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The Periodic Table
The first periodic table was devised by Dmitri Mendeleev and published in 1869.
Mendeleev found he could arrange the 65 elements that were then known in a grid
or table so that each element had:
1. A higher atomic weight than the one on its left.
2. Similar chemical properties to other elements in the same column.
He realized that the table in front of him lay at the very heart of chemistry. In his
table he noted gaps - spaces where elements should be but none had yet been
discovered.
In 1913, Henry Moseley, who worked with Rutherford, showed that it is atomic
number (electric charge) which is most fundamental to the chemical properties of
any element. Mendeleev had believed chemical properties were determined by
atomic weight. Moseley correctly predicted the existence of new elements based on
atomic numbers.
Today the chemical elements are still arranged in order of increasing atomic
number (Z) as you go from left to right across the table. We call the horizontal rows
periods and the vertical rows groups.
We also know now that an element's chemistry is determined by the way its
electrons are arranged - its electron configuration.
The noble gases are found in group 18, on the far right of each period. The
reluctance of the noble gases to undergo chemical reactions indicates that the atoms
of these gases strongly prefer their own electron configurations - featuring a full
outer shell of electrons - to any other.
In contrast to the noble gases, the elements with the highest reactivity are those
with the greatest need to gain or lose electrons in order to achieve a full outer shell
of electrons.
Elements that sit in the same group (e.g. the alkali metals in Group 1) all have the
same number of outer electrons, leading to similar chemical properties.
Likewise the halogens in Group 17 also have similar properties to one another.
When halogens react, they gain an electron to form negatively charged ions. Each
ion has the same electron configuration as the noble gas in the same period. The
ions are therefore more chemically stable than the elements from which they
formed.
There is a progression from metals to non-metals across each period.
The block of elements in groups 3 - 12 contains the transition metals. These are
similar to one another in many ways: they produce colored compounds, have
variable valency and are often used as catalysts.
Then we come to the lanthanides (elements 58 - 71) and actinides (elements 90 -
103). The lanthanides are often called the rare earth elements, although in fact these
elements are not rare. The actinides include most of the well-known elements that
take part in or are produced by nuclear reactions. No element with atomic number
higher than 92 occurs naturally in large quantities. Tiny amounts of plutonium and
neptunium exist in nature as decay products of uranium. These elements, and higher
elements, are also produced artificially in nuclear reactors or particle accelerators
The periodic table of the chemical elements (also known as the periodic
table or periodic table of the elements) is a tabular display of the 118
known chemical elements organized by selected properties of their atomic
structures. Elements are presented by increasing atomic number, the number
of protons in an atom's atomic nucleus. While rectangular in general outline, gaps
are included in the horizontal rows (known as periods) as needed to keep elements
with similar properties together in vertical columns (known as groups), e.g. alkali
metals, alkali earths, halogens, noble gases. The following is the periodic table as
defined by the International Union of Pure and Applied Chemistry (IUPAC):
Group # 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Period
1
1
H
2
He
2
3
Li
4
Be
5
B
6
C
7
N
8
O
9
F
10
Ne
3
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
4
19
K
20
Ca
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
5
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
6
55
Cs
56
Ba
*
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
7
87
Fr
88
Ra
**
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Cn
113
Uut
114
Uuq
115
Uup
116
Uuh
117
Uus
118
Uuo
* Lanthanides
(Lanthanoids)
57
La
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
** Actinides
(Actinoids)
89
Ac
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
Organizing principles
The main value of the periodic table is the ability to predict the chemical properties
of an element based on its location on the table. It should be noted that the
properties vary differently when moving vertically along the columns of the table
than when moving horizontally along the rows.
The layout of the periodic table demonstrates recurring ("periodic") chemical
properties. Elements are listed in order of increasing atomic number (i.e., the
number of protons in the atomic nucleus). Rows are arranged so that elements with
similar properties fall into the same columns (groups or families). According
to quantum mechanical theories of electron configuration within atoms, each row
(period) in the table corresponded to the filling of a quantum shell of electrons.
