The periodic table is a tabular arrangement of the chemical elements,
ordered by their atomic number, electron configuration, and recurring
chemical properties, whose adopted structure shows periodic trends.
Generally, within one row (period) the elements are metals on the
left, and non-metals on the right, with the elements having similar
chemical behaviours being placed in the same column. Table rows are
commonly called periods and columns are called groups. Six groups have
accepted names as well as assigned numbers: for example, group 17
elements are halogens; and group 18 are noble gases. Also displayed
are four simple rectangular areas or blocks associated with the
filling of different atomic orbitals.
Importantly, the organization of the periodic table can be utilized to
derive relationships between various element properties, but also
predicted chemical properties and behaviours of undiscovered or newly
synthesized elements. Russian chemist
Dmitri Mendeleev was first to
publish a recognizable periodic table in 1869, developed mainly to
illustrate periodic trends of the then-known elements. He also
predicted some properties of unidentified elements that were expected
to fill gaps within this table. Most of his forecasts proved to be
correct. Mendeleev's idea has been slowly expanded and refined with
the discovery or synthesis of further new elements and by developing
new theoretical models to explain chemical behaviour. The modern
periodic table now provides a useful framework for analyzing chemical
reactions, and continues to be widely adopted in chemistry, nuclear
physics and other sciences.
All elements ranging from atomic numbers 1 (hydrogen) to 118
(oganesson) have been either discovered or synthesized, completing the
first seven rows of the periodic table. The first 94 elements
exist naturally, although some are found only in trace amounts and
were synthesized in laboratories before being found in nature.[n 1]
Atomic numbers for elements 95 to 118 have only been synthesized in
laboratories or nuclear reactors. The synthesis of elements having
higher atomic numbers is currently being pursued: these elements would
begin an eighth row, and theoretical work has been done to suggest
possible appearances for this extension. Numerous synthetic
radionuclides of naturally occurring elements have also been produced
2 Grouping methods
2.4 Metals, metalloids and nonmetals
Periodic trends and patterns
3.2 Atomic radii
3.3 Ionization energy
3.6 Metallic character
3.7 Linking or bridging groups
4.1 First systemization attempts
4.2 Mendeleev's table
4.3 Second version and further development
5 Different periodic tables
5.1 The long- or 32-column table
5.2 Tables with different structures
6 Open questions and controversies
6.1 Placement of hydrogen and helium
6.2 Group 3 and its elements in periods 6 and 7
Lanthanum and actinium
Lutetium and lawrencium
Lanthanides and actinides
6.3 Groups included in the transition metals
6.4 Elements with unknown chemical properties
6.5 Further periodic table extensions
6.6 Element with the highest possible atomic number
6.7 Optimal form
7 See also
11 External links
For a more detailed periodic table, see
Periodic table (large cells).
Alkaline earth metals
1 (red)=Gas 3 (black)=Solid 80 (green)=Liquid 109 (gray)=Unknown Color
of the atomic number shows state of matter (at 0 °C and
Primordial From decay Synthetic Border shows natural
occurrence of the element
Standard atomic weight (Ar)
Ca: 7001400780000000000♠40.078 — Formal short value, rounded
Po:  — mass number of the most stable isotope
Background color shows subcategory in the metal–metalloid–nonmetal
Alkaline earth metal
Each chemical element has a unique atomic number (Z) representing the
number of protons in its nucleus.[n 2] Most elements have differing
numbers of neutrons among different atoms, with these variants being
referred to as isotopes. For example, carbon has three naturally
occurring isotopes: all of its atoms have six protons and most have
six neutrons as well, but about one per cent have seven neutrons, and
a very small fraction have eight neutrons. Isotopes are never
separated in the periodic table; they are always grouped together
under a single element. Elements with no stable isotopes have the
atomic masses of their most stable isotopes, where such masses are
shown, listed in parentheses.
In the standard periodic table, the elements are listed in order of
increasing atomic number Z (the number of protons in the nucleus of an
atom). A new row (period) is started when a new electron shell has its
first electron. Columns (groups) are determined by the electron
configuration of the atom; elements with the same number of electrons
in a particular subshell fall into the same columns (e.g. oxygen and
selenium are in the same column because they both have four electrons
in the outermost p-subshell). Elements with similar chemical
properties generally fall into the same group in the periodic table,
although in the f-block, and to some respect in the d-block, the
elements in the same period tend to have similar properties, as well.
Thus, it is relatively easy to predict the chemical properties of an
element if one knows the properties of the elements around it.
As of 2016[update], the periodic table has 118 confirmed elements,
from element 1 (hydrogen) to 118 (oganesson). Elements 113, 115, 117
and 118, the most recent discoveries, were officially confirmed by the
International Union of Pure and Applied
Chemistry (IUPAC) in December
2015. Their proposed names, nihonium (Nh), moscovium (Mc), tennessine
(Ts) and oganesson (Og) respectively, were announced by the IUPAC in
June 2016 and made official in November 2016.
The first 94 elements occur naturally; the remaining 24, americium to
oganesson (95–118), occur only when synthesized in laboratories. Of
the 94 naturally occurring elements, 83 are primordial and 11 occur
only in decay chains of primordial elements. No element heavier
than einsteinium (element 99) has ever been observed in macroscopic
quantities in its pure form, nor has astatine (element 85); francium
(element 87) has been only photographed in the form of light emitted
from microscopic quantities (300,000 atoms).
Main article: Group (periodic table)
A group or family is a vertical column in the periodic table. Groups
usually have more significant periodic trends than periods and blocks,
explained below. Modern quantum mechanical theories of atomic
structure explain group trends by proposing that elements within the
same group generally have the same electron configurations in their
valence shell. Consequently, elements in the same group tend to
have a shared chemistry and exhibit a clear trend in properties with
increasing atomic number. In some parts of the periodic table,
such as the d-block and the f-block, horizontal similarities can be as
important as, or more pronounced than, vertical
Under an international naming convention, the groups are numbered
numerically from 1 to 18 from the leftmost column (the alkali metals)
to the rightmost column (the noble gases). Previously, they were
known by roman numerals. In America, the roman numerals were followed
by either an "A" if the group was in the s- or p-block, or a "B" if
the group was in the d-block. The roman numerals used correspond to
the last digit of today's naming convention (e.g. the group 4 elements
were group IVB, and the group 14 elements were group IVA). In Europe,
the lettering was similar, except that "A" was used if the group was
before group 10, and "B" was used for groups including and after group
10. In addition, groups 8, 9 and 10 used to be treated as one
triple-sized group, known collectively in both notations as group
VIII. In 1988, the new IUPAC naming system was put into use, and the
old group names were deprecated.
