Periodic table of elements

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The periodic table of the elements is a tabular method of displaying the chemical elements. A chemical element is a fundamental classification of atomic matter where differentiation of particles is based on the number of protons found in their nuclei, the so-called atomic number. So far, 118 elements are known to exist either by producing them artificially or finding them naturally in the environment.

The component parts of the table are known as Rows and Groups (the later referring to columns). In the table, elements are sorted in ascending order of atomic number. The physical state of matter that they exist in at a temperature of 273.15 K (0 °C) and at a pressure of 100,000 Pa (≈ 1 atm) is usually also provided.

Typical version

The table shown below is only one of many versions. A number of others—often structured in a different manner—are available on the Gallery subpage and elsewhere online. In the upper left hand corner of an entry the atomic mass (formerly known as atomic weight) is given. It is an average over isotopic masses weighted by natural abundance. In the lower left hand corner the atomic number Z is given. In the lower right hand corner the electronegativity of the element is listed.

1.0
H
1 2.1
4.0
He
2 n/a
6.9
Li
3 1.0
9.0
Be
4 1.6
10.8
B
5 2.0
12.0
C
6 2.6
14.0
N
7 3.0
16.0
O
8 3.4
19.0
F
9 4.0
20.2
Ne
10 n/a
23.0
Na
11 0.9
24.3
Mg
12 1.3
27.0
Al
13 1.6
28.1
Si
14 1.9
31.0
P
15 2.2
32.1
S
16 2.6
35.5
Cl
17 3.2
39.9
Ar
18 n/a
39.1
K
19 0.8
40.1
Ca
20 1.0
45.0
Sc
21 1.4
47.9
Ti
22 1.5
50.9
V
23 1.6
52.0
Cr
24 1.6
54.9
Mn
25 1.6
55.8
Fe
26 1.8
58.9
Co
27 1.9
58.7
Ni
28 1.9
63.5
Cu
29 1.9
65.4
Zn
30 1.7
69.7
Ga
31 1.8
72.6
Ge
32 2.0
74.9
As
33 2.2
79.0
Se
34 2.6
79.9
Br
35 3.0
83.8
Kr
36 3.0
85.5
Rb
37 0.8
87.6
Sr
38 1.0
88.9
Y
39 1.2
91.2
Zr
40 1.3
92.9
Nb
41 1.6
95.9
Mo
42 2.2
98.9
Tc
43 1.9
101.1
Ru
44 2.2
102.9
Rh
45 2.3
106.4
Pd
46 2.2
107.9
Ag
47 1.9
112.4
Cd
48 1.7
114.8
In
49 1.8
118.7
Sn
50 2.0
121.8
Sb
51 2.1
127.6
Te
52 2.1
126.9
I
53 2.7
131.3
Xe
54 2.6
132.9
Cs
55 0.8
137.3
Ba
56 0.9
178.5
Hf
72 1.3
180.9
Ta
73 1.5
183.8
W
74 2.4
186.2
Re
75 1.9
190.2
Os
76 2.2
192.2
Ir
77 2.2
195.1
Pt
78 2.3
197.0
Au
79 2.5
200.6
Hg
80 2.0
204.4
Tl
81 1.6
207.2
Pb
82 2.3
209.0
Bi
83 2.0
209.0
Po
84 2.0
210.0
At
85 2.2
222.0
Rn
86 n/a
223.0
Fr
87 0.7
226.0
Ra
88 0.9
265.0
Rf
104
268.0
Db
105
271.0
Sg
106
272.0
Bh
107
270.0
Hs
108
276.0
Mt
109
281.0
Ds
110
280.0
Rg
111
285.0
Cn
112
 
138.9
La
57 1.1
140.1
Ce
58 1.1
140.9
Pr
59 1.1
144.2
Nd
60 1.1
145.0
Pm
61
150.4
Sm
62 1.2
152.0
Eu
63
157.3
Gd
64 1.2
158.9
Tb
65
162.5
Dy
66 1.2
164.9
Ho
67 1.2
167.3
Er
68 1.2
168.9
Tm
69 1.3
173.0
Yb
70 1.1
175.0
Lu
71 1.3
227.0
Ac
89 1.1
232.0
Th
90 1.3
231.0
Pa
91 1.5
238.0
U
92 1.4
237.0
Np
93 1.4
244.0
Pu
94 1.3
243.0
Am
95 1.3
247.0
Cm
96 1.3
247.0
Bk
97 1.3
251.0
Cf
98 1.3
252.0
Es
99 1.3
257.0
Fm
100 1.3
258.0
Md
101 1.3
259.0
No
102 1.3
262.0
Lr
103

Elements by periodic table group (vertical column)

See Atomic electron configuration for the orbital occupancies of ground state atoms.