There are progressively longer periods further down the table, grouping the
elements into s-, p-, d- and f-blocks to reflect their electron configuration
Atomic number
By definition, each chemical element has a unique atomic number, the number
of protons in its nucleus. Different atoms of many elements have different numbers
of neutrons, which differentiates between isotopes of an element. For example, all
atoms of hydrogen have one proton, and no atoms of any other element have
exactly one proton. On the other hand, a hydrogen atom can have one or two
neutrons in its nucleus, or none at all, yet all of these cases are isotopes of
hydrogen, not instances of some other element. (A hydrogen atom with no neutrons
in addition to its sole proton is called protium, one with one neutron in addition to
its proton is called deuterium, and one with two additional neutrons, tritium.)
In the modern periodic table, the elements are placed progressively in each row
(period) from left to right in the sequence of their atomic numbers, with each new
row starting with the next atomic number following the last number in the previous
row. No gaps or duplications exist. Since the elements can be uniquely sequenced
by atomic number, conventionally from lowest to highest, sets of elements are
sometimes specified by such notation as "through", "beyond", or "from ... through",
as in "through iron", "beyond uranium", or "from lanthanum through lutetium". The
terms "light" and "heavy" are sometimes also used informally to indicate relative
atomic numbers (not densities), as in "lighter than carbon" or "heavier than lead",
although technically the weight or mass of atoms of an element (their atomic
weights or atomic masses) do not always increase monotonically with their atomic
numbers.
The significance of atomic numbers to the organization of the periodic table was
not appreciated until the existence and properties of protons and neutrons became
understood. Mendeleev's periodic tables instead used atomic weights, information
determinable to fair precision in his time, which worked well enough in most cases
to give a powerfully predictive presentation far better than any other comprehensive
portrayal of the chemical elements' properties then possible. Substitution of atomic
numbers, once understood, gave a definitive, integer-based sequence for the
elements, still used today even as new synthetic elements are being produced and
studied.
Periodicity of chemical properties
The primary determinant of an element's chemical properties is its electron
configuration, particularly the valence shell electrons. For instance, any atoms with
four valence electrons occupying p orbitals will exhibit some similarity. The type of
orbital in which the atom's outermost electrons reside determines the "block" to
which it belongs. The number of valence shell electrons determines the family, or
group, to which the element belongs.
Subshell S G F D P
Period
1 1s
2 2s
2p
3 3s
3p
4 4s
3d 4p
5 5s
4d 5p
6 6s
4f 5d 6p
7 7s
5f 6d 7p
8 8s 5g 6f 7d 8p
The total number of electron shells an atom has determines the period to which it
belongs. Each shell is divided into different subshells, which as atomic number
increases are filled in roughly this order (the Aufbau principle) (see table).[5]
Hence
the structure of the periodic table. Since the outermost electrons determine chemical
properties, those with the same number of valence electrons are generally grouped
together.
Progressing through a group from lightest element to heaviest element, the outer-
shell electrons (those most readily accessible for participation in chemical
reactions) are all in the same type of orbital, with a similar shape, but with
increasingly higher energy and average distance from the nucleus. For instance, the
outer-shell (or "valence") electrons of the first group, headed by hydrogen, all have
one electron in an s orbital. In hydrogen, that s orbital is in the lowest possible
energy state of any atom, the first-shell orbital (and represented by hydrogen's
position in the first period of the table). In francium, the heaviest element of the
group, the outer-shell electron is in the seventh-shell orbital, significantly further
out on average from the nucleus than those electrons filling all the shells below it in
energy. As another example, both carbon and lead have four electrons in their outer
shell orbitals.