Some of these groups have been given trivial (unsystematic) names, as
seen in the table below, although some are rarely used. Groups 3–10
have no trivial names and are referred to simply by their group
numbers or by the name of the first member of their group (such as
"the scandium group" for group 3), since they display fewer
similarities and/or vertical trends.
Elements in the same group tend to show patterns in atomic radius,
ionization energy, and electronegativity. From top to bottom in a
group, the atomic radii of the elements increase. Since there are more
filled energy levels, valence electrons are found farther from the
nucleus. From the top, each successive element has a lower ionization
energy because it is easier to remove an electron since the atoms are
less tightly bound. Similarly, a group has a top-to-bottom decrease in
electronegativity due to an increasing distance between valence
electrons and the nucleus. There are exceptions to these trends:
for example, in group 11, electronegativity increases farther down the
Groups in the Periodic table
CAS (US, A-B-A)
old IUPAC (Europe, A-B)
Alkaline earth metalsr
Name by elementr
h H h
a Group 3 has scandium (Sc) and yttrium (Y). For the rest of the
group, sources differ as either being (1) lutetium (Lu) and lawrencium
(Lr), or (2) lanthanum (La) and actinium (Ac), or (3) the whole set of
15+15 lanthanides and actinides. IUPAC has initiated a project to
standardize the definition as either (1) Sc, Y, Lu and Lr, or (2) Sc,
Y, La and Ac.
b Group 18, the noble gases, were not discovered at the time of
Mendeleev's original table. Later (1902), Mendeleev accepted the
evidence for their existence, and they could be placed in a new "group
0", consistently and without breaking the periodic table principle.
r Group name as recommended by IUPAC.
Hydrogen (H), while placed in group 1, is not considered to be part
of the alkali metals.
Main article: Period (periodic table)
A period is a horizontal row in the periodic table. Although groups
generally have more significant periodic trends, there are regions
where horizontal trends are more significant than vertical group
trends, such as the f-block, where the lanthanides and actinides form
two substantial horizontal series of elements.
Elements in the same period show trends in atomic radius, ionization
energy, electron affinity, and electronegativity. Moving left to right
across a period, atomic radius usually decreases. This occurs because
each successive element has an added proton and electron, which causes
the electron to be drawn closer to the nucleus. This decrease in
atomic radius also causes the ionization energy to increase when
moving from left to right across a period. The more tightly bound an
element is, the more energy is required to remove an electron.
Electronegativity increases in the same manner as ionization energy
because of the pull exerted on the electrons by the nucleus.
Electron affinity also shows a slight trend across a period. Metals
(left side of a period) generally have a lower electron affinity than
nonmetals (right side of a period), with the exception of the noble
Main article: Block (periodic table)
Left to right: s-, f-, d-, p-block in the periodic table
Specific regions of the periodic table can be referred to as blocks in
recognition of the sequence in which the electron shells of the
elements are filled. Each block is named according to the subshell in
which the "last" electron notionally resides.[n 3] The s-block
comprises the first two groups (alkali metals and alkaline earth
metals) as well as hydrogen and helium. The p-block comprises the last
six groups, which are groups 13 to 18 in IUPAC group numbering
(3A to 8A in American group numbering) and contains, among other
elements, all of the metalloids. The d-block comprises groups 3 to 12
(or 3B to 2B in American group numbering) and contains all of the
transition metals. The f-block, often offset below the rest of the
periodic table, has no group numbers and comprises lanthanides and
Metals, metalloids and nonmetals
Metals, metalloids, nonmetals, and
elements with unknown chemical properties in the periodic
table. Sources disagree on the classification of some of these
According to their shared physical and chemical properties, the
elements can be classified into the major categories of metals,
metalloids and nonmetals. Metals are generally shiny, highly
conducting solids that form alloys with one another and salt-like
ionic compounds with nonmetals (other than noble gases). A majority of
nonmetals are coloured or colourless insulating gases; nonmetals that
form compounds with other nonmetals, feature covalent bonding. In
between metals and nonmetals are metalloids, which have intermediate
or mixed properties.
Metal and nonmetals can be further classified into subcategories that
show a gradation from metallic to non-metallic properties, when going
left to right in the rows. The metals may be subdivided into the
highly reactive alkali metals, through the less reactive alkaline
earth metals, lanthanides and actinides, via the archetypal transition
metals, and ending in the physically and chemically weak
Nonmetals may be simply subdivided into the
polyatomic nonmetals, being nearer to the metalloids and show some
incipient metallic character; the essentially nonmetallic diatomic
nonmetals, nonmetallic and the almost completely inert, nonatomic
noble gases. Specialized groupings such as refractory metals and noble
metals, are examples of subsets of transition metals, also known
and occasionally denoted.
Placing elements into categories and subcategories based just on
shared properties is imperfect. There is a large disparity of
properties within each category with notable overlaps at the
boundaries, as is the case with most classification schemes.
Beryllium, for example, is classified as an alkaline earth metal
although its amphoteric chemistry and tendency to mostly form covalent
compounds are both attributes of a chemically weak or post-transition
Radon is classified as a nonmetallic noble gas yet has some
cationic chemistry that is characteristic of metals. Other
classification schemes are possible such as the division of the
elements into mineralogical occurrence categories, or crystalline
structures. Categorizing the elements in this fashion dates back to at
least 1869 when Hinrichs wrote that simple boundary lines could be
placed on the periodic table to show elements having shared
properties, such as metals, nonmetals, or gaseous elements.
Periodic trends and patterns
Main article: Periodic trends
Main article: Electronic configuration
Approximate order in which shells and subshells are arranged by
increasing energy according to the Madelung rule
The electron configuration or organisation of electrons orbiting
neutral atoms shows a recurring pattern or periodicity. The electrons
occupy a series of electron shells (numbered 1, 2, and so on). Each
shell consists of one or more subshells (named s, p, d, f and g). As
atomic number increases, electrons progressively fill these shells and
subshells more or less according to the
Madelung rule or energy
ordering rule, as shown in the diagram. The electron configuration for
neon, for example, is 1s2 2s2 2p6. With an atomic number of ten, neon
has two electrons in the first shell, and eight electrons in the
second shell; there are two electrons in the s subshell and six in the
p subshell. In periodic table terms, the first time an electron
occupies a new shell corresponds to the start of each new period,
these positions being occupied by hydrogen and the alkali
Periodic table trends (arrows show an increase)
Since the properties of an element are mostly determined by its
electron configuration, the properties of the elements likewise show
recurring patterns or periodic behaviour, some examples of which are
shown in the diagrams below for atomic radii, ionization energy and
electron affinity. It is this periodicity of properties,
manifestations of which were noticed well before the underlying theory
was developed, that led to the establishment of the periodic law (the
properties of the elements recur at varying intervals) and the
formulation of the first periodic tables.
