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Periodic table (47994 views - Periodic Table Of Elements)
The periodic table is a tabular arrangement of the chemical elements, ordered by their atomic number, electron configurations, and recurring chemical properties. This ordering shows periodic trends, such as elements with similar behaviour in the same column. It also shows four rectangular blocks with some approximately similar chemical properties. In general, within one row (period) the elements are metals on the left, and non-metals on the right. The rows of the table are called periods; the columns are called groups. Six groups have generally accepted names as well as numbers: for example, group 17 elements are the halogens; and group 18, the noble gases. The periodic table can be used to derive relationships between the properties of the elements, and predict the properties of new elements yet to be discovered or synthesized. The periodic table provides a useful framework for analyzing chemical behaviour, and is widely used in chemistry and other sciences. The Russian chemist Dmitri Mendeleev published the first widely recognized periodic table in 1869. He developed his table to illustrate periodic trends in the properties of the then-known elements. Mendeleev also predicted some properties of then-unknown elements that would be expected to fill gaps in this table. Most of his predictions were proved correct when the elements in question were subsequently discovered. Mendeleev's periodic table has since been expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour. All elements from atomic numbers 1 (hydrogen) to 118 (oganesson) have been discovered or synthesized, with the most recent additions (nihonium, moscovium, tennessine, and oganesson) being confirmed by the International Union of Pure and Applied Chemistry (IUPAC) in 2015 and officially named in 2016: they complete the first seven rows of the periodic table. The first 94 elements exist naturally, although some are found only in trace amounts and were synthesized in laboratories before being found in nature. Elements with atomic numbers from 95 to 118 have only been synthesized in laboratories or nuclear reactors. Synthesis of elements having higher atomic numbers is being pursued. Numerous synthetic radionuclides of naturally occurring elements have also been produced in laboratories.
PARTcloud - Periodic table
This article is about the table used in chemistry. For other uses, see Periodic table (disambiguation).
Periodic table forms
18-column
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Chemical Galaxy
Janet's left step
Beyond period 7
Pyykkö
Periodic table history
Dmitri Mendeleev
Discovery of elements
Naming & etymology
for places
(in East Asia)
Systematic element names
Sets of elements
By periodic table structure
1 (alkali metals)
2 (alkaline earth metals)
15 (pnictogens)
16 (chalcogens)
17 (halogens)
18 (noble gases)
Periods (1–7 ...)
Blocks (s, p, d, f, ...)
Atomic orbitals
Aufbau principle
By metallic classification
alkaline earth
post-transition
actinide (super-/eka-super-)
dividing metals & nonmetals
Nonmetals
polyatomic
monatomic (noble gas)
By other characteristics
Coinage metals
Platinum-group metals
Light metals
Native metals
Main-group elements
Rare-earth elements
Transuranium, transplutonium elements
Major, minor & trans- actinides
List of chemical elements ...
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Chemistry Portal
The periodic table is a tabular arrangement of the chemical elements, ordered by their atomic number, electron configurations, and recurring chemical properties. This ordering shows periodic trends, such as elements with similar behaviour in the same column. It also shows four rectangular blocks with some approximately similar chemical properties. In general, within one row (period) the elements are metals on the left, and non-metals on the right.
The rows of the table are called periods; the columns are called groups. Six groups have generally accepted names as well as numbers: for example, group 17 elements are the halogens; and group 18, the noble gases. The periodic table can be used to derive relationships between the properties of the elements, and predict the properties of new elements yet to be discovered or synthesized. The periodic table provides a useful framework for analyzing chemical behaviour, and is widely used in chemistry and other sciences.
The Russian chemist Dmitri Mendeleev published the first widely recognized periodic table in 1869. He developed his table to illustrate periodic trends in the properties of the then-known elements. Mendeleev also predicted some properties of then-unknown elements that would be expected to fill gaps in this table. Most of his predictions were proved correct when the elements in question were subsequently discovered. Mendeleev's periodic table has since been expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour.