Note that as atomic number (i.e., charge on the atomic nucleus) increases, this leads
to greater spin-orbit coupling between the nucleus and the electrons, reducing the
validity of the quantum mechanical orbital approximation model, which considers
each atomic orbital as a separate entity.[citation needed]
Groups
A group or family is a vertical column in the periodic table. Groups are considered
the most important method of classifying the elements. In some groups, the
elements have very similar properties and exhibit a clear trend in properties down
the group. Under the international naming system, the groups are numbered
numerically 1 through 18 from the left most column (the alkali metals) to the right
most column (the noble gases).[7]
The older naming systems differed slightly
between Europe and America (the table shown in this section shows the old
American Naming System).[8]
Some of these groups have been given trivial (unsystematic) names, such as
the alkali metals, alkaline earth metals, halogens, pnictogens, chalcogens, and noble
gases. However, some other groups, such as group 7, have no trivial names and are
referred to simply by their group numbers, since they display fewer similarities
and/or vertical trends.[7]
Modern quantum mechanical theories of atomic structure explain group trends by
proposing that elements within the same group generally have the same electron
configurations in their valence shell, which is the most important factor in
accounting for their similar properties.[1]
Elements in the same group show patterns in atomic radius, ionization energy,
and electronegativity. From top to bottom in a group, the atomic radii of the
elements increase. Since there are more filled energy levels, valence electrons are
found farther from the nucleus. From the top, each successive element has a lower
ionization energy because it is easier to remove an electron since the atoms are less
tightly bound. Similarly, a group has a top to bottom decrease in electronegativity
due to an increasing distance between valence electrons and the nucleus.[9]
Periods
A period is a horizontal row in the periodic table. Although groups are the most
common way of classifying elements, there are some regions of the periodic table
where the horizontal trends and similarities in properties are more significant than
vertical group trends. This can be true in the d-block (or "transition metals"), and
especially for the f-block, where the lanthanides and actinides form two substantial
horizontal series of elements.
Elements in the same period show trends in atomic radius, ionization
energy, electron affinity, and electronegativity. Moving left to right across a period,
atomic radius usually decreases. This occurs because each successive element has
an added proton and electron which causes the electron to be drawn closer to the
nucleus.[10]
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.[9]
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.[11]
Blocks
Because of the importance of the outermost electron shell, the different regions of
the periodic table are sometimes referred to asperiodic table blocks, named
according to the subshell in which the "last" electron resides. 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 through 18 in IUPAC (3A through 8A in American) and contains, among
others, all of the semimetals. The d-block comprises groups 3 through 12 in IUPAC
(or 3B through 8B in American group numbering) and contains all of the transition
metals. The f-block, usually offset below the rest of the periodic table, comprises
the lanthanides and actinides.[12]
Effective nuclear charge
Effective Nuclear Charge Diagram
The effective nuclear charge is the net positive charge experienced by
an electron in a multi-electron atom. The term "effective" is used because
the shielding effect of negatively charged electrons prevents higher orbital electrons
from experiencing the fullnuclear charge by the repelling effect of inner-layer
electrons. The effective nuclear charge experienced by the outer shell electron is
also called the core charge. It is possible to determine the strength of the nuclear
charge by looking at the oxidation number of the atom.
Calculating the effective nuclear charge
In an atom with one electron, that electron experiences the full charge of the
positive nucleus. In this case, the effective nuclear charge can be calculated
from Coulomb's law.
However, in an atom with many electrons the outer electrons are simultaneously
attracted to the positive nucleus and repelled by the negatively charged electrons.
The effective nuclear charge on such an electron is given by the following equation:
Zeff = Z − S
where
Z is the number of protons in the nucleus (atomic number), and
S is the average number of electrons between the nucleus and the electron in
question (the number of nonvalence electrons).
S can be found by the systematic application of various rule sets, the
simplest of which is known as "Slater's rules" (named after John C.