Main article: Atomic radius
Atomic number plotted against atomic radius[n 4]
Atomic radii vary in a predictable and explainable manner across the
periodic table. For instance, the radii generally decrease along each
period of the table, from the alkali metals to the noble gases; and
increase down each group. The radius increases sharply between the
noble gas at the end of each period and the alkali metal at the
beginning of the next period. These trends of the atomic radii (and of
various other chemical and physical properties of the elements) can be
explained by the electron shell theory of the atom; they provided
important evidence for the development and confirmation of quantum
The electrons in the 4f-subshell, which is progressively filled across
the lanthanide series, are not particularly effective at shielding the
increasing nuclear charge from the sub-shells further out. The
elements immediately following the lanthanides have atomic radii that
are smaller than would be expected and that are almost identical to
the atomic radii of the elements immediately above them. Hence
hafnium has virtually the same atomic radius (and chemistry) as
zirconium, and tantalum has an atomic radius similar to niobium, and
so forth. This is known as the lanthanide contraction. The effect of
the lanthanide contraction is noticeable up to platinum (element 78),
after which it is masked by a relativistic effect known as the inert
pair effect. The d-block contraction, which is a similar effect
between the d-block and p-block, is less pronounced than the
lanthanide contraction but arises from a similar cause.
Ionization energy: each period begins at a minimum for the alkali
metals, and ends at a maximum for the noble gases
Main article: Ionization energy
The first ionization energy is the energy it takes to remove one
electron from an atom, the second ionization energy is the energy it
takes to remove a second electron from the atom, and so on. For a
given atom, successive ionization energies increase with the degree of
ionization. For magnesium as an example, the first ionization energy
is 738 kJ/mol and the second is 1450 kJ/mol. Electrons in
the closer orbitals experience greater forces of electrostatic
attraction; thus, their removal requires increasingly more energy.
Ionization energy becomes greater up and to the right of the periodic
Large jumps in the successive molar ionization energies occur when
removing an electron from a noble gas (complete electron shell)
configuration. For magnesium again, the first two molar ionization
energies of magnesium given above correspond to removing the two 3s
electrons, and the third ionization energy is a much larger
7730 kJ/mol, for the removal of a 2p electron from the very
stable neon-like configuration of Mg2+. Similar jumps occur in the
ionization energies of other third-row atoms.
Main article: Electronegativity
Graph showing increasing electronegativity with growing number of
Electronegativity is the tendency of an atom to attract a shared pair
of electrons. An atom's electronegativity is affected by both its
atomic number and the distance between the valence electrons and the
nucleus. The higher its electronegativity, the more an element
attracts electrons. It was first proposed by
Linus Pauling in
1932. In general, electronegativity increases on passing from left
to right along a period, and decreases on descending a group. Hence,
fluorine is the most electronegative of the elements,[n 5] while
caesium is the least, at least of those elements for which substantial
data is available.
There are some exceptions to this general rule.
Gallium and germanium
have higher electronegativities than aluminium and silicon
respectively because of the d-block contraction. Elements of the
fourth period immediately after the first row of the transition metals
have unusually small atomic radii because the 3d-electrons are not
effective at shielding the increased nuclear charge, and smaller
atomic size correlates with higher electronegativity. The
anomalously high electronegativity of lead, particularly when compared
to thallium and bismuth, appears to be an artifact of data selection
and data availability. Methods of calculation other than the Pauling
method show the normal periodic trends for these elements.
Dependence of electron affinity on atomic number. Values generally
increase across each period, culminating with the halogens before
decreasing precipitously with the noble gases. Examples of localized
peaks seen in hydrogen, the alkali metals and the group 11 elements
are caused by a tendency to complete the s-shell (with the 6s shell of
gold being further stabilized by relativistic effects and the presence
of a filled 4f sub shell). Examples of localized troughs seen in the
alkaline earth metals, and nitrogen, phosphorus, manganese and rhenium
are caused by filled s-shells, or half-filled p- or d-shells.
The electron affinity of an atom is the amount of energy released when
an electron is added to a neutral atom to form a negative ion.
Although electron affinity varies greatly, some patterns emerge.
Generally, nonmetals have more positive electron affinity values than
Chlorine most strongly attracts an extra electron. The
electron affinities of the noble gases have not been measured
conclusively, so they may or may not have slightly negative
Electron affinity generally increases across a period. This is caused
by the filling of the valence shell of the atom; a group 17 atom
releases more energy than a group 1 atom on gaining an electron
because it obtains a filled valence shell and is therefore more
A trend of decreasing electron affinity going down groups would be
expected. The additional electron will be entering an orbital farther
away from the nucleus. As such this electron would be less attracted
to the nucleus and would release less energy when added. In going down
a group, around one-third of elements are anomalous, with heavier
elements having higher electron affinities than their next lighter
congenors. Largely, this is due to the poor shielding by d and f
electrons. A uniform decrease in electron affinity only applies to
group 1 atoms.
The lower the values of ionization energy, electronegativity and
electron affinity, the more metallic character the element has.
Conversely, nonmetallic character increases with higher values of
these properties. Given the periodic trends of these three
properties, metallic character tends to decrease going across a period
(or row) and, with some irregularities (mostly) due to poor screening
of the nucleus by d and f electrons, and relativistic effects,
tends to increase going down a group (or column or family). Thus, the
most metallic elements (such as caesium and francium) are found at the
bottom left of traditional periodic tables and the most nonmetallic
elements (oxygen, fluorine, chlorine) at the top right. The
combination of horizontal and vertical trends in metallic character
explains the stair-shaped dividing line between metals and nonmetals
found on some periodic tables, and the practice of sometimes
categorizing several elements adjacent to that line, or elements
adjacent to those elements, as metalloids.
