Copper (Cu): “Shiny, reddish copper was the first metal ever manipulated by humans, and it remains an important metal in industry today” (1). Chemistry: With an atomic number of 29, Copper (Cu) is a very unique looking metal by appearance. The exterior does hold a “red-orange metallic luster,” and copper is a “soft, malleable, and ductile metal” as well. Basically, as “a freshly exposed surface of pure copper has a reddish orange color” to it, the metal is also utilized as the pure driver to produce heat and electricity. Below is a picture of what copper looks like in its natural state. To continue being factual, Copper’s atomic number, 29, represents the number of protons in the nucleus. Its atomic weight, which is the average mass of the atom, is 63.55 grams. It has a density of 8.92 grams per cubic centimeter, and clearly, as the picture indicates, copper is solid at room temperature. The pure metal melts at 1,984.32 degrees Fahrenheit and boils at 5,301 degrees Fahrenheit. Before use, copper “must be smelted for purity,” and most often occurring is ores. De facto of Mother Nature, “natural chemical reactions do sometimes release native copper,” and this enlightens us why humans have been using copper “for at least 8,000 years” to make tools and creating new technologies. Increasing supplies with copper, “people figured out how to smelt copper by about 4500 B.C.” As advancements started happening, copper alloys were made, by adding tin, “people made a harder metal: bronze.” An intriguing statistic says “about two-thirds of the copper on Earth is found in igneous (volcanic) rocks.” According to the USGS, roughly a quarter of copper is stated in sedimentary rocks. While it is a metal that carries characteristics of being ductile and malleable, this can also explain its use in electronics and wiring. Copper is known to turn green sometimes; this is a result of an oxidation reaction. This means that it is losing elections when it is vulnerable to air and water. Like stated, “the resulting copper oxide is a dull green.” The reason the Statue of Liberty has a green appearance rather than a red-orange color is from the oxidation reaction that happened to the original copper. In accordance with the Copper Development Association, “a weathered layer of copper oxide only 0.005 inches thick coats Lady Liberty (2).” Take a look below at oxidized copper (3), and an electron configuration for copper to get a more molecular idea (2). Business: Currently, copper is listed at $3.04/lb. and equivalently $6,702.93/t. Just trading over $3 a pound, the fine metal “is up close to 28 percent year-to-date and far outperforming its five-year average from 2012 to 2016.” There are many factors that are influencing the price of the metal as we speak. As represented by the purchasing manager’s index (PMI), manufacturing activity is growing at a rate that hasn’t been witnessed in years in the U.S., Eu.

1. Copper (Cu):
“Shiny, reddish copper was the first metal ever manipulated by
humans, and it remains an important metal in industry today”
(1).
Chemistry:
With an atomic number of 29, Copper (Cu) is a very
unique looking metal by appearance. The exterior does hold a
“red-orange metallic luster,” and copper is a “soft, malleable,
and ductile metal” as well. Basically, as “a freshly exposed
surface of pure copper has a reddish orange color” to it, the
metal is also utilized as the pure driver to produce heat and
electricity. Below is a picture of what copper looks like in its
natural state.
To continue being factual, Copper’s atomic number, 29,
represents the number of protons in the nucleus. Its atomic
weight, which is the average mass of the atom, is 63.55 grams.
It has a density of 8.92 grams per cubic centimeter, and clearly,
as the picture indicates, copper is solid at room temperature.
The pure metal melts at 1,984.32 degrees Fahrenheit and boils
at 5,301 degrees Fahrenheit.

2. Before use, copper “must be smelted for purity,” and most often
occurring is ores. De facto of Mother Nature, “natural chemical
reactions do sometimes release native copper,” and this
enlightens us why humans have been using copper “for at least
8,000 years” to make tools and creating new technologies.
Increasing supplies with copper, “people figured out how to
smelt copper by about 4500 B.C.” As advancements started
happening, copper alloys were made, by adding tin, “people
made a harder metal: bronze.”
An intriguing statistic says “about two-thirds of the copper on
Earth is found in igneous (volcanic) rocks.” According to the
USGS, roughly a quarter of copper is stated in sedimentary
rocks. While it is a metal that carries characteristics of being
ductile and malleable, this can also explain its use in electronics
and wiring.
Copper is known to turn green sometimes; this is a result
of an oxidation reaction. This means that it is losing elections
when it is vulnerable to air and water. Like stated, “the
resulting copper oxide is a dull green.” The reason the Statue
of Liberty has a green appearance rather than a red-orange color
is from the oxidation reaction that happened to the original
copper. In
accordance with the Copper Development Association, “a
weathered layer of copper oxide only 0.005 inches thick coats
Lady Liberty (2).” Take a look below at oxidized copper (3),
and an electron configuration for copper to get a more
molecular idea (2).

3. Business:
Currently, copper is listed at $3.04/lb. and equivalently
$6,702.93/t. Just trading over $3 a pound, the fine metal “is up
close to 28 percent year-to-date and far outperforming its five-
year average from 2012 to 2016.” There are many factors that
are influencing the price of the metal as we speak. As
represented by the purchasing manager’s index (PMI),
manufacturing activity is growing at a rate that hasn’t been
witnessed in years in the U.S., Eurozone, and China. September
of 2017 marked the 100th straight month of expansion,
conquering a 13-year high of 60.8. Reflect the graph below to
see how copper outperformed its five-year average.
Another belief that is influencing the price of copper is the
shortages that are happening in China; despite September 2017,

4. “imports of the metal rising to its highest level since March”
2017. The world’s second-largest economy “took in 1.47
million metrics of copper ore and concentrates” in September
2017 as well, which is an amount that equates to six percent
more than the same month in 2016.
Another reason so much copper is entering China is
because of battery electric vehicles (BEVs); these demand
“three to four times as much copper as traditional fossil fuel-
powered vehicles.” While China has a tight and the most
profitable grasp on the BEV market, according to the Financial
Times, Beijing is working on putting a stop and ultimately
prohibit the retailing of fossil fuel-powered vehicles.
Nonetheless, just because of the vertical magnitude of the
Chinese market, “this move is sure to delight copper bulls and
investors in any metal that’s set to benefit from higher BEV
production (4).” 54 percent of all new car transactions by 2040
will be BEVs, according to Bloomberg New Energy Finance.
Expectedly, China, Europe, and the U.S. are accounted to make
up 60 percent of the worldwide BEV fleet. With the rise in
BEV automobiles, this predicts a huge effect on copper prices
over the span of the next ten years and more. Take a look at the
graph below that charts the driving demand for copper due to
electric vehicles in the coming years.