Slater). Douglas Hartree defined the effective Z of a Hartree-Fock orbital
to be:
where
<r>H is the mean radius of the orbital for hydrogen, and
<r>Z is the mean radius of the orbital for an electron configuration with
nuclear charge Z.
Note: Zeff is also often written Z*.
Example
Consider a sodium cation, a fluorine anion, and a
neutral neon atom. Each has 10 electrons, and the number of
nonvalence electrons is 2 (10 total electrons - 8 valence) but the
effective nuclear charge varies because each has a different
atomic number:
Zeff(F-) = 9 − 2 = 7 +
Zeff(Ne) = 10 − 2 = 8 +
Zeff(Na+) = 11 − 2 = 9 +
So, the sodium cation has the largest effective nuclear
charge, and thus the smallest atomic radius.
Values Shielding effect
The shielding effect describes the decrease in attraction between an electron and
the nucleus in any atom with more than one electron shell. It is also referred to as
the screening effect or atomic shielding.
Slater's rules
In quantum chemistry, Slater's rules provide numerical values for the effective
nuclear charge concept. In a many-electron atom, each electron is said to
experience less than the actual nuclear chargeowning to shielding or screening by
the other electrons. For each electron in an atom, Slater's rules provide a value for
the screening constant, denoted by s, S, or σ, which relates the effective and actual
nuclear charges as
The rules were devised semi-empirically by John C. Slater and published in
1930.[1]
Rules
Firstly,[1][4]
the electrons are arranged in to a sequence of groups in order of
increasing principal quantum number n, and for equal n in order of
increasing azimuthal quantum number l, except that s- and p- orbitals are kept
together.
[1s] [2s,2p] [3s,3p] [3d] [4s,4p] [4d] [4f] [5s, 5p] [5d] etc.
Each group is given a different shielding constant which depends upon the
number and types of electrons in those groups preceding it.
The shielding constant for each group is formed as the sum of the following
contributions:
1. An amount of 0.35 from each other electron within the same group
except for the [1s] group, where the other electron contributes only
0.30.
2. If the group is of the [s p] type, an amount of 0.85 from each electron
with principal quantum number (n) one less and an amount of 1.00 for
each electron with an even smaller principal quantum number
3. If the group is of the [d] or [f], type, an amount of 1.00 for each
electron inside it. This includes i) electrons with a smaller principal
quantum number and ii) electrons with an equal principal quantum
number and a smaller azimuthal quantum number (l)
In tabular form, the rules are summarized as:
Group
Other
electrons in
the same
group
Electrons in group(s)
with principal quantum
number n
andazimuthal quantum
number < l
Electrons in
group(s)
with principal
quantum
number n-1
Electrons in all
group(s)
with principal
quantum number <
n-1
[1s] 0.3 N/A N/A N/A
[ns,np] 0.35 N/A 0.85 1
[nd] or [nf] 0.35 1 1 1
Atomic Radius
The atomic radius of an element is half of the distance between the centers of two
atoms of that element that are just touching each other. Generally, the atomic radius
decreases across a period from left to right and increases down a given group. The
atoms with the largest atomic radii are located in Group I and at the bottom of
groups.
Moving from left to right across a period, electrons are added one at a time to the
outer energy shell. Electrons within a shell cannot shield each other from the
attraction to protons. Since the number of protons is also increasing, the effective
nuclear charge increases across a period. This causes the atomic radius to decrease.
Moving down a group in the periodic table, the number of electrons and filled
electron shells increases, but the number of valence electrons remains the same.
The outermost electrons in a group are exposed to the same effective nuclear
charge, but electrons are found farther from the nucleus as the number of filled
energy shells increases. Therefore, the atomic radii increase.