All elements from atomic numbers 1 (hydrogen) to 118 (oganesson) have been discovered or synthesized, with the most recent additions (nihonium, moscovium, tennessine, and oganesson) being confirmed by the International Union of Pure and Applied Chemistry (IUPAC) in 2015 and officially named in 2016: they complete the first seven rows of the periodic table.[1][2] The first 94 elements exist naturally, although some are found only in trace amounts and were synthesized in laboratories before being found in nature.[n 1] Elements with atomic numbers from 95 to 118 have only been synthesized in laboratories or nuclear reactors.[3] Synthesis of elements having higher atomic numbers is being pursued. Numerous synthetic radionuclides of naturally occurring elements have also been produced in laboratories.
2 Grouping methods
2.1 Groups
2.2 Periods
2.4 Metals, metalloids and nonmetals
3 Periodic trends and patterns
3.1 Electron configuration
3.2 Atomic radii
3.3 Ionization energy
3.4 Electronegativity
3.5 Electron affinity
3.6 Metallic character
3.7 Linking or bridging groups
4.1 First systemization attempts
4.2 Mendeleev's table
4.3 Second version and further development
5 Different periodic tables
5.1 The long- or 32-column table
5.2 Tables with different structures
6 Open questions and controversies
6.1 Placement of hydrogen and helium
6.2 Group 3 and its elements in periods 6 and 7
6.2.1 Lanthanum and actinium
6.2.2 Lutetium and lawrencium
6.2.3 Lanthanides and actinides
6.3 Groups included in the transition metals
6.4 Elements with unknown chemical properties
6.5 Further periodic table extensions
6.6 Element with the highest possible atomic number
6.7 Optimal form
For large cell versions, see Periodic table (large cells).
Alkali metals Alkaline earth metals Pnictogens Chalcogens Halogens Noble gases
Hydrogen1H Helium2He
Lithium3Li Beryllium4Be Boron5B Carbon6C Nitrogen7N Oxygen8O Fluorine9F Neon10Ne
Sodium11Na Magnesium12Mg Aluminium13Al Silicon14Si Phosphorus15P Sulfur16S Chlorine17Cl Argon18Ar
Potassium19K Calcium20Ca Scandium21Sc Titanium22Ti Vanadium23V Chromium24Cr Manganese25Mn Iron26Fe Cobalt27Co Nickel28Ni Copper29Cu Zinc30Zn Gallium31Ga Germanium32Ge Arsenic33As Selenium34Se Bromine35Br Krypton36Kr
Rubidium37Rb Strontium38Sr Yttrium39Y Zirconium40Zr Niobium41Nb Molybdenum42Mo Technetium43Tc Ruthenium44Ru Rhodium45Rh Palladium46Pd Silver47Ag Cadmium48Cd Indium49In Tin50Sn Antimony51Sb Tellurium52Te Iodine53I Xenon54Xe
Caesium55Cs Barium56Ba Lanthanum57La Hafnium72Hf Tantalum73Ta Tungsten74W Rhenium75Re Osmium76Os Iridium77Ir Platinum78Pt Gold79Au Mercury80Hg Thallium81Tl Lead82Pb Bismuth83Bi Polonium84Po Astatine85At Radon86Rn
Francium87Fr Radium88Ra Actinium89Ac Rutherfordium104Rf Dubnium105Db Seaborgium106Sg Bohrium107Bh Hassium108Hs Meitnerium109Mt Darmstadtium110Ds Roentgenium111Rg Copernicium112Cn Nihonium113Nh Flerovium114Fl Moscovium115Mc Livermorium116Lv Tennessine117Ts Oganesson118Og
Cerium58Ce Praseodymium59Pr Neodymium60Nd Promethium61Pm Samarium62Sm Europium63Eu Gadolinium64Gd Terbium65Tb Dysprosium66Dy Holmium67Ho Erbium68Er Thulium69Tm Ytterbium70Yb Lutetium71Lu
Thorium90Th Protactinium91Pa Uranium92U Neptunium93Np Plutonium94Pu Americium95Am Curium96Cm Berkelium97Bk Californium98Cf Einsteinium99Es Fermium100Fm Mendelevium101Md Nobelium102No Lawrencium103Lr
black=solid green=liquid red=gas gray=unknown Color of the atomic number shows state of matter (at 0 °C and 1 atm)
Primordial From decay Synthetic Border shows natural occurrence of the element
Background color shows subcategory in the metal–metalloid–nonmetal trend:
Metal Metalloid Nonmetal Unknown
Alkali metal Alkaline earth metal Lanthanide Actinide Transition metal Post-transition metal Polyatomic nonmetal Diatomic 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.[4]