Linking or bridging groups
32-column periodic table showing, from left to right, the location of
group 3; lutetium and lawrencium; groups 11–12; and the noble gases
From left to right across the four blocks of the long- or 32-column
form of the periodic table are a series of linking or bridging groups
of elements, located approximately between each block. These
groups, like the metalloids, show properties in between, or that are a
mixture of, groups to either side. Chemically, the group 3 elements,
scandium, yttrium, lanthanum and actinium behave largely like the
alkaline earth metals or, more generally, s block metals
but have some of the physical properties of d block transition
Lutetium and lawrencium, at the end of the end of the f
block, may constitute another linking or bridging group. Lutetium
behaves chemically as a lanthanide but shows a mix of lanthanide and
transition metal physical properties. Lawrencium, as an
analogue of lutetium, would presumable display like characteristics.[n
6] The coinage metals in group 11 (copper, silver, and gold) are
chemically capable of acting as either transition metals or main group
metals. The volatile group 12 metals, zinc, cadmium and mercury
are sometimes regarded as linking the d block to the p block.
Notionally they are d block elements but they have few transition
metal properties and are more like their p block neighbors in group
13. The relatively inert noble gases, in group 18, bridge the
most reactive groups of elements in the periodic table—the halogens
in group 17 and the alkali metals in group 1.
Main article: History of the periodic table
First systemization attempts
The discovery of the elements mapped to significant periodic table
development dates (pre-, per- and post-)
Antoine Lavoisier published a list of 33 chemical elements,
grouping them into gases, metals, nonmetals, and earths. Chemists
spent the following century searching for a more precise
classification scheme. In 1829,
Johann Wolfgang Döbereiner
Johann Wolfgang Döbereiner observed
that many of the elements could be grouped into triads based on their
chemical properties. Lithium, sodium, and potassium, for example, were
grouped together in a triad as soft, reactive metals. Döbereiner also
observed that, when arranged by atomic weight, the second member of
each triad was roughly the average of the first and the third;
this became known as the Law of Triads. German chemist Leopold
Gmelin worked with this system, and by 1843 he had identified ten
triads, three groups of four, and one group of five. Jean-Baptiste
Dumas published work in 1857 describing relationships between various
groups of metals. Although various chemists were able to identify
relationships between small groups of elements, they had yet to build
one scheme that encompassed them all.
In 1857, German chemist August Kekulé observed that carbon often has
four other atoms bonded to it. Methane, for example, has one carbon
atom and four hydrogen atoms. This concept eventually became known
as valency; different elements bond with different numbers of
In 1862, Alexandre-Emile Béguyer de Chancourtois, a French geologist,
published an early form of periodic table, which he called the
telluric helix or screw. He was the first person to notice the
periodicity of the elements. With the elements arranged in a spiral on
a cylinder by order of increasing atomic weight, de Chancourtois
showed that elements with similar properties seemed to occur at
regular intervals. His chart included some ions and compounds in
addition to elements. His paper also used geological rather than
chemical terms and did not include a diagram; as a result, it received
little attention until the work of Dmitri Mendeleev.
In 1864, Julius Lothar Meyer, a German chemist, published a table with
44 elements arranged by valency. The table showed that elements with
similar properties often shared the same valency. Concurrently,
William Odling (an English chemist) published an arrangement of 57
elements, ordered on the basis of their atomic weights. With some
irregularities and gaps, he noticed what appeared to be a periodicity
of atomic weights among the elements and that this accorded with
"their usually received groupings". Odling alluded to the idea of
a periodic law but did not pursue it. He subsequently proposed (in
1870) a valence-based classification of the elements.
Newlands' periodic table, as presented to the
Chemical Society in
1866, and based on the law of octaves
English chemist John Newlands produced a series of papers from 1863 to
1866 noting that when the elements were listed in order of increasing
atomic weight, similar physical and chemical properties recurred at
intervals of eight; he likened such periodicity to the octaves of
music. This so termed
Law of Octaves
Law of Octaves was ridiculed by
Newlands' contemporaries, and the
Chemical Society refused to publish
his work. Newlands was nonetheless able to draft a table of the
elements and used it to predict the existence of missing elements,
such as germanium. The
Chemical Society only acknowledged the
significance of his discoveries five years after they credited
In 1867, Gustavus Hinrichs, a Danish born academic chemist based in
America, published a spiral periodic system based on atomic spectra
and weights, and chemical similarities. His work was regarded as
idiosyncratic, ostentatious and labyrinthine and this may have
militated against its recognition and acceptance.
Dmitri Mendeleev, watercolour by Ilya Repin
Mendeleev's periodic table from his book An Attempt Towards a Chemical
Conception of the Ether
A version of Mendeleev's 1869 periodic table: An experiment on a
system of elements based on their atomic weights and chemical
similarities. This early arrangement presents the periods vertically,
and the groups horizontally.
Russian chemistry professor
Dmitri Mendeleev and German chemist Julius
Lothar Meyer independently published their periodic tables in 1869 and
1870, respectively. Mendeleev's table was his first published
version; that of Meyer was an expanded version of his (Meyer's) table
of 1864. They both constructed their tables by listing the
elements in rows or columns in order of atomic weight and starting a
new row or column when the characteristics of the elements began to
The recognition and acceptance afforded to Mendeleev's table came from
two decisions he made. The first was to leave gaps in the table when
it seemed that the corresponding element had not yet been
discovered. Mendeleev was not the first chemist to do so, but he
was the first to be recognized as using the trends in his periodic
table to predict the properties of those missing elements, such as
gallium and germanium. The second decision was to occasionally
ignore the order suggested by the atomic weights and switch adjacent
elements, such as tellurium and iodine, to better classify them into
chemical families. Later, in 1913,
Henry Moseley determined
experimental values of the nuclear charge or atomic number of each
element, and showed that Mendeleev's ordering actually corresponds to
the order of increasing atomic number.
The significance of atomic numbers to the organization of the periodic
table was not appreciated until the existence and properties of
protons and neutrons became understood. Mendeleev's periodic tables
used atomic weight instead of atomic number to organize the elements,
information determinable to fair precision in his time. Atomic weight
worked well enough in most cases to (as noted) give a presentation
that was able to predict the properties of missing elements more
accurately than any other method then known. Substitution of atomic
numbers, once understood, gave a definitive, integer-based sequence
for the elements, and Moseley predicted (in 1913) that the only
elements still missing between aluminium (Z=13) and gold (Z=79) were Z
= 43, 61, 72 and 75, all of which were later discovered. The sequence
of atomic numbers is still used today even as new synthetic elements
are being produced and studied.
Second version and further development
Mendeleev's 1871 periodic table with eight groups of elements. Dashes
represented elements unknown in 1871.