5. Conclusion:
Since copper occurs directly in nature, this led to very
early human use; “it was the first metal to be smelted from its
ore, the first metal to be cast into a shape in a mold, and the
first metal to be purposefully alloyed with another metal (2).”
The characteristics of this metal made it so versatile for early
humans to make tools and get jobs done. It is truly amazing
how a metal like copper revolutionized technology for humans
and brought so much innovation and opportunities.
As there is a current market for copper today, it is getting
traded at just over $3 a pound, and it is used today in
electronics and wiring. Also, copper is way outperforming its
five-year average, and the need for copper in electric vehicles
exponentially rises for the next ten years, continued to 2040.
This means that the demand for copper is not slowing down, and
if anything, the price will go up because the demand for the
copper has gone up as well. Copper plays a vital and also low-
key factor in our economy, and the green on your pennies
symbolize the oxidization that has occurred over time to your
copper penny.
Works Cited (Sources)

6. 1. https://www.livescience.com/29377-copper.html
2. https://en.wikipedia.org/wiki/Copper
3. https://www.youtube.com/watch?v=JoO8TbXebls
4. http://www.businessinsider.com/copper-is-the-metal-of-the-
future-2017-10
CHE 102: LECTURE 2 Metals, Minerals and Money
The Illiad is an ancient Greek epic poem, traditionally
attributed to Homer. Set during the Trojan War, the ten-year
siege of the city of Troy (a city in Turkey, the modern-day
Hisarlik) by a coalition of Greek states, it tells of battles and
events during the weeks of a quarrel between King Agamemnon
and the celebrated warrior Achilles. Above is the death mask of
Agamemnon, cast in GOLD, created ~ 1550–1500 B.C. and
discovered in 1876 in Mycenae, Greece by Heinrich
Schliemann. It has been referred to as the “Mona Lisa of
Prehistory.”
Metals and minerals form some of the most beautiful crystals in
nature. For example,
Gold as it is found in Nature
has the following structure at the atomic level:

7. A mineral , iron pyrite, infamous in the California Gold Rush [
that began on January 24, 1848, when gold was found by James
W. Marshall at Sutter's Mill in Coloma, California ] was known
as “ Fool’s Gold.” The chemical formula is Fe2S [ Iron↔Fe,
Sulphur ↔ S ].
Its crystal structure is very different from that of Gold.
The element Sodium, Na, has the following atomic crystal
structure. The structure
is termed body-centered cubic.
The ionic compound, Halite or table salt, NaCl, has the atomic
structure below. Here,
the smaller gray balls are sodium ions, Na+, and the larger
green balls are chlorine ions,
Cl-. Notice that the Cl- ions are larger than the Na+ ion, an
important distinction which
will be explained later in the course. The structure of the
mineral is octahedral.
An example of an element that forms molecular (as opposed to
ionic) crystals is Carbon.
There are four “flavors” of Carbon found in Nature; these are
called allotropes. The two
most common allotropes are diamond and graphite, which have
very different atomic structures. Diamond has a tetrahedral
structure whereas graphite has a structure like stacked sheets of
chicken wire (planar sheets of hexagons).

8. ALLOYS
Different metals can form solid state mixtures of variable
composition, called alloys. Alloys are analogous to liquid
solutions, e.g. salt dissolved in water. There are two types of
alloys, substitutional and interstitial.
When a molten metal is mixed with another substance, there are
two mechanisms that can cause an alloy to form, called atom
exchange and the interstitial mechanism. The relative size of
each element in the mix plays a primary role in determining
which mechanism will occur. When the atoms are relatively
similar in size, the atom exchange method usually happens,
where some of the atoms composing the metallic crystals are
substituted with atoms of the other constituent. This is called
a substitutional alloy. Examples of substitutional alloys include
BRONZE and BRASS, in which some of the copper atoms are
substituted with either tin or zinc atoms respectively.
In the case of the interstitial mechanism, one atom is usually
much smaller than the other and can not successfully substitute
for the other type of atom in the crystals of the base metal.
Instead, the smaller atoms become trapped in the spaces
between the atoms of the crystal matrix, called the interstices.
This is referred to as an interstitial alloy. STEEL is an example
of an interstitial alloy, because the very small carbon atoms fit
into interstices of the iron matrix.
STAINLESS STEEL is an example of a combination of
interstitial and substitutional alloys, because the carbon atoms
fit into the interstices, but some of the iron atoms are

9. substituted by nickel and chromium atoms.
2. Materials in Nature:
Materials in nature can exist in three phases, as a solid, a liquid
and a gas. The state
observed depends on the pressure and temperature. So, for
example, water can pass
from a solid (ice), to a liquid (water) to a gas (steam) by raising
the temperature.
Moreover, the temperature at which a particular “phase
transition” occurs is very
specific. For water, ice forms at exactly 0o Centigrade, and gas
appears at exactly
100oC. These signatures are very specific for every substance
found in nature; they
are like a “fingerprint” and there are volumes in libraries and
data banks with fingerprints for thousands (+) of substances.
Importantly, solids are the most ordered state of matter (see
above), liquids
more disorganized, and gases are totally disorganized. Atoms
in a solid are
locked in place, atoms in a liquid can break free and move
around in the same
volume, and the motion of atoms in a gas is totally random,
with the gas occupying
not only the original volume of the precursor liquid phase, but
all the nooks and
crannies of a container.
Phase transitions, such as the ones cited above for water, occur
because heat

10. (which is energy) is either added to or subtracted from the
physical system.
A driving force for a transformation to occur is the tendency of
all systems in
Nature to try to be in the most disorganized state possible. This
tendency
to become disorganized (or more random) is characterized by
the increase
in a property called the entropy.
The drive, or perhaps better the drift, of any system in Nature to
reach a stable
state is therefore a consequence of two factors. One is the
tendency toward the
most stable, lowest energy state. The other is the tendency
toward maximum
disorder (maximum entropy)
These two observations have been canonized in by two ,
fundamental generalizations,
called Conservation Laws. Textbook statements of these two
laws follow:
First Law of Thermodynamics
The Law of Conservation of Energy states that the
total energy of an isolated system is constant; energy can be
transformed from one form to another, but can be neither
created nor destroyed.
Second Law of Thermodynamics
This law states that the total entropy of an isolated system can

11. never decrease over time. The total entropy can remain constant
in ideal cases where the system is in a steady state
(equilibrium), or is undergoing a reversible process. In
all spontaneous processes, the total entropy increases and the
process is irreversible. The increase in entropy accounts for the
irreversibility of natural processes, and the asymmetry between
future and past time.
3. The Periodic Table
Over the course of centuries, and particularly in the late Middle
Ages, new substances were discovered (for example,
phosphorus) that expanded the list of elements known since
antiquity.
With the publication in 1800 of Volta’s discovery of the battery,
a plethora of new
elements and new perspectives were discovered/reported. These
led, eventually,
to the publication in 1869 of Mendeleev’s “periodic table”
which is the single
most important organizing principle in Chemistry.
The periodic table of elements is organized so that one can
quickly discern the properties of individual elements such as
their mass, electron number, electron configuration and their
unique chemical properties. Metals reside on the left side of the
table, while non-metals reside on the right.
The foundational study on the periodicity of properties of
elements was reported by the Russian chemist Mendeleev (1834
– 1907). He used the Periodic Law not only to correct the then-
accepted properties of some known elements, such as the
valence and atomic weight of uranium, but also to predict the
properties of eight elements that were yet to be discovered. His

12. original Periodic Table, reported in 1869, follows.
The modern version is below.
At the end of this file I have reproduced the Wikipedia website
on the Periodic Table.
It is worth spending time on this website, as it will be useful
throughout the course.
4. Why GOLD ???