In the standard periodic table, the elements are listed in order of increasing atomic number Z (the number of protons in the nucleus of an atom). A new row (period) is started when a new electron shell has its first electron. Columns (groups) are determined by the electron configuration of the atom; elements with the same number of electrons in a particular subshell fall into the same columns (e.g. oxygen and selenium are in the same column because they both have four electrons in the outermost p-subshell). Elements with similar chemical properties generally fall into the same group in the periodic table, although in the f-block, and to some respect in the d-block, the elements in the same period tend to have similar properties, as well. Thus, it is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it.[5]
As of 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 announced by the IUPAC in June 2016 and made official in November 2016.[6][7][8][9]
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).[10]
Grouping methods
Main article: Group (periodic table)
A group or family is a vertical column in the periodic table. Groups usually have more significant periodic trends than periods and blocks, explained below. Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group generally have the same electron configurations in their valence shell.[11] Consequently, elements in the same group tend to have a shared chemistry and exhibit a clear trend in properties with increasing atomic number.[12] However, 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.[13][14][15]
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).[16] 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.[17]
Some of these groups have been given trivial (unsystematic) names, as seen in the table below, although some are rarely used. Groups 3–10 have no trivial names and are referred to simply by their group numbers or by the name of the first member of their group (such as "the scandium group" for group 3), since they display fewer similarities and/or vertical trends.[16]
Elements in the same group tend to show patterns in atomic radius, ionization energy, and electronegativity. From top to bottom in a group, the atomic radii of the elements increase. Since there are more filled energy levels, valence electrons are found farther from the nucleus. From the top, each successive element has a lower ionization energy because it is easier to remove an electron since the atoms are less tightly bound. Similarly, a group has a top to bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus.[18] There are exceptions to these trends, however: for example, in group 11, electronegativity increases farther down the group.[19]
Groups in the periodic table
IUPAC group
Mendeleev (I–VIII) IA IIA IIIB IVB VB VIB VIIB VIIIB IB IIB IIIB IVB VB VIB VIIB b
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 VA VIA VIIA VIIIB IB IIB IIIB IVB VB VIB VIIB 0
Trivial name Alkali metals Alkaline earth metals Coinage metalsc Volatile metalsc Icosagensc Crystallogensc Pnictogens Chalcogens Halogens Noble gases
Name by element Lithium group Beryllium group Scandium group 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 d H d He
Period 2 Li Be B C N O F Ne
Period 3 Na Mg Al Si 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 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 Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
a Group 3 has scandium (Sc) and yttrium (Y). For the rest of the group, sources differ as either being (1) lutetium (Lu) and lawrencium (Lr), or (2) lanthanum (La) and actinium (Ac), or (3) the whole set of 15+15 lanthanides and actinides. IUPAC has initiated a project to standardize the definition as either (1) Sc, Y, Lu and Lr, or (2) Sc, Y, La and Ac.[20]
b Group 18, the noble gases, were not discovered at the time of Mendeleev's original table. Later (1902), Mendeleev accepted the evidence for their existence, and they could be placed in a new "group 0", consistently and without breaking the periodic table principle.
c This group name is not recommended by IUPAC.
d Hydrogen (H), while placed in column 1, is not considered to be in the group alkali metals.