Eight-column form of periodic table, updated with all elements
discovered to 2016
In 1871, Mendeleev published his periodic table in a new form, with
groups of similar elements arranged in columns rather than in rows,
and those columns numbered I to VIII corresponding with the element's
oxidation state. He also gave detailed predictions for the properties
of elements he had earlier noted were missing, but should exist.
These gaps were subsequently filled as chemists discovered additional
naturally occurring elements. It is often stated that the last
naturally occurring element to be discovered was francium (referred to
by Mendeleev as eka-caesium) in 1939. Plutonium, produced
synthetically in 1940, was identified in trace quantities as a
naturally occurring element in 1971.
The popular periodic table layout, also known as the common or
standard form (as shown at various other points in this article), is
attributable to Horace Groves Deming. In 1923, Deming, an American
chemist, published short (Mendeleev style) and medium (18-column) form
periodic tables.[n 7] Merck and Company prepared a handout form of
Deming's 18-column medium table, in 1928, which was widely circulated
in American schools. By the 1930s Deming's table was appearing in
handbooks and encyclopaedias of chemistry. It was also distributed for
many years by the Sargent-Welch Scientific Company.
With the development of modern quantum mechanical theories of electron
configurations within atoms, it became apparent that each period (row)
in the table corresponded to the filling of a quantum shell of
electrons. Larger atoms have more electron sub-shells, so later tables
have required progressively longer periods.
Glenn T. Seaborg, in 1945, suggested a new periodic table showing the
actinides as belonging to a second f-block series
In 1945, Glenn Seaborg, an American scientist, made the suggestion
that the actinide elements, like the lanthanides, were filling an f
sub-level. Before this time the actinides were thought to be forming a
fourth d-block row. Seaborg's colleagues advised him not to publish
such a radical suggestion as it would most likely ruin his career. As
Seaborg considered he did not then have a career to bring into
disrepute, he published anyway. Seaborg's suggestion was found to be
correct and he subsequently went on to win the 1951
Nobel Prize in
chemistry for his work in synthesizing actinide elements.[n 8]
Although minute quantities of some transuranic elements occur
naturally, they were all first discovered in laboratories. Their
production has expanded the periodic table significantly, the first of
these being neptunium, synthesized in 1939. Because many of the
transuranic elements are highly unstable and decay quickly, they are
challenging to detect and characterize when produced. There have been
controversies concerning the acceptance of competing discovery claims
for some elements, requiring independent review to determine which
party has priority, and hence naming rights. In 2010, a joint
Russia–US collaboration at Dubna, Moscow Oblast, Russia, claimed to
have synthesized six atoms of tennessine (element 117), making it the
most recently claimed discovery. It, along with nihonium (element
113), moscovium (element 115), and oganesson (element 118), are the
four most recently named elements, whose names all became official on
28 November 2016.
Different periodic tables
The long- or 32-column table
The periodic table in 32-column format
The modern periodic table is sometimes expanded into its long or
32-column form by reinstating the footnoted f-block elements into
their natural position between the s- and d-blocks. Unlike the
18-column form this arrangement results in "no interruptions in the
sequence of increasing atomic numbers". The relationship of the
f-block to the other blocks of the periodic table also becomes easier
to see. Jensen advocates a form of table with 32 columns on the
grounds that the lanthanides and actinides are otherwise relegated in
the minds of students as dull, unimportant elements that can be
quarantined and ignored. Despite these advantages the 32-column
form is generally avoided by editors on account of its undue
rectangular ratio (compared to a book page ratio), and the
familiarity of chemists with the modern form (as introduced by
Tables with different structures
Main article: Alternative periodic tables
Within 100 years of the appearance of Mendeleev's table in 1869 it has
been estimated that around 700 different periodic table versions were
published. As well as numerous rectangular variations, other
periodic table formats have been shaped, for example,[n 9] like a
circle, cube, cylinder, building, spiral, lemniscate, octagonal
prism, pyramid, sphere, or triangle. Such alternatives are often
developed to highlight or emphasize chemical or physical properties of
the elements that are not as apparent in traditional periodic
Theodor Benfey's spiral periodic table
A popular alternative structure is that of
Theodor Benfey (1960).
The elements are arranged in a continuous spiral, with hydrogen at the
centre and the transition metals, lanthanides, and actinides occupying
Most periodic tables are two-dimensional; three-dimensional tables
are known to as far back as at least 1862 (pre-dating Mendeleev's
two-dimensional table of 1869). More recent examples include
Courtines' Periodic Classification (1925), Wringley's Lamina
System (1949), Giguère's Periodic helix (1965) and Dufour's
Periodic Tree (1996). Going one further, Stowe's Physicist's
Periodic Table (1989) has been described as being
four-dimensional (having three spatial dimensions and one colour
The various forms of periodic tables can be thought of as lying on a
chemistry–physics continuum. Towards the chemistry end of the
continuum can be found, as an example, Rayner-Canham's "unruly"
Inorganic Chemist's Periodic Table (2002), which emphasizes
trends and patterns, and unusual chemical relationships and
properties. Near the physics end of the continuum is Janet's Left-Step
Periodic Table (1928). This has a structure that shows a closer
connection to the order of electron-shell filling and, by association,
quantum mechanics. A somewhat similar approach has been taken by
Alper, albeit criticized by
Eric Scerri as disregarding the need
to display chemical and physical periodicity. Somewhere in the
middle of the continuum is the ubiquitous common or standard form of
periodic table. This is regarded as better expressing empirical trends
in physical state, electrical and thermal conductivity, and oxidation
numbers, and other properties easily inferred from traditional
techniques of the chemical laboratory. Its popularity is thought
to be a result of this layout having a good balance of features in
terms of ease of construction and size, and its depiction of atomic
order and periodic trends.
Left-step periodic table (by Charles Janet)
3d 4p 5s
4d 5p 6s
4f 5d 6p 7s
5f 6d 7p 8s
This form of periodic table is more congruent with the order in which
electron shells are ideally filled according to the Madelung rule, as
shown in the accompanying sequence in the left margin (read from top
to bottom, left to right). In reality, the filling of electron shells
is characterized by a number of irregularities.
Open questions and controversies
Placement of hydrogen and helium
Simply following electron configurations, hydrogen (electronic
configuration 1s1) and helium (1s2) should be placed in groups 1 and
2, above lithium (1s22s1) and beryllium (1s22s2). While such a
placement is common for hydrogen, it is rarely used for helium outside
of the context of electron configurations: When the noble gases (then
called "inert gases") were first discovered around 1900, they were
known as "group 0", reflecting no chemical reactivity of these
elements known at that point, and helium was placed on the top of that
group, as it did share the extreme chemical inertness seen throughout
the group. As the group changed its formal number, many authors
continued to assign helium directly above neon, in group 18; one of
the examples of such placing is the current IUPAC table.