13. From the desk mask of Tutankhamun (below) to that of
Agamemnon (see top of file),
from jewelry
[ A Moche gold necklace depicting feline heads. Larco Museum
Collection, Lima, Peru ]
to the representation of religious or cultural themes [See below
the Musica raft, ~600-1600 AD. The figure refers to the
ceremony of the legend of El Dorado. The zipa used to cover
his body in gold dust, and from his raft, he offered treasures to
the Guatavita goddess in the middle of the sacred lake. This old
Muisca tradition became the origin of the legend of El Dorado.
On display in the Gold Museum in Bogota, Columbia.]
to money,
[ Gold coin of Eucraides I (171–145 BC), one of the Hellenistic
rulers of ancient Ai-Khanoum. This is the largest known gold
coin minted in antiquity (169,20 g; 58 mm) ]
GOLD has been a benchmark of wealth for the World’s
civilizations.

14. [ An Indian tribute-bearer at Apadana, from the Achaemenid
satrapy of Hindush, carrying gold on a yoke, circa 500 BC. ]
Gold can stimulate a subjective personal experience, but gold
can also be objectified if it's adopted as a system of exchange.
This duplicity is a conundrum that is unique to gold as
a commodity. Gold can be something quantitative and tangible,
like money, and at the same time, it can embody something
ephemeral, like a feeling, even a host of feelings. So, part of
the reason that gold has always had value lies in the psychology
and nature of the human experience.
But, why?
Well, for openers it never changes its appearance. The death
mask of Agamemnon looks
the same today as the day it was cast. Chemically, this is a
profound attribute not shared
by any other element to the same degree. For example, Silver is
a beautiful metal, but can
tarnish. Gold is inert chemically to gases/liquids/solids in its
environment.
Secondly, it is the most malleable of elements. Physically, this
is an attribute unmatched by any other element. A gold nugget
of 0.5 cm (0.20 in) in size can be hammered into a gold foil of
about 0.5 m2 (5.4 sq ft). It is this property which accounts for
the use of Gold since antiquity for jewelry and the beautiful
objects created by gifted craftsmen.
The central role of Gold in our monetary system is more
problematical. Hunter Gatherer societies functioned without

15. money. Such societies were characterized by Thomas Hobbes [
English philosopher and political theorist best known for his
book “Leviathan” (1651) ], as“solitary, poor, nasty, brutish
and short.”
But, it is important to note that sophisticated societies have also
functioned without money.
Five hundred years ago, the most sophisticated society in South
America, the Inca Empire, was moneyless.
Labor was the unit of value in the Inca Empire, just as it was
supposed to be in a
Communist society. Recall Karl Marx (1818-1883), who
sought to demonstrate in
“Das Kaptial,” that money is commoditized labor the surplus
generated by honest toil, appropriated and ‘reified’ in order to
satisfy the capitalist class’s insatiable lust for accumulation.
If, however, one adopts a substance (say, an element) as the
basis of a monetary system, what are the chemical and physical
properties that are desirable? Sanat Kumar [Chair of the
Department of Chemical Engineering at Columbia University]
has provided insight. He focuses on four qualities that an
element must meet to stand alone as a currency.
First, it can't be a gas — gases simply are not practical for
currency exchange. That knocks out a bunch of contenders from
the right side of the periodic table, including the Noble gases,
which would meet the other three qualifications.
Second, it can't be corrosive or reactive — pure lithium (Li), for
example, ignites when exposed to water or air. Iron (Fe) rusts.
This qualification knocks out 38 elements.
Third, it can't be radioactive. For one thing, your money would

16. eventually radiate away to nothing. For another, the radiation
would eventually kill you. This eliminates the two rows that are
separate from table, the elements known as actinides and
lanthanides.
Any of the 30 or so remaining elements would make nice, stable
forms of currency if they met the fourth qualification: They
must be rare enough to be valuable, but not so rare that it’s
impossible to find them.
That brings us to five elements: rhodium (Rh), palladium (Pd) ,
platinum (Pt), silver (Ag) and gold (Au).
Although silver has been used for currency, it tarnishes easily,
so it's out. Rhodium and Palladium were discovered only in the
1800s, so they'd have been of no use to early civilizations. That
leaves Gold and Platinum. Platinum, however, has a melting
point around 3,000 degrees Fahrenheit (about 1,600 degrees
Celsius), which could only be attained in a modern furnace, so
early civilizations would not have been able to conveniently
shape it into uniform units.
That leaves Gold, which is solid but malleable, doesn't react,
and won't kill you.
So what does this process of elemental elimination tell us about
what makes a good currency?
A currency needs to be stable, portable and non-toxic. And it
needs to be fairly rare - you might be surprised just how little
gold there is in the world.
If you were to collect together every earring, every gold
sovereign, the tiny traces gold in every computer chip, every
pre-Columbian statuette, every wedding ring and melt it down,
it's guesstimated that you'd be left with approximately one
volume of (20 meter)^3 . In English units, 20 meters = 21.87
yards. Think of the “20 yard line” on a football field;
a volume ~20 yards on each of three sides is the estimate given.
But scarcity and stability aren't the whole story. Gold has one

17. other quality that makes it the stand-out contender for currency
in the periodic table. Gold is... golden.
All the other metals in the periodic table are silvery-colored
except for copper.
BUT copper corrodes, turning green when exposed to moist air.
That makes gold very distinctive.
At the end of the day, that's the secret of gold's success as a
currency. It satisfies the
advantages and attributes of “MONEY,” and it is “beautiful.”
Check out the spot value of Gold in today’s “coronavirus”
environment. Gold and Palladium are the most actively traded
metals in the market. In today’s Wall
Street Journal, Gold is valued at $1578.20 a troy ounce.
Periodic table
From Wikipedia, the free encyclopedia
Jump to navigationJump to search
This article is about the table used in chemistry and physics.
For other uses, see Periodic table (disambiguation).
The periodic table, also known as the periodic table of
elements, is a tabular display of the chemical elements, which
are arranged by atomic number, electron configuration, and
recurring chemical properties. The structure of the table
shows periodic trends. The seven rows of the table,
called periods, generally have metals on the left
and nonmetals on the right. The columns, called groups, contain
elements with similar chemical behaviours. Six groups have
accepted names as well as assigned numbers: for example,
group 17 elements are the halogens; and group 18 are the noble
gases. Also displayed are four simple rectangular areas
or blocks associated with the filling of different atomic orbitals.

18. The elements from atomic numbers 1 (hydrogen) through 118
(oganesson) have been discovered or synthesized, completing
seven full rows of the periodic table.[1][2] The first 94
elements, hydrogen through plutonium, all occur naturally,
though some are found only in trace amounts and a few were
discovered in nature only after having first been synthesized.[n
1] Elements 95 to 118 have only been synthesized in
laboratories or nuclear reactors.[3] 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 candidates for this extension.
Numerous synthetic radioisotopes of naturally occurring
elements have also been produced in laboratories.
The organization of the periodic table can be used to derive
relationships between the various element properties, and also
to predict chemical properties and behaviours of undiscovered
or newly synthesized elements. Russian chemist Dmitri
Mendeleev published the first 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
the 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 the
development of 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 used in chemistry, nuclear physics and other sciences.
Some discussion remains ongoing regarding the placement and
categorisation of specific elements, the future extension and
limits of the table, and whether there is an optimal form of the
table.
Part of a series on the
Periodic table
Periodic table forms[show]
Periodic table history[show]