Main article: Period (periodic table)
A period is a horizontal row in the periodic table. Although groups generally have more significant periodic trends, there are regions where horizontal trends are more significant than vertical group trends, such as the f-block, where the lanthanides and actinides form two substantial horizontal series of elements.[21]
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.[22] 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.[18] 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.[23]
Main article: Block (periodic table)
Specific regions of the periodic table can be referred to as blocks in recognition of the sequence in which the electron shells of the elements are filled. Each block is named according to the subshell in which the "last" electron notionally resides.[24][n 3] The s-block comprises the first two groups (alkali metals and alkaline earth metals) as well as hydrogen and helium. The p-block comprises the last six groups, which are groups 13 to 18 in IUPAC group numbering (3A to 8A in American group numbering) and contains, among other elements, all of the metalloids. The d-block comprises groups 3 to 12 (or 3B to 2B in American group numbering) and contains all of the transition metals. The f-block, often offset below the rest of the periodic table, has no group numbers and comprises lanthanides and actinides.[25]
Metals, metalloids and nonmetals
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 the noble gases). The 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.[26]
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 are 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. The nonmetals are simply subdivided into the polyatomic nonmetals, which, being nearest to the metalloids, show some incipient metallic character; the diatomic nonmetals, which are essentially nonmetallic; and the monatomic noble gases, which are nonmetallic and almost completely inert. Specialized groupings such as the refractory metals and the noble metals, which are subsets (in this example) of the transition metals, are also known[27] and occasionally denoted.[28]
Placing the elements into categories and subcategories based on shared properties is imperfect. There is a spectrum of properties within each category and it is not hard to find overlaps at the boundaries, as is the case with most classification schemes.[29] 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 nonmetal and a noble gas yet has some cationic chemistry that is more characteristic of a metal. 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[30] wrote that simple boundary lines could be drawn on the periodic table to show elements having like properties, such as the metals and the nonmetals, or the gaseous elements.
Periodic trends and patterns
Main article: Periodic trends
Main article: Electronic configuration
Approximate order in which shells and subshells are arranged by increasing energy according to the Madelung rule
The electron configuration or organisation of electrons orbiting neutral atoms shows a recurring pattern or periodicity. The electrons occupy a series of electron shells (numbered shell 1, shell 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—two in the s subshell and six in the p subshell. In periodic table terms, the first time an electron occupies a new shell corresponds to the start of each new period, these positions being occupied by hydrogen and the alkali metals.[31][32]
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.[31][32]
Atomic radii
Main article: Atomic radius
Atomic number plotted against atomic radius[n 4]
Atomic radii vary in a predictable and explainable manner across the periodic table. For instance, the radii generally decrease along each period of the table, from the alkali metals to the noble gases; and increase down each group. The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory.[33]
The electrons in the 4f-subshell, which is progressively filled from cerium (element 58) to ytterbium (element 70), are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii that are smaller than would be expected and that are almost identical to the atomic radii of the elements immediately above them.[34] Hence hafnium has virtually the same atomic radius (and chemistry) as zirconium, and tantalum has an atomic radius similar to niobium, and so forth. This is known as the lanthanide contraction. The effect of the lanthanide contraction is noticeable up to platinum (element 78), after which it is masked by a relativistic effect known as the inert pair effect.[35] 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.[34]
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.[35]
Large jumps in the successive molar ionization energies occur when removing an electron from a noble gas (complete electron shell) configuration. For magnesium again, the first two molar ionization energies of magnesium given above correspond to removing the two 3s electrons, and the third ionization energy is a much larger 7730 kJ/mol, for the removal of a 2p electron from the very stable neon-like configuration of Mg2+. Similar jumps occur in the ionization energies of other third-row atoms.[35]
Main article: Electronegativity
Electronegativity is the tendency of an atom to attract electrons.[36] 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.[37] In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, fluorine is the most electronegative of the elements,[n 5] while caesium is the least, at least of those elements for which substantial data is available.[19]
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.[19] The anomalously high electronegativity of lead, particularly when compared to thallium and bismuth, appears to be an artifact of data selection (and data availability)—methods of calculation other than the Pauling method show the normal periodic trends for these elements.[38]
Main article: Electron affinity
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.[41]
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.[41]
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. However, 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.[42]