Elements in any one group behave in a similar way and show the same overall general properties. The number of electrons in the outer shells of the electron orbitals is the same within a group, only the principal quantum number n describing the orbitals increases. For instance, the right column is occupied by the noble gases. Because they have an outer shell of electrons that is completely filled, they are inert in behavior. (See the article atomic orbital for a more detailed explanation of the building up of the electronic shells.) 

On the extreme left there is a group of metals, called the alkali metals. They are characterized by having a single electron in their outer valence orbital. Going down from Li, Na, K to Fr the reactivity decreases, but these metals react easily – for example, with water, in a very exothermic way. This lowered reactivity is related to the atomic weight. As the mass of the nucleus increases with each neutron and proton, so also does the number of electrons (to balance the electric charge). Because these electrons are all at a greater distance from the nucleus, the energy gained from removing one electron diminishes according to the proportionality of n−2. The one free electron in the outer valence shell is in a so-called ns-orbital. In all these metals there is only one free electron and these ball-shaped orbitals are denoted by 1s, 2s, 3s, etc., all with an increasing average radius from the nucleus.

For all the subsequent groups, characterizations can be formulated based upon similarities in reactivity—the way they react, how many electrons can be shared, etc. As a result, the likelihood of reactants to be able to react together can be determined from the Periodic Table of Elements.

The way the element at the top of a column reacts and the way it produces new chemicals give a clue to how other elements in that same group (column) will react. For instance carbon and hydrogen (H2) react to form methane. The bonds of the carbon (C) take a tetrahedral or pyramidal shape. Silicon (Si) responds in the same way and silicon chemistry is to a significant extent similar to carbon chemistry.

Elements by periodic table row

Elements in a row show different periodicities – such as electronegativities, increasing atomic mass, increasing number of protons and neutrons. The stability of these elements diminishes however with an increasing number of nucleons (protons and neutrons). These instabilities lead to radioactive decay as is evident from the actinides and lanthanides, the groups of the rare earths. They contain elements such as plutonium, radon and uranium, and, notorious since the poisoning scandal in the UK, polonium.

The ratio of protons to neutrons is important in determining the stability of elements. If the ratio is greater than 1, the probability increases that an element (or an isotope of an element) will be radioactive and unstable. The greater the ratio, the greater the probability that the element will be radioactive.

Other properties vary differently from the column-wise view of the periodic table of elements. Going downward, the reactivity decreases. Going along a row from left to right, the reactivity increases, and properties such as electronegativity increase as well. For the latter reason, the tendency to lose an electron to form a covalent bond decreases and turns into the tendency of needing one or more electrons to form a stable covalent bond, producing a more stable chemical substance or molecule.

The alkali metals are the most eager to lose an electron (they have the lowest electronegativity in the row) making them very reactive with water. The halogens are reactive for the opposite reason – they need to attract that electron to obtain a more stable electron-configuration. hydrofluoric acid is among the strongest acids known to mankind, where sodium hydroxide is one of the strongest bases — as an example. Electronegativity also is the source for the occurrence of electrical polarity in substances. In every molecule of water (H2O), the hydrogen is less inclined to keep its electron tightly bound. The oxygen, on the other hand, being more electronegative, likes to share electrons to create a more stable electron configuration. This results in the protons'(hydrogen) side of the molecule being slightly more positively charged ∆+, whereas the oxygen side is slightly more negatively charged 2∆−. Therefore water is a highly polar fluid with the capability of facilitating hydrogen bonding in solutions, which is a somewhat non-standard behaviour.

Elements classified alphabetically

See Chemical elements

History

In the early part of the 19th century it was known that matter was composed of a limited number of types of atoms or elementary substances ("elements")[1] [2] and it was observed that many of those elements resembled one another. So it was only natural for chemists to search for a method of classification of the elements. Mass being a property common to all matter, and the relative masses of the elements being known, it was natural to seek some relationship between mass and properties. Early attempts at classification on this basis were hampered by the fact that a number of elements had not been discovered, and by a number of gross inaccuracies in the established atomic weights. In spite of these handicaps, the earlier attempts at classification were so obviously based on some fundamental truth that the modern chemist can only wonder at the negligible attention they attracted, or the ridicule with which they were received.