19. Sets of elements
By periodic table structure[show]
By metallic classification[show]
By other characteristics[show]
Elements
List of chemical elements[show]
Properties of elements[show]
Data pages for elements[show]
· Book
· Category
· Chemistry Portal
· v
· t
· e
Contents
· 1Overview
· 2Grouping methods
· 2.1Groups
· 2.2Periods
· 2.3Blocks
· 2.4Metals, metalloids and nonmetals
· 3Periodic trends and patterns
· 3.1Electron configuration
· 3.2Atomic radii
· 3.3Ionization energy
· 3.4Electronegativity
· 3.5Electron affinity
· 3.6Metallic character
· 3.7Oxidation number
· 3.8Linking or bridging groups
· 4History
· 4.1First systemization attempts
· 4.2Mendeleev's table
· 4.3Second version and further development
· 5Different periodic tables
· 5.1The long- or 32-column table

20. · 5.2Tables with different structures
· 6Open questions and controversies
· 6.1Placement of hydrogen and helium
· 6.2Group 3 and its elements in periods 6 and 7
· 6.2.1Lanthanum and actinium
· 6.2.2Lutetium and lawrencium
· 6.2.3Lanthanides and actinides
· 6.3Groups included in the transition metals
· 6.4Elements with unknown chemical properties
· 6.5Further periodic table extensions
· 6.6Element with the highest possible atomic number
· 6.6.1Bohr model
· 6.6.2Relativistic Dirac equation
· 6.7Optimal form
· 7Other
· 8See also
· 9Notes
· 10References
· 10.1Bibliography
· 11Further reading
· 12External linksOverview
· v
· t
· e
Periodic table
Group
1
2
3
4
5
6
7
8
9

21. 10
11
12
13
14
15
16
17
18
Alkali metals
Alkaline earth metals
Pnictogens
Chalcogens
Halogens
Noble gases
Period
1
Hydrogen1H1.008
Helium2He4.0026
2
Lithium3Li6.94

22. Beryllium4Be9.0122
Boron5B10.81
Carbon6C12.011
Nitrogen7N14.007
Oxygen8O15.999
Fluorine9F18.998
Neon10Ne20.180
3
Sodium11Na22.990
Magnesium12Mg24.305
Aluminium13Al26.982
Silicon14Si28.085
Phosphorus15P30.974
Sulfur16S32.06
Chlorine17Cl35.45
Argon18Ar39.95
4
Potassium19K39.098
Calcium20Ca40.078
Scandium21Sc44.956
Titanium22Ti47.867
Vanadium23V50.942
Chromium24Cr51.996
Manganese25Mn54.938
Iron26Fe55.845
Cobalt27Co58.933
Nickel28Ni58.693
Copper29Cu63.546
Zinc30Zn65.38
Gallium31Ga69.723
Germanium32Ge72.630
Arsenic33As74.922
Selenium34Se78.971

23. Bromine35Br79.904
Krypton36Kr83.798
5
Rubidium37Rb85.468
Strontium38Sr87.62
Yttrium39Y88.906
Zirconium40Zr91.224
Niobium41Nb92.906
Molybdenum42Mo95.95
Technetium43Tc[97]
Ruthenium44Ru101.07
Rhodium45Rh102.91
Palladium46Pd106.42
Silver47Ag107.87
Cadmium48Cd112.41
Indium49In114.82
Tin50Sn118.71
Antimony51Sb121.76
Tellurium52Te127.60
Iodine53I126.90
Xenon54Xe131.29
6
Caesium55Cs132.91
Barium56Ba137.33
Lanthanum57La138.91
Hafnium72Hf178.49
Tantalum73Ta180.95
Tungsten74W183.84
Rhenium75Re186.21
Osmium76Os190.23
Iridium77Ir192.22
Platinum78Pt195.08
Gold79Au196.97
Mercury80Hg200.59

24. Thallium81Tl204.38
Lead82Pb207.2
Bismuth83Bi208.98
Polonium84Po[209]
Astatine85At[210]
Radon86Rn[222]
7
Francium87Fr[223]
Radium88Ra[226]
Actinium89Ac[227]
Rutherfordium104Rf[267]
Dubnium105Db[268]
Seaborgium106Sg[269]
Bohrium107Bh[270]
Hassium108Hs[269]
Meitnerium109Mt[278]
Darmstadtium110Ds[281]
Roentgenium111Rg[282]
Copernicium112Cn[285]
Nihonium113Nh[286]
Flerovium114Fl[289]
Moscovium115Mc[290]
Livermorium116Lv[293]
Tennessine117Ts[294]
Oganesson118Og[294]
Cerium58Ce140.12
Praseodymium59Pr140.91
Neodymium60Nd144.24
Promethium61Pm[145]

25. Samarium62Sm150.36
Europium63Eu151.96
Gadolinium64Gd157.25
Terbium65Tb158.93
Dysprosium66Dy162.50
Holmium67Ho164.93
Erbium68Er167.26
Thulium69Tm168.93
Ytterbium70Yb173.05
Lutetium71Lu174.97
Thorium90Th232.04
Protactinium91Pa231.04
Uranium92U238.03
Neptunium93Np[237]
Plutonium94Pu[244]
Americium95Am[243]
Curium96Cm[247]
Berkelium97Bk[247]
Californium98Cf[251]
Einsteinium99Es[252]
Fermium100Fm[257]
Mendelevium101Md[258]
Nobelium102No[259]
Lawrencium103Lr[266]
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 1 atm)
PrimordialFrom decaySyntheticBorder shows natural occurrence
of the element
Standard atomic weightAr, std(E)[4]
· Ca: 40.078 — Formal short value, rounded (no uncertainty)[5]
· Po: [209] — mass number of the most stable isotope

26. Background color shows subcategory in the metal–metalloid–
nonmetal trend:
Metal
Metalloid
Nonmetal
Unknown
chemical
properties
Alkali metal
Alkaline earth metal
Lanthanide
Actinide
Transition metal
Post-transition metal
Reactive nonmetal
Noble gas
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.[7]
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

27. 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.[8]
Since 2016, 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 made official in November 2016 by
IUPAC.[9][10][11][12]
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.[3] 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).[13]Grouping methods
Groups
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

28. shell.[14] Consequently, elements in the same group tend to
have a shared chemistry and exhibit a clear trend in properties
with increasing atomic number.[15] 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 similarities.[16][17][18]
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).[19] 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.[20]
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),[19] 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.

29. Similarly, a group has a top-to-bottom decrease in
electronegativity due to an increasing distance between valence
electrons and the nucleus.[21] There are exceptions to these
trends: for example, in group 11, electronegativity increases
farther down the group.[22]
hide
· v
· t
· e
Groups in the Periodic table
IUPAC group
1a
2
3b
n/ab
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
Mendeleev (I–VIII)
I
II
III
IV

30. V
VI
VII
VIII
I
II
III
IV
V
VI
VII
c
CAS (US, A-B-A)
IA
IIA
IIIB
IVB
VB
VIB
VIIB
VIIIB
IB
IIB
IIIA
IVA
VA
VIA
VIIA
VIIIA
old IUPAC (Europe, A-B)
IA
IIA
IIIA
IVA

31. VA
VIA
VIIA
VIII
IB
IIB
IIIB
IVB
VB
VIB
VIIB
0
Trivial name
H and Alkali metalsr
Alkaline earth metalsr
Coinage metals
Triels
Tetrels
Pnictogensr
Chalcogensr
Halogensr
Noble gasesr
Name by elementr
Lithium group
Beryllium group
Scandium group