Ionization Energy
The ionization energy, or ionization potential, is the energy required to completely
remove an electron from a gaseous atom or ion. The closer and more tightly bound
an electron is to the nucleus, the more difficult it will be to remove, and the higher
its ionization energy will be. The first ionization energy is the energy required to
remove one electron from the parent atom. The second ionization energy is the
energy required to remove a second valence electron from the univalent ion to form
the divalent ion, and so on. Successive ionization energies increase. The second
ionization energy is always greater than the first ionization energy. Ionization
energies increase moving from left to right across a period (decreasing atomic
radius). Ionization energy decreases moving down a group (increasing atomic
radius). Group I elements have low ionization energies because the loss of an
electron forms a stable octet.
The property is alternately still often called the ionization potential, measured in
volts. In chemistry it is often referred to one mole of substance (molar ionization
energy or enthalpy) and reported in kJ/mol. In atomic physics the ionization energy
is typically measured in the unit electron volt (eV).
The ionization energy is different for electrons of different atomic or molecular
orbitals. More generally, the nth ionization energy is the energy required to strip off
the nth electron after the first n − 1 electrons have been removed. It is considered a
measure of the tendency of an atom or ion to surrender an electron, or the strength
of the electron binding; the greater the ionization energy, the more difficult it is to
remove an electron. The ionization energy may be an indicator of the reactivity of
an element. Elements with a low ionization energy tend to be reducing agents and
form cations, which in turn combine with anions to form salts.
Electron binding energy (BE), more accurately, is the energy required to release
an electron from its atomic or molecular orbital when adsorbed to a surface rather
than a free atom. Binding energy values are normally reported as positive values
with units of eV. The binding energies of 1s electrons are roughly proportional to
(Z-1)² (Moseley's law).
Values and trends
Main article: Molar ionization energies of the elements
Generally the (n+1)th ionization energy is larger than the nth ionization energy.
Always, the next ionization energy involves removing an electron from an
orbital closer to the nucleus. Electrons in the closer orbitals experience greater
forces of electrostatic attraction; thus, their removal requires increasingly more
energy.
Some values for elements of the third period are given in the following table:
Successive molar ionization energies in kJ/mol
(96.485 kJ/mol = 1 eV/particle)
Element First Second Third Fourth Fifth Sixth Seventh
Na 496 4,560
Mg 738 1,450 7,730
Al 577 1,816 2,881 11,600
Si 786 1,577 3,228 4,354 16,100
P 1,060 1,890 2,905 4,950 6,270 21,200
S 999.6 2,260 3,375 4,565 6,950 8,490 27,107
Cl 1,256 2,295 3,850 5,160 6,560 9,360 11,000
Ar 1,520 2,665 3,945 5,770 7,230 8,780 12,000
Large jumps in the successive molar ionization energies occur when
passing noble gas configurations. For example, as can be seen in the table above,
the first two molar ionization energies of magnesium (stripping the two 3s
electrons from a magnesium atom) are much smaller than the third, which
requires stripping off a 2p electron from the very stable neon configuration of
Mg2+
.
Periodic trend for ionization energy. Each period begins at a minimum for the
alkali metals, and ends at a maximum for the noble gases.
Ionization energy is also a periodic trend within the periodic table organization.
Moving left to right within aperiod or upward within a group, the first ionization
energy generally increases. As the atomic radiusdecreases, it becomes harder to
remove an electron that is closer to a more positively charged nucleus. ionization
enthalpy increases from left to right in a period and decreases from top to bottom
in a group.
Electron Affinity
Electron affinity reflects the ability of an atom to accept an electron. It is the energy
change that occurs when an electron is added to a gaseous atom. Atoms with
stronger effective nuclear charge have greater electron affinity. Some
generalizations can be made about the electron affinities of certain groups in the
periodic table. The Group IIA elements, the alkaline earths, have low electron
affinity values. These elements are relatively stable because they have
filled s subshells. Group VIIA elements, the halogens, have high electron affinities
because the addition of an electron to an atom results in a completely filled shell.
Group VIII elements, noble gases, have electron affinities near zero, since each
atom possesses a stable octet and will not accept an electron readily. Elements of
other groups have low electron affinities.
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