As early as 1829[3] Johann Wolfgang Döbereiner (1780–1849) pointed out that there were "triads" of similar elements in which the atomic weight of one was the arithmetic mean of the other two, e.g., calcium (Ca), strontium (Sr), and barium (Ba); lithium (Li),sodium (Na), and potassium (K); chlorine (Cl), bromine (Br), and iodine (I). In 1850 Max Josef von Pettenkofer (1818–1901) developed this idea further by drawing attention to series of similar elements in which the atomic weight of any element differed from that of the next lighter or next heavier in the series by 8 or a multiple of 8, e.g., Li, Na, K; or magnesium (Mg), Sr, Ba; oxygen,sulfur,selenium, tellurium. Until then little notice was taken of such regularities.

In 1860 a chemical conference (the first ever) was held in Karlsruhe (Germany) where Stanislao Cannizzaro had presented a more accurate list of atomic weights than had previously been available. Not only had some values been slightly wrong through inaccurate measurements but some were half or a third of the correct value through false reasoning. The availability of these corrected weights inspired the introduction of the first "periodic table" of the now familiar type. It was drawn up by Alexandre-Emile Béguyer de Chancourtois in 1863, who arranged the then known elements in sixteen columns in order of atomic weight, and found that by doing so many similar elements were in the same column (group). Independently of Chancourtois, John A.R. Newlands was thinking along similar lines. He also arranged the known elements in order of atomic weight, and then noticed that in many cases there was a repetition of chemical properties at each 8th element; he called this the "Law of Octaves". When he described his ideas in a lecture at a Chemical Society meeting[4] the lack of spaces for undiscovered elements and the placing of two elements in one box were criticized, and Professor G. F. Foster humorously inquired if he had examined the elements according to their initial letter. Apparently the idea of a periodic table was in the air, for not only Newlands drew up a table, but simultaneously William Odling, also in the UK, designed one. Both tables show already some of the now familiar characteristics of the periodic table.

While Newlands and Odling were developing their table in England, Lothar Meyer and Dmitri Mendeleev in Germany and Russia, respectively, were thinking on similar lines, and independently in 1868–1870 they published tables now known as the Mendeleev table. Mendeleev published his table in Russian, but an abstract appeared in German[5]. In this work Mendeleev placed groups of similar elements in the same horizontal line and not, as later became the custom, in the same column. Meyer's table is known to have been in existence in manuscript form at least as early as 1868, but it was not published until 1870.

In publishing his table Mendeleev recognized that there must be a number of undiscovered elements, and he predicted the properties of some of these. The spectacular confirmation of these predictions by the discovery of gallium (1875) and scandium (1879) had excellent publicity value, and in 1882 the Royal Society of London awarded the Davy Medal jointly to Mendeleev and Lothar Meyer in recognition of their work, ignoring the claims of Newlands. At a later date (1887) they made some reparation to Newlands by awarding him the Davy Medal.

The Mendeleev classification proved to be extremely useful, but it contained some anomalies, that were resolved much later (by Bohr in 1922[6]) when it was realized that the classification should be based on atomic number and not on atomic weight; it turned out that some heavier isotopes were intruding in the table at the wrong places.

References

  1. Nash LK. (1956) The Origin of Dalton's Atomic Theory. Isis 47:101-116. See pp. 108-116 for the author's version of how John Dalton originated the chemical atomic theory.
  2. Rocke AJ. (2005) In Search of El Dorado: John Dalton and the Origins of the Atomic Theory. Social Research 72(1):125-158 (Spring Issue), 34p. | Online Full-Text from BNET
  3. J. W. Döbereiner, Versuch zu einer Gruppirung der Elementaren Stoffe nach ihrer Analogie, Poggendorf's Annalen der Physik und Chemie, vol. 15, pp. 301–307 (1829) Google books (Go to p. 301)
  4. Scanned page of Proceedings of Newlands' lecture
  5. D. Mendelejeff, Ueber die Beziehungen der Eigenschaften zu den Atomgewichten der Elemente, Zeitschrift für Chemie vol. 12, 405-406 (1869)
  6. N. Bohr, Der Bau der Atome und die physikalischen und chemischen Eigenschaften der Elemente, Zeitschrift für Physik, vol. 9, 1 (1922)