32. Titanium group
Vanadium group
Chromium group
Manganese group
Iron group
Cobalt group
Nickel group
Copper group
Zinc group
Boron group
Carbon group
Nitrogen group
Oxygen group
Fluorine group
Helium or Neon group
Period 1
H
He

33. Period 2
Li
Be
B
C
N
O
F
Ne
Period 3
Na
Mg
Al
Si

34. P
S
Cl
Ar
Period 4
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Period 5
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd

35. Ag
Cd
In
Sn
Sb
Te
I
Xe
Period 6
Cs
Ba
La
Ce–Lu
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
Period 7
Fr
Ra
Ac
Th–Lr
Rf
Db
Sg

36. Bh
Hs
Mt
Ds
Rg
Cn
Nh
Fl
Mc
Lv
Ts
Og
a Group 1 is composed of hydrogen (H) and the alkali metals.
Elements of the group have one s-electron in the outer electron
shell. Hydrogen is not considered to be an alkali metal as it
rarely exhibits behaviour comparable to theirs, though it is more
analogous to them than any other group. This makes the group
somewhat exceptional.
n/a Do not have a group number
b 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.[23]
c 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.
Periods
Main article: Period (periodic table)
A period is a horizontal row in the periodic table. Although

37. 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.[24]
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.[25] 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.[21] 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 gases.[26]
Blocks
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. Elements are assigned to
blocks by what orbitals their valence electrons or vacancies lie
in.[27] 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

38. the transition metals. The f-block, often offset below the rest of
the periodic table, has no group numbers and comprises most of
the lanthanides and actinides. A hypothetical g-block is
expected to begin around element 121, a few elements away
from what is currently known.[28]
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 elements.
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.[29]
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 post-transition metals.
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, monatomic noble gases. Specialized groupings
such as refractory metals and noble metals, are examples of
subsets of transition metals, also known[30] and occasionally
denoted.[31]

39. 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.[32] 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 metal. 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[33] 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
Electron configuration
Main article: Electron 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

40. occupies a new shell corresponds to the start of each new
period, these positions being occupied by hydrogen and
the alkali metals.[34][35]
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.[34][35]
Atomic radii
Main article: Atomic radius
Atomic number plotted against atomic radius[n 3]
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 theory.[36]
The electrons in the 4f-subshell, which is progressively filled
from lanthanum (element 57) to ytterbium (element 70),[38] 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

41. are smaller than would be expected and that are almost identical
to the atomic radii of the elements immediately above
them.[39] Hence lutetium has virtually the same atomic radius
(and chemistry) as yttrium, 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 an effect
of the lanthanide contraction: a similar actinide contraction also
exists. 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.[40] 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.[39]
Ionization energy
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 table.[40]
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

42. jumps occur in the ionization energies of other third-row
atoms.[40]
Electronegativity
Main article: Electronegativity
Graph showing increasing electronegativity with growing
number of selected groups
Electronegativity is the tendency of an atom to attract a shared
pair of electrons.[41] 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.[42] 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 4] while caesium is the least,
at least of those elements for which substantial data is
available.[22]
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.[22] The anomalously high
electronegativity of lead, particularly when compared
to thallium and bismuth, is an artifact of electronegativity
varying with oxidation state: its electronegativity conforms
better to trends if it is quoted for the +2 state instead of the +4
state.[43]
Electron affinity
Main article: Electron affinity

43. Dependence of electron affinity on atomic number.[44] 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.[45]
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 metals. 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 values.[46]
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 stable.[46]
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.[47]
Metallic character
The lower the values of ionization energy, electronegativity and
electron affinity, the more metallic character the element has.

44. Conversely, nonmetallic character increases with higher values
of these properties.[48] 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,[49] 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.[50][51]
Oxidation number
With some minor exceptions, oxidation numbers among the
elements show four main trends according to their periodic table
geographic location: left; middle; right; and south. On the left
(groups 1 to 3), the highest most stable oxidation number is the
group number, with lower oxidation states being less stable. In
the middle (groups 4 to 11), higher oxidation states become
more stable going down each group. Group 12 is an exception to
this trend; they behave as if they were located on the left side of
the table. On the right, higher oxidation states tend to become
less stable going down a group.[52] The shift between these
trends is continuous: for example, group 3 also has lower
oxidation states most stable in its lightest member (scandium,
with CsScCl3 for example known in the +2 state),[53] and
group 12 is predicted to have copernicium more readily showing
oxidation states above +2.[54]
The lanthanides and actinides positioned along the south of the
table are distinguished by having the +3 oxidation state in
common; this is the most stable state for the lanthanides. The
early actinides show a pattern of oxidation states somewhat

45. similar to those of their period 6 and 7 transition metal
congeners; the later actinides are more similar to the
lanthanides.[55]
Linking or bridging groups
Sc, Y, La, Ac, Zr, Hf, Rf, Nb, Ta, Db, Lu, Lr, Cu, Ag, Au, Zn,
Cd, Hg, He, Ne, Ar, Kr, Xe, Rn
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Sodium
Magnesium
Aluminium
Silicon
Phosphorus
Sulfur
Chlorine
Argon
Potassium
Calcium
Scandium
Titanium
Vanadium
Chromium
Manganese

46. Iron
Cobalt
Nickel
Copper
Zinc
Gallium
Germanium
Arsenic
Selenium
Bromine
Krypton
Rubidium
Strontium
Yttrium
Zirconium
Niobium
Molybdenum
Technetium
Ruthenium
Rhodium
Palladium
Silver
Cadmium
Indium
Tin
Antimony
Tellurium
Iodine
Xenon
Caesium
Barium
Lanthanum
Cerium
Praseodymium

47. Neodymium
Promethium
Samarium
Europium
Gadolinium
Terbium
Dysprosium
Holmium
Erbium
Thulium
Ytterbium
Lutetium
Hafnium
Tantalum
Tungsten
Rhenium
Osmium
Iridium
Platinum
Gold
Mercury (element)
Thallium
Lead
Bismuth
Polonium
Astatine
Radon
Francium
Radium
Actinium
Thorium
Protactinium
Uranium
Neptunium
Plutonium
Americium

48. Curium
Berkelium
Californium
Einsteinium
Fermium
Mendelevium
Nobelium
Lawrencium
Rutherfordium
Dubnium
Seaborgium
Bohrium
Hassium
Meitnerium
Darmstadtium
Roentgenium
Copernicium
Nihonium
Flerovium
Moscovium
Livermorium
Tennessine
Oganesson
32-column periodic table showing, from left to right, the
location of group 3; the heavy group 4 and 5 elements; 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. In general, groups at the peripheries of blocks
display similarities to the groups of the neighbouring blocks as
well as to the other groups in their own blocks, as expected as
most periodic trends are continuous.[56] These groups, like the
metalloids, show properties in between, or that are a mixture of,
groups to either side. Chemically, the group 3 elements,
lanthanides, and heavy group 4 and 5 elements show some