Hydrogen's chemical properties are not very close to those of the
alkali metals, which occupy group 1. On this basis it is sometimes
placed elsewhere. A common alternative is at the top of group 17
given hydrogen's strictly univalent and largely non-metallic
chemistry, and the strictly univalent and non-metallic chemistry of
fluorine (the element otherwise at the top of group 17). Sometimes, to
show hydrogen has properties corresponding to both those of the alkali
metals and the halogens, it is shown at the top of the two columns
simultaneously. Another suggestion is above carbon in group 14:
placed that way, it fits well into the trends of increasing ionization
potential values and electron affinity values, and is not too far from
the electronegativity trend, even though hydrogen cannot show the
tetravalence characteristic of the heavier group 14 elements.
Finally, hydrogen is sometimes placed separately from any group; this
is based on how general properties of hydrogen differ from that of any
group. The other period 1 element, helium, is sometimes placed
separately from any group as well. The property that
distinguishes helium from the rest of the noble gases (even though the
extraordinary inertness of helium is extremely close to that of neon
and argon) is that in its closed electron shell, helium has only
two electrons in the outermost electron orbital, while the rest of the
noble gases have eight.
Group 3 and its elements in periods 6 and 7
Although scandium and yttrium are always the first two elements in
group 3, the identity of the next two elements is not completely
settled. They are commonly lanthanum and actinium, and less often
lutetium and lawrencium. The two variants originate from historical
difficulties in placing the lanthanides in the periodic table, and
arguments as to where the f block elements start and end.[n 10][n
11] It has been claimed that such arguments are proof that, "it is a
mistake to break the [periodic] system into sharply delimited
blocks". A third variant shows the two positions below yttrium as
being occupied by the lanthanides and the actinides.
Chemical and physical arguments have been made in support of lutetium
and lawrencium but the majority of authors seem
unconvinced. Most working chemists are not aware there is any
controversy. In December 2015 an IUPAC project was established to
make a recommendation on the matter.
Lanthanum and actinium
La and Ac below Y
Lanthanum and actinium are commonly depicted as the remaining group 3
members.[n 12] It has been suggested that this layout originated
in the 1940s, with the appearance of periodic tables relying on the
electron configurations of the elements and the notion of the
differentiating electron. The configurations of caesium, barium and
lanthanum are [Xe]6s1, [Xe]6s2 and [Xe]5d16s2.
Lanthanum thus has a 5d
differentiating electron and this establishes it "in group 3 as the
first member of the d-block for period 6". A consistent set of
electron configurations is then seen in group 3: scandium [Ar]3d14s2,
yttrium [Kr]4d15s2 and lanthanum [Xe]5d16s2. Still in period 6,
ytterbium was assigned an electron configuration of [Xe]4f135d16s2 and
lutetium [Xe]4f145d16s2, "resulting in a 4f differentiating electron
for lutetium and firmly establishing it as the last member of the
f-block for period 6". Later spectroscopic work found that the
electron configuration of ytterbium was in fact [Xe]4f146s2. This
meant that ytterbium and lutetium—the latter with
[Xe]4f145d16s2—both had 14 f-electrons, "resulting in a d- rather
than an f- differentiating electron" for lutetium and making it an
"equally valid candidate" with [Xe]5d16s2 lanthanum, for the group 3
periodic table position below yttrium.
Lanthanum has the
advantage of incumbency since the 5d1 electron appears for the first
time in its structure whereas it appears for the third time in
lutetium, having also made a brief second appearance in
In terms of chemical behaviour, and trends going down group 3 for
properties such as melting point, electronegativity and ionic
radius, scandium, yttrium, lanthanum and actinium are
similar to their group 1–2 counterparts. In this variant, the number
of f electrons in the most common (trivalent) ions of the f-block
elements consistently matches their position in the f-block. For
example, the f-electron counts for the trivalent ions of the first
three f-block elements are Ce 1, Pr 2 and Nd 3.
Lutetium and lawrencium
Lu and Lr below Y
In other tables, lutetium and lawrencium are the remaining group 3
members.[n 13] Early techniques for chemically separating scandium,
yttrium and lutetium relied on the fact that these elements occurred
together in the so-called "yttrium group" whereas La and Ac occurred
together in the "cerium group". Accordingly, lutetium rather than
lanthanum was assigned to group 3 by some chemists in the 1920s and
30s.[n 14] Several physicists in the 1950s and '60s favoured lutetium,
in light of a comparison of several of its physical properties with
those of lanthanum. This arrangement, in which lanthanum is the
first member of the f-block, is disputed by some authors since
lanthanum lacks any f-electrons. It has been argued that this is not
valid concern given other periodic table anomalies—thorium, for
example, has no f-electrons yet is part of the f-block. As for
lawrencium, its gas phase atomic electron configuration was confirmed
in 2015 as [Rn]5f147s27p1. Such a configuration represents another
periodic table anomaly, regardless of whether lawrencium is located in
the f-block or the d-block, as the only potentially applicable p-block
position has been reserved for nihonium with its predicted
configuration of [Rn]5f146d107s27p1.[n 15]
Chemically, scandium, yttrium and lutetium (and presumably lawrencium)
behave like trivalent versions of the group 1–2 metals. On the
other hand, trends going down the group for properties such as melting
point, electronegativity and ionic radius, are similar to those found
among their group 4–8 counterparts. In this variant, the number
of f electrons in the gaseous forms of the f-block atoms usually
matches their position in the f-block. For example, the f-electron
counts for the first five f-block elements are La 0, Ce 1, Pr 3, Nd 4
and Pm 5.
Lanthanides and actinides
Markers below Y
A few authors position all thirty lanthanides and actinides in the two
positions below yttrium (usually via footnote markers). This variant
emphasizes similarities in the chemistry of the 15 lanthanide elements
(La–Lu), possibly at the expense of ambiguity as to which elements
occupy the two group 3 positions below yttrium, and a 15-column wide f
block (there can only be 14 elements in any row of the f block).[n 16]
Groups included in the transition metals
The definition of a transition metal, as given by IUPAC, is an element
whose atom has an incomplete d sub-shell, or which can give rise to
cations with an incomplete d sub-shell. By this definition all of
the elements in groups 3–11 are transition metals. The IUPAC
definition therefore excludes group 12, comprising zinc, cadmium and
mercury, from the transition metals category.
