-rays are characteristic of the difference in energy between the two shells and of the atomic structure of the emitting element, EDS allows the elemental composition of the specimen to be measured.