49. behaviour similar to the alkaline earth metals[57] or, more
generally, s block metals[58][59][60] but have some of the
physical properties of d block transition metals;[61] meanwhile,
lutetium behaves chemically as a lanthanide (with which it is
often classified) but shows a mix of lanthanide and transition
metal physical properties (as does
yttrium).[62][63] Lawrencium, as an analogue of lutetium,
would presumably display like characteristics.[n 5] The coinage
metals in group 11 (copper, silver, and gold) are chemically
capable of acting as either transition metals or main group
metals.[66] 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.[67][68] 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.[56]History
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-)
In 1789, Antoine Lavoisier published a list of 33 chemical
elements, grouping them into gases, metals, nonmetals,
and earths.[69] Chemists spent the following century searching
for a more precise classification scheme. In 1829, 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.[70] This became known as the Law of
Triads.[71] German chemist Leopold Gmelin worked with this

50. 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.[70] 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.[72] This concept
eventually became known as valency, where different elements
bond with different numbers of atoms.[73]
In 1862, the French geologist Alexandre-Émile Béguyer de
Chancourtois published an early form of the 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.[74]
Julius Lothar Meyer's periodic table, published in "Die
modernen Theorien der Chemie" (1864)[75]
In 1864, Julius Lothar Meyer, a German chemist, published a
table with 28 elements. Realizing that an arrangement according
to atomic weight did not exactly fit the observed periodicity in
chemical properties he gave valency priority over minor
differences in atomic weight. A missing element between Si and
Sn was predicted with atomic weight 73 and valency
4.[75] Concurrently, English chemist William Odling 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

51. groupings".[76] Odling alluded to the idea of a periodic law but
did not pursue it.[77] He subsequently proposed (in 1870) a
valence-based classification of the elements.[78]
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.[79][80] This so
termed Law of Octaves was ridiculed by Newlands'
contemporaries, and the Chemical Society refused to publish his
work.[81] Newlands was nonetheless able to draft a table of the
elements and used it to predict the existence of missing
elements, such as germanium.[82] The Chemical Society only
acknowledged the significance of his discoveries five years
after they credited Mendeleev.[83]
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.[84][85]
Mendeleev's table
Periodic table of elements. Vienna, 1885. University of St
Andrews
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

52. 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.[86] Mendeleev's
table, dated March 1 [O.S. February 17] 1869,[87] was his first
published version. That of Meyer was an expanded version of
his (Meyer's) table of 1864.[88] 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 repeat.[89]
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.[90] 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.[91] 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.
Mendeleev published in 1869, using atomic weight to organize
the elements, information determinable to fair precision in his
time. Atomic weight worked well enough to allow Mendeleev to
accurately predict the properties of missing elements.
Mendeleev took the unusual step of naming missing elements
using the Sanskrit numerals eka (1), dvi (2), and tri (3) to
indicate that the element in question was one, two, or three rows
removed from a lighter congener. It has been suggested that
Mendeleev, in doing so, was paying homage to ancient Sanskrit
grammarians, in particular Pāṇini, who devised a periodic
alphabet for the language.[92]
Henry Moseley (1887–1915)
Following the discovery of the atomic nucleus by Ernest
Rutherford in 1911, it was proposed that the integer count of the

53. nuclear charge is identical to the sequential place of each
element in the periodic table. In 1913, English physicist Henry
Moseley using X-ray spectroscopy confirmed this proposal
experimentally. Moseley determined the value of the nuclear
charge of each element and showed that Mendeleev's ordering
actually places the elements in sequential order by nuclear
charge.[93] Nuclear charge is identical to proton count and
determines the value of the atomic number (Z) of each element.
Using atomic number gives a definitive, integer-based sequence
for the elements. 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 atomic number is the absolute definition of
an element and gives a factual basis for the ordering of the
periodic table.[94]
Second version and further development
Mendeleev's 1871 periodic table with eight groups of elements.
Dashes represented elements unknown in 1871.
Eight-group 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.[95] These gaps were
subsequently filled as chemists discovered additional naturally
occurring elements.[96] It is often stated that the last naturally
occurring element to be discovered was francium (referred to by
Mendeleev as eka-caesium) in 1939, but it was technically only
the last element to be discovered in nature as opposed to by
synthesis.[97]Plutonium, produced synthetically in 1940, was
identified in trace quantities as a naturally occurring element in

54. 1971.[98]
The popular[99] 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.[100][n
6] 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 encyclopedias of chemistry. It was
also distributed for many years by the Sargent-Welch Scientific
Company.[101][102][103]
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.[104]
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.[105][106][n 7]
Although minute quantities of some transuranic elements occur
naturally,[3] they were all first discovered in laboratories. Their
production has expanded the periodic table significantly, the
first of these being neptunium, synthesized in

55. 1939.[107] 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.[108] 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.[109]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, as
proposed by Alfred Werner.[110] Unlike the 18-column form
this arrangement results in "no interruptions in the sequence of
increasing atomic numbers".[111] The relationship of the f-
block to the other blocks of the periodic table also becomes
easier to see.[112]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.[113] 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,[114] and the familiarity of
chemists with the modern form, as introduced by Seaborg.[115]
show
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56. · e
Periodic table (large cells, 32-column layout)
Tables with different structures
Main article: Alternative periodic tables
Within 100 years of the appearance of Mendeleev's table in
1869, Edward G. Mazurs had collected an estimated 700
different published versions of the periodic
table.[113][118][119] As well as numerous rectangular
variations, other periodic table formats have been shaped, for
example,[n 8] like a circle, cube, cylinder, building,
spiral, lemniscate,[120] 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 tables.[119]
Theodor Benfey's spiral periodic table
A popular[121] alternative structure is that of Otto Theodor
Benfey (1960). The elements are arranged in a continuous
spiral, with hydrogen at the centre and the transition metals,
lanthanides, and actinides occupying peninsulas.[122]
Most periodic tables are two-dimensional;[3] 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),[123] Wringley's Lamina System (1949),[124]Giguère's
Periodic helix (1965)[125] and Dufour's Periodic Tree
(1996).[126] Going one further, Stowe's Physicist's Periodic
Table (1989)[127] has been described as being four-dimensional
(having three spatial dimensions and one colour
dimension).[128]
The various forms of periodic tables can be thought of as lying
on a chemistry–physics continuum.[129] Towards the chemistry
end of the continuum can be found, as an example, Rayner-
Canham's "unruly"[130] Inorganic Chemist's Periodic Table
(2002),[131] which emphasizes trends and patterns, and unusual

57. 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.[132] A somewhat similar approach has been taken
by Alper,[133] albeit criticized by Eric Scerri as disregarding
the need to display chemical and physical
periodicity.[134] 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.[135] 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.[77][136]
hide
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· e
Left-step periodic table (by Charles Janet)
f1
f2
f3
f4
f5
f6
f7
f8
f9
f10
f11
f12
f13
f14

58. d1
d2
d3
d4
d5
d6
d7
d8
d9
d10
p1
p2
p3
p4
p5
p6
s1
s2
1s
H
He
2s
Li
Be
2p 3s
B
C
N
O
F
Ne
Na
Mg

59. 3p 4s
Al
Si
P
S
Cl
Ar
K
Ca
3d 4p 5s
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Rb
Sr
4d 5p 6s
Y
Zr
Nb
Mo

60. Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Cs
Ba
4f 5d 6p 7s
La
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
Hf
Ta
W
Re
Os
Ir

61. Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
Fr
Ra
5f 6d 7p 8s
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
Cn

62. Nh
Fl
Mc
Lv
Ts
Og
119
120
f-block
d-block
p-block
s-block
This form of periodic table is 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). The experimentally
determined ground-state electron configurations of the elements
differ from the configurations predicted by the Madelung rule in
twenty instances, but the Madelung-predicted configurations are
always at least close to the ground state. The last two elements
shown, elements 119 and 120, have not yet been
synthesized.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).[137] 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