Some chemists treat the categories "d-block elements" and "transition
metals" interchangeably, thereby including groups 3–12 among the
transition metals. In this instance the group 12 elements are treated
as a special case of transition metal in which the d electrons are not
ordinarily involved in chemical bonding. The 2007 report of
mercury(IV) fluoride (HgF4), a compound in which mercury would use its
d electrons for bonding, has prompted some commentators to suggest
that mercury can be regarded as a transition metal. Other
commentators, such as Jensen, have argued that the formation of a
compound like HgF4 can occur only under highly abnormal conditions;
indeed, its existence is currently disputed. As such, mercury could
not be regarded as a transition metal by any reasonable interpretation
of the ordinary meaning of the term.
Still other chemists further exclude the group 3 elements from the
definition of a transition metal. They do so on the basis that the
group 3 elements do not form any ions having a partially occupied d
shell and do not therefore exhibit any properties characteristic of
transition metal chemistry. In this case, only groups 4–11 are
regarded as transition metals. Though the group 3 elements show few of
the characteristic chemical properties of the transition metals, they
do show some of their characteristic physical properties (on account
of the presence in each atom of a single d electron).
Elements with unknown chemical properties
Although all elements up to oganesson have been discovered, of the
elements above hassium (element 108), only copernicium (element 112),
nihonium (element 113), and flerovium (element 114) have known
chemical properties, and only for copernicium is there enough evidence
for a conclusive categorisation at present. The other elements may
behave differently from what would be predicted by extrapolation, due
to relativistic effects; for example, flerovium has been predicted to
possibly exhibit some noble-gas-like properties, even though it is
currently placed in the carbon group. The current experimental
evidence still leaves open the question of whether flerovium behaves
more like a metal or a noble gas.
Further periodic table extensions
Main article: Extended periodic table
Periodic table with eight rows, extended to element 172
It is unclear whether new elements will continue the pattern of the
current periodic table as period 8, or require further adaptations or
adjustments. Seaborg expected the eighth period to follow the
previously established pattern exactly, so that it would include a
two-element s-block for elements 119 and 120, a new g-block for the
next 18 elements, and 30 additional elements continuing the current
f-, d-, and p-blocks, culminating in element 168, the next noble
gas. More recently, physicists such as
Pekka Pyykkö have
theorized that these additional elements do not follow the Madelung
rule, which predicts how electron shells are filled and thus affects
the appearance of the present periodic table. There are currently
several competing theoretical models for the placement of the elements
of atomic number less than or equal to 172. In all of these it is
element 172, rather than element 168, that emerges as the next noble
gas after oganesson, although these must be regarded as speculative as
no complete calculations have been done beyond element 122.
Element with the highest possible atomic number
The number of possible elements is not known. A very early suggestion
made by Elliot Adams in 1911, and based on the arrangement of elements
in each horizontal periodic table row, was that elements of atomic
weight greater than 256± (which would equate to between elements 99
and 100 in modern-day terms) did not exist. A higher—more
recent—estimate is that the periodic table may end soon after the
island of stability, which is expected to centre around element
126, as the extension of the periodic and nuclides tables is
restricted by proton and neutron drip lines. Other predictions of
an end to the periodic table include at element 128 by John Emsley,
at element 137 by Richard Feynman, and at element 155 by Albert
Bohr model exhibits difficulty for atoms with atomic number
greater than 137, as any element with an atomic number greater than
137 would require 1s electrons to be travelling faster than c, the
speed of light. Hence the non-relativistic
Bohr model is
inaccurate when applied to such an element.
Relativistic Dirac equation
The relativistic Dirac equation has problems for elements with more
than 137 protons. For such elements, the wave function of the Dirac
ground state is oscillatory rather than bound, and there is no gap
between the positive and negative energy spectra, as in the Klein
paradox. More accurate calculations taking into account the
effects of the finite size of the nucleus indicate that the binding
energy first exceeds the limit for elements with more than 173
protons. For heavier elements, if the innermost orbital (1s) is not
filled, the electric field of the nucleus will pull an electron out of
the vacuum, resulting in the spontaneous emission of a positron;
This does not happen if the innermost orbital is filled, so that
element 173 is not necessarily the end of the periodic table.
The many different forms of periodic table have prompted the question
of whether there is an optimal or definitive form of periodic table.
The answer to this question is thought to depend on whether the
chemical periodicity seen to occur among the elements has an
underlying truth, effectively hard-wired into the universe, or if any
such periodicity is instead the product of subjective human
interpretation, contingent upon the circumstances, beliefs and
predilections of human observers. An objective basis for chemical
periodicity would settle the questions about the location of hydrogen
and helium, and the composition of group 3. Such an underlying truth,
if it exists, is thought to have not yet been discovered. In its
absence, the many different forms of periodic table can be regarded as
variations on the theme of chemical periodicity, each of which
explores and emphasizes different aspects, properties, perspectives
and relationships of and among the elements.[n 18]
Abundance of the chemical elements
Atomic electron configuration table
List of chemical elements
List of periodic table-related articles
Names for sets of chemical elements
Table of nuclides
Template:Spectral lines of the elements
The Mystery of Matter: Search for the Elements (PBS film)
Timeline of chemical element discoveries
Book: Periodic table
^ The elements discovered initially by synthesis and later in nature
are technetium (Z=43), promethium (61), astatine (85), neptunium (93),
and plutonium (94).
^ An element zero (i.e. a substance composed purely of neutrons), is
included in a few alternate presentations, for example, in the
^ There is an inconsistency and some irregularities in this
convention. Thus, helium is shown in the p-block but is actually an
s-block element, and (for example) the d-subshell in the d-block is
actually filled by the time group 11 is reached, rather than group 12.
^ The noble gases, astatine, francium, and all elements heavier than
americium were left out as there is no data for them.
^ While fluorine is the most electronegative of the elements under the
Pauling scale, neon is the most electronegative element under other
scales, such as the Allen scale.
^ While Lr is thought to have p rather than d electron in its
ground-state electron configuration, and would therefore be expected
to be volatile metal capable of forming a +1 cation in solution, no
evidence of either of these properties has been able to be obtained
despite experimental attempts to do so. It was originally expected
to have a d electron in its electron configuration and this may
still be the case for metallic lawrencium, whereas gas phase atomic
lawrencium is very likely thought to have a p electron.