63. 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.[138]
The position of hydrogen in group 1 is reasonably well settled.
Its usual oxidation state is +1 as is the case for its heavier alkali
metal congeners. Like lithium, it has a significant covalent
chemistry.[139][140] It can stand in for alkali metals in typical
alkali metal structures.[141] It is capable of forming alloy-like
hydrides, featuring metallic bonding, with some transition
metals.[142]
Nevertheless, it is sometimes placed elsewhere. A common
alternative is at the top of group 17[134] 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.[143] 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.[144] Finally, hydrogen is
sometimes placed separately from any group; this is based on its
general properties being regarded as sufficiently different from
those of the elements in any other group.
The other period 1 element, helium, is most often placed in
group 18 with the other noble gases, as its extraordinary
inertness is extremely close to that of the other light noble
gases neon and argon.[145] Nevertheless, it is occasionally
placed separately from any group as well.[146] The property
that distinguishes helium from the rest of the noble gases 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. Some authors, such as Henry Bent (the eponym

64. of Bent's rule), Wojciech Grochala, and Felice Grandinetti, have
argued that helium would be correctly placed in group 2, over
beryllium; Charles Janet's left-step table also contains this
assignment. The normalized ionization potentials and electron
affinities show better trends with helium in group 2 than in
group 18; helium is expected to be slightly more reactive than
neon (which breaks the general trend of reactivity in the noble
gases, where the heavier ones are more reactive); predicted
helium compounds often lack neon analogues even theoretically,
but sometimes have beryllium analogues; and helium over
beryllium better follows the trend of first-row anomalies in the
table (s >> p > d > f).[147][148][149]
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.[150][n 9] It has been claimed that such arguments are
proof that, "it is a mistake to break the [periodic] system into
sharply delimited blocks".[151] A third common variant shows
the two positions below yttrium as being occupied by the
lanthanides and the actinides.[29]
Chemical and physical arguments have been made in support of
lutetium and lawrencium[152][153] but the majority of authors
seem either unconvinced by them or unaware of
them.[154][155] Most working chemists are not aware there is
any controversy.[155] In December 2015 an IUPAC project was
established to make a recommendation on the matter,
considering only the first two alternatives as
possibilities.[156]Lanthanum and actinium

65. La and Ac below Y
Lanthanum and actinium are commonly depicted as the
remaining group 3 members.[157][n 10] It has been suggested
that this layout originated in the 1940s, with the appearance of
periodic tables relying on the ground-state electron
configurations of the elements and the notion of the
differentiating electron. The ground-state configurations
of caesium, barium and lanthanum are [Xe]6s1, [Xe]6s2 and
[Xe]5d16s2. Lanthanum thus emerges with a 5d differentiating
electron and on these grounds it was considered to be "in group
3 as the first member of the d-block for period 6".[158] A
superficially 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".[158] 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.[158] 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
gadolinium[159] (though similar logic would also lead to
thorium getting the 6d2 position, having incumbency
over rutherfordium).
In terms of chemical behaviour,[160] and trends going down
group 3 (if Sc-Y-La is chosen) for properties such as melting
point, electronegativity and ionic radius,[161][162] scandium,
yttrium, lanthanum and actinium are similar to their group 1–2
counterparts. In this variant, the number of f electrons in the

66. most common (trivalent) ions of the f-block elements
consistently matches their position in the f-block.[163] 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.[164] However,
outside the lanthanides there does not exist a typical oxidation
state across any period of a block.[36]Lutetium and lawrencium
Lu and Lr below Y
In other tables, lutetium and lawrencium are the remaining
group 3 members.[n 11] 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".[158] Accordingly, lutetium rather than lanthanum was
assigned to group 3 by some chemists in the 1920s and 30s.[n
12] Several physicists in the 1950s and '60s favoured lutetium,
in light of a comparison of several of its physical properties
with those of lanthanum.[158]
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 a valid concern
given other periodic table anomalies—thorium, for example, has
no f-electrons yet is part of the f-block.[165] Karl Gschneidner,
analysing the melting points of the lanthanides in a 1971 article,
reached the conclusion that it was likely that 4f, 5d, 6s, and 6p
electrons were all involved in the bonding of lanthanide metals
except for lutetium.[166] The fact that lanthanum was
demonstrated to be a 4f-band metal (with about 0.17 electrons
per atom)[37] whereas the 4f shell appears to have no influence
on the metallic properties of lutetium, has been used as an
argument to place lutetium in group 3 instead of
lanthanum.[167] Scandium, yttrium, and lutetium show a more
consistent set of electron configurations matching the global
trend on the periodic table: the 5d metals then all add a closed
4f14 shell. (For example, the shift from yttrium [Kr]4d15s2 to

67. lutetium [Xe]4f145d16s2 exactly parallels that from zirconium
[Kr]4d25s2 to hafnium [Xe]4f145d26s2.)[158] The inclusion of
lutetium rather than lanthanum also homogenises the 5d
transition series: trends in atomic size, coordination number,
and relative abundance of metal–oxygen bonds all reveal that
lutetium is closer than lanthanum to the behaviour of the
uncontroversial 5d metals hafnium through mercury.[168]
As for lawrencium, its gas phase ground-state 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.[169] However, it is expected that in
the condensed phase and in chemical environments lawrencium
has the expected 6d occupancy, and simple modelling studies
suggest it will behave like a lanthanide,[170] as do the rest of
the late actinides, in particular being a homologue of
lutetium.[164]
While scandium, yttrium and lutetium (and presumably
lawrencium) do often behave like trivalent versions of the group
1–2 metals, being hard class-A cations mostly restricted to the
group oxidation state, they are not the only elements in the d-
block or f-block that do so. The early transition metals
zirconium and hafnium in group 4 also display such behaviour,
as does the actinide thorium.[171][172] The physical properties
of the group 3 elements are affected by the presence of a d
electron, which forms more localised bonds within the metals
than the p electrons in the similar group 13 metals;[61] exactly
the same situation is found comparing group 4 to group
14.[60] Trends going down group 3 (if Sc-Y-Lu is chosen) for
properties such as melting point, electronegativity and ionic
radius, are similar to those found among their group 4–8
counterparts.[158] 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

68. the first five f-block elements are La 0, Ce 1, Pr 3, Nd 4 and Pm
5.[158]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, which is stated in the 2005 Red Book to be the
IUPAC-agreed version as of 2005 (a number of later versions
exist, and the last update is from 1 December 2018),[173][n
13] 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 14]
Groups included in the transition metals
The definition of a transition metal, as given by IUPAC in
the Gold Book, is an element whose atom has an incomplete d
sub-shell, or which can give rise to cations with an incomplete d
sub-shell.[175] 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. However, the 2005 IUPAC
nomenclature as codified in the Red Book gives both the group
3–11 and group 3–12 definitions of the transition metals as
alternatives.
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 given up for chemical
bonding (they can sometimes contribute to the valence bonding
orbitals even so, as in zinc fluoride).[176] The 2007 report
of mercury(IV) fluoride (HgF4), a compound in which mercury
would use its d electrons for bonding, has prompted some

69. commentators to suggest that mercury can be regarded as a
transition metal.[177] Other commentators, such as
Jensen,[178] 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.[178]
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 properties
characteristic of transition metal chemistry.[179] In this case,
only groups 4–11 are regarded as transition metals. This
categorisation is however not one of the alternatives considered
by IUPAC. Though the group 3 elements show few of the
characteristic chemical properties of the transition metals, the
same is true of the heavy members of groups 4 and 5, which
also are mostly restricted to the group oxidation state in their
chemistry. Moreover, the group 3 elements show characteristic
physical properties of transition metals (on account of the
presence in each atom of a single d electron).[61]
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.[180] The current experimental evidence still leaves open
the question of whether flerovium behaves more like a metal or
a noble gas.[181]

70. Further periodic table extensions
Main article: Extended periodic table
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Sodium
Magnesium
Aluminium
Silicon
Phosphorus
Sulfur
Chlorine
Argon
Potassium
Calcium
Scandium
Titanium
Vanadium
Chromium
Manganese
Iron
Cobalt
Nickel
Copper

71. Zinc
Gallium
Germanium
Arsenic
Selenium
Bromine
Krypton
Rubidium
Strontium
Yttrium
Zirconium
Niobium
Molybdenum
Technetium
Ruthenium
Rhodium
Palladium
Silver
Cadmium
Indium
Tin
Antimony
Tellurium
Iodine
Xenon
Caesium
Barium
Lanthanum
Cerium
Praseodymium
Neodymium
Promethium
Samarium

72. Europium
Gadolinium
Terbium
Dysprosium
Holmium
Erbium
Thulium
Ytterbium
Lutetium
Hafnium
Tantalum
Tungsten
Rhenium
Osmium
Iridium
Platinum
Gold
Mercury (element)
Thallium
Lead
Bismuth
Polonium
Astatine
Radon
Francium
Radium
Actinium
Thorium
Protactinium
Uranium
Neptunium
Plutonium
Americium
Curium
Berkelium

73. Californium
Einsteinium
Fermium
Mendelevium
Nobelium
Lawrencium
Rutherfordium
Dubnium
Seaborgium
Bohrium
Hassium
Meitnerium
Darmstadtium
Roentgenium
Copernicium
Nihonium
Flerovium
Moscovium
Livermorium
Tennessine
Oganesson
Ununennium
Unbinilium
Unbiunium
Unquadquadium
Unquadpentium
Unquadhexium
Unquadseptium
Unquadoctium
Unquadennium
Unpentnilium
Unpentunium
Unpentbium
Unpenttrium

74. Unpentquadium
Unpentpentium
Unpenthexium
Unpentseptium
Unpentoctium
Unpentennium
Unhexnilium
Unhexunium
Unhexbium
Unhextrium
Unhexquadium
Unhexpentium
Unhexhexium
Unhexseptium
Unhexoctium
Unhexennium
Unseptnilium
Unseptunium
Unseptbium

75. Unbibium
Unbitrium
Unbiquadium
Unbipentium
Unbihexium
Unbiseptium
Unbioctium
Unbiennium
Untrinilium
Untriunium
Untribium
Untritrium
Untriquadium
Untripentium
Untrihexium
Untriseptium
Untrioctium
Untriennium

76. Unquadnilium
Unquadunium
Unquadbium
Unquadtrium
Periodic table with eight rows, extended to element 172[182]
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.[183] More
recently, physicists such as Pekka Pyykkö have theorized that
these additional elements do not exactly 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 123.[184][185]
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 circa 256

77. (which would equate to between elements 99 and 100 in
modern-day terms) did not exist.[186] A higher, more recent
estimate is that the periodic table may end soon after the island
of stability,[187] whose centre is predicted to lie
between element 110 and element 126, as the extension of the
periodic and nuclide tables is restricted by proton and
neutron drip lines as well as decreasing stability
towards spontaneous fission.[188][189] Other predictions of an
end to the periodic table include at element 128 by John
Emsley,[3] at element 137 by Richard Feynman,[190] at element
146 by Yogendra Gambhir,[191] and at element 155 by Albert
Khazan.[3][n 15]Bohr model
The 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.[192] Hence the non-relativistic
Bohr model is inaccurate when applied to such an
element.Relativistic Dirac equation
The relativisticDirac 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.[193] 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.[194] This
does not happen if the innermost orbital is filled, so that
element 173 is not necessarily the end of the periodic
table.[190]
Optimal form
The many different forms of periodic table have prompted the
question of whether there is an optimal or definitive form of

78. periodic table.[195] 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
16]Other
In celebration of the periodic table's 150th anniversary,
the United Nations declared the year 2019 as the International
Year of the Periodic Table, celebrating "one of the most
significant achievements in science".[198]See also
· Chemistry portal
· Abundance of the chemical elements
· Atomic electron configuration table
· Element collecting
· List of chemical elements
· List of periodic table-related articles
· Names for sets of chemical elements
· Standard model
· Table of nuclides
· Template:Spectral lines of the elements
· The Mystery of Matter: Search for the Elements (PBS film)
· Timeline of chemical element discoveriesNotes
1. ^ The elements discovered initially by synthesis and later in
nature are technetium (Z = 43), promethium (61), astatine (85),
neptunium (93), and plutonium (94).
2. ^ An element zero (i.e. a substance composed purely of
neutrons), is included in a few alternate presentations, for

79. example, in the Chemical Galaxy.[6]
3. ^ The noble gases, astatine, francium, and all elements
heavier than americiumwere left out as there is no data for
them.
4. ^ 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.
5. ^ While Lr is thought to have a p rather than d electron in its
ground-state electron configuration, and would therefore be
expected to be a volatile metal capable of forming a +1 cation
in solution like thallium, no evidence of either of these
properties has been able to be obtained despite experimental
attempts to do so.[64] It was originally expected to have a d
electron in its electron configuration[64] and this may still be
the case for metallic lawrencium, whereas gas phase atomic
lawrencium is very likely thought to have a p electron.[65]
6. ^ 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): 684–
688 (687).
7. ^ 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. pp. 315–316, ISBN0-444-

80. 40776-6.
8. ^ See The Internet database of periodic tables for depictions
of these kinds of variants.
9. ^ 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 (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.
10. ^ For examples of this table see Atkins et al. (2006). Shriver
& Atkins Inorganic Chemistry (4th ed.). Oxford: Oxford

81. University Press • Myers et al. (2004). Holt Chemistry.
Orlando: Holt, Rinehart & Winston • Chang R. (2000). Essential
Chemistry (2nd ed.). Boston: McGraw-Hill
11. ^ 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: McGraw-Hill Kogakusha
12. ^ 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.
(1989). Chemistry with Inorganic Qualitative Analysis (3rd
ed.). SanDiego: Harcourt Brace Jovanovich, pp. 955–956, 958.
13. ^ Notwithstanding, an IUPAC member subsequently wrote
that, "IUPAC has not approved any specific form of the periodic
table, and an IUPAC-approved form does not exist, though even
members of IUPAC themselves have published diagrams titled
“IUPAC Periodic Table of the Elements". However, the only
specific recommendation IUPAC has made concerning the
periodic table covers the Group numbering of 1–18."[174]
14. ^ 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 Company
15. ^ Karol (2002, p. 63) contends that gravitational effects
would become significant when atomic numbers become
astronomically large, thereby overcoming other super-massive