^ An antecedent of Deming's 18-column table may be seen in Adams'
16-column Periodic Table of 1911. Adams omits the rare earths and the
"radioactive elements" (i.e. the actinides) from the main body of his
table and instead shows them as being "careted in only to save space"
(rare earths between Ba and eka-Yt; radioactive elements between
eka-Te and eka-I). See: Elliot Q. A. (1911). "A modification of the
periodic table". Journal of the American Chemical Society. 33(5):
^ A second extra-long periodic table row, to accommodate known and
undiscovered elements with an atomic weight greater than bismuth
(thorium, protactinium and uranium, for example), had been postulated
as far back as 1892. Most investigators considered that these elements
were analogues of the third series transition elements, hafnium,
tantalum and tungsten. The existence of a second inner transition
series, in the form of the actinides, was not accepted until
similarities with the electron structures of the lanthanides had been
established. See: van Spronsen, J. W. (1969). The periodic system of
chemical elements. Amsterdam: Elsevier. p. 315–316,
^ See The Internet database of periodic tables for depictions of these
kinds of variants.
^ But for the existence of the lanthanides the composition of group 3
would not have been a source of any special interest, since scandium,
yttrium, lanthanum and actinium exhibit the same gradual change in
properties as do calcium, strontium, barium and radium in group
^ The detachment of the lanthanides from the main body of the periodic
table has been attributed to the Czech chemist
Bohuslav Brauner who,
in 1902, allocated all of them ("Ce etc.") to one position in group 4,
below zirconium. This arrangement was referred to as the "asteroid
hypothesis", in analogy to asteroids occupying a single orbit in the
solar system. Before this time the lanthanides were generally (and
unsuccessfully) placed throughout groups I to VIII of the older
8-column form of periodic table. Although predecessors of Brauner's
1902 arrangement are recorded from as early as 1895, he is known to
have referred to the "chemistry of asteroids" in an 1881 letter to
Mendeleev. Other authors assigned all of the lanthanides to either
group 3, groups 3 and 4, or groups 2, 3 and 4. In 1922 Niels Bohr
continued the detachment process by locating the lanthanides between
the s- and d-blocks. In 1949
Glenn T. Seaborg
Glenn T. Seaborg (re)introduced the form
of periodic table that is popular today, in which the lanthanides and
actinides appear as footnotes. Seaborg first published his table in a
classified report dated 1944. It was published again by him in 1945 in
Chemical and Engineering News, and in the years up to 1949 several
authors commented on, and generally agreed with, Seaborg's proposal.
In that year he noted that the best method for presenting the
actinides seemed to be by positioning them below, and as analogues of,
the lanthanides. See: Thyssen P. and Binnemans K. (2011).
"Accommodation of the Rare Earths in the Periodic Table: A Historical
Analysis". In K. A. Gschneider Jr. (ed). Handbook on the Physics and
Chemistry of the Rare Earths. 41. Amsterdam: Elsevier, pp. 1–94;
Seaborg G. T. (1994). Origin of the
Actinide Concept'. In K. A.
Gschneider Jr. (ed). Handbook on the Physics and
Chemistry of the Rare
Earths. 18. Amsterdam: Elsevier, pp. 1–27.
^ For examples of this table see Atkins et al. (2006). Shriver &
Chemistry (4th ed.). Oxford: Oxford University Press
• Myers et al. (2004). Holt Chemistry. Orlando: Holt, Rinehart &
Winston • Chang R. (2000). Essential
Chemistry (2nd ed.). Boston:
^ For examples of the group 3 = Sc-Y-Lu-Lr table see Rayner-Canham G.
& Overton T. (2013). Descriptive Inorganic
Chemistry (6th ed.).
New York: W. H. Freeman and Company • Brown et al. (2009).
Chemistry: The Central Science (11th ed.). Upper Saddle River, New
Jersey: Pearson Education • Moore et al. (1978). Chemistry. Tokyo:
^ The phenomenon of different separation groups is caused by
increasing basicity with increasing radius, and does not constitute a
fundamental reason to show Lu, rather than La, below Y. Thus, among
the Group 2 alkaline earth metals, Mg (less basic) belongs in the
"soluble group" and Ca, Sr and Ba (more basic) occur in the "ammonium
carbonate group". Nevertheless, Mg, Ca, Sr and Ba are routinely
collocated in Group 2 of the periodic table. See: Moeller et al.
Chemistry with Inorganic Qualitative Analysis (3rd ed.).
SanDiego: Harcourt Brace Jovanovich, pp. 955–956, 958.
^ Even if metallic lawrencium has a p electron, simple modelling
studies suggest it will behave like a lanthanide, as do the rest
of the late actinides.
^ For examples of the group 3 = Ln and An table see Housecroft C. E.
& Sharpe A. G. (2008). Inorganic
Chemistry (3rd ed.). Harlow:
Pearson Education • Halliday et al. (2005). Fundamentals of Physics
(7th ed.). Hoboken, NewJersey: John Wiley & Sons • Nebergall et.
al. (1980). General
Chemistry (6th ed.). Lexington: D. C. Heath and
^ Karol (2002, p. 63) contends that gravitational effects would
become significant when atomic numbers become astronomically large,
thereby overcoming other super-massive nuclei instability phenomena,
and that neutron stars (with atomic numbers on the order of 1021) can
arguably be regarded as representing the heaviest known elements in
the universe. See: Karol P. J. (2002). "The Mendeleev–Seaborg
periodic table: Through Z = 1138 and beyond". Journal of Chemical
Education 79 (1): 60–63.
^ Scerri, one of the foremost authorities on the history of the
periodic table, favoured the concept of an optimal form of
periodic table but has recently changed his mind and now supports the
value of a plurality of periodic tables.
^ "Chemistry: Four elements added to periodic table". BBC News. 4
January 2016. Archived from the original on 4 January 2016.
^ St. Fleur, Nicholas (1 December 2016). "Four New Names Officially
Added to the Periodic Table of Elements". New York Times. Archived
from the original on 14 August 2017.
^ a b c d e f Emsley, J. (2011). Nature's Building Blocks: An A-Z
Guide to the Elements (New ed.). New York, NY: Oxford University
Press. ISBN 978-0-19-960563-7.
^ Meija, J.; et al. (2016). "
Atomic weights of the elements 2013
(IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3):
^ IUPAC 2016, Table 2, 3 combined; uncertainty removed.
^ Greenwood & Earnshaw, pp. 24–27
^ Gray, p. 6
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^ Messler, R. W. (2010). The essence of materials for engineers.
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