Metallic character
The lower the values of ionization energy, electronegativity and electron affinity, the more metallic character the element has. Conversely, nonmetallic character increases with higher values of these properties.[43] 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,[44] 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.[45][46]
Linking or bridging groups
32-column periodic table showing, from left to right, the location of group 3; lutetium and lawrencium; groups 11–12; and the noble gases
From left to right across the four blocks of the long- or 32-column form of the periodic table are a series of linking or bridging groups of elements, located approximately between each block.[47] These groups, like the metalloids, show properties in between, or that are a mixture of, groups to either side. Chemically, the group 3 elements, scandium, yttrium, lanthanum and actinium behave largely like the alkaline earth metals[48] or, more generally, s block metals[49][50] but have some of the physical properties of d block transition metals.[51] Lutetium and lawrencium, at the end of the end of the f block, may constitute another linking or bridging group. Lutetium behaves chemically as a lanthanide but shows a mix of lanthanide and transition metal physical properties.[52][53] Lawrencium, as an analogue of lutetium, would presumable display like characteristics.[n 6] The coinage metals in group 11 (copper, silver, and gold) are chemically capable of acting as either transition metals or main group metals.[56] 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.[57][58] 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.[47]
Main article: History of the periodic table
First systemization attempts
In 1789, Antoine Lavoisier published a list of 33 chemical elements, grouping them into gases, metals, nonmetals, and earths.[59] 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;[60] this became known as the Law of Triads.[61] German chemist Leopold Gmelin worked with this system, and by 1843 he had identified ten triads, three groups of four, and one group of five. Jean-Baptiste Dumas published work in 1857 describing relationships between various groups of metals. Although various chemists were able to identify relationships between small groups of elements, they had yet to build one scheme that encompassed them all.[60]
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.[62] This concept eventually became known as valency; different elements bond with different numbers of atoms.[63]
In 1862, Alexandre-Emile Béguyer de Chancourtois, a French geologist, published an early form of periodic table, which he called the telluric helix or screw. He was the first person to notice the periodicity of the elements. With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois showed that elements with similar properties seemed to occur at regular intervals. His chart included some ions and compounds in addition to elements. His paper also used geological rather than chemical terms and did not include a diagram; as a result, it received little attention until the work of Dmitri Mendeleev.[64]
In 1864, Julius Lothar Meyer, a German chemist, published a table with 44 elements arranged by valency. The table showed that elements with similar properties often shared the same valency.[65] Concurrently, William Odling (an English chemist) published an arrangement of 57 elements, ordered on the basis of their atomic weights. With some irregularities and gaps, he noticed what appeared to be a periodicity of atomic weights among the elements and that this accorded with "their usually received groupings".[66] Odling alluded to the idea of a periodic law but did not pursue it.[67] He subsequently proposed (in 1870) a valence-based classification of the elements.[68]
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.[69][70] This so termed Law of Octaves, however, was ridiculed by Newlands' contemporaries, and the Chemical Society refused to publish his work.[71] Newlands was nonetheless able to draft a table of the elements and used it to predict the existence of missing elements, such as germanium.[72] The Chemical Society only acknowledged the significance of his discoveries five years after they credited Mendeleev.[73]
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.[74][75]
Mendeleev's table
Dmitri Mendeleev, watercolour by Ilya Repin
Mendeleev's periodic table from his book An Attempt Towards a Chemical Conception of the Ether
Russian chemistry professor Dmitri Mendeleev and German chemist Julius Lothar Meyer independently published their periodic tables in 1869 and 1870, respectively.[76] Mendeleev's table was his first published version; that of Meyer was an expanded version of his (Meyer's) table of 1864.[77] 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.[78]
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.[79] 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.[80] The second decision was to occasionally ignore the order suggested by the atomic weights and switch adjacent elements, such as tellurium and iodine, to better classify them into chemical families. Later in 1913, Henry Moseley determined experimental values of the nuclear charge or atomic number of each element, and showed that Mendeleev's ordering actually corresponds to the order of increasing atomic number.[81]
The significance of atomic numbers to the organization of the periodic table was not appreciated until the existence and properties of protons and neutrons became understood. Mendeleev's periodic tables used atomic weight instead of atomic number to organize the elements, information determinable to fair precision in his time. Atomic weight worked well enough in most cases to (as noted) give a presentation that was able to predict the properties of missing elements more accurately than any other method then known. Substitution of atomic numbers, once understood, gave a definitive, integer-based sequence for the elements, and Moseley predicted (in 1913) that the only elements still missing between aluminium (Z=13) and gold (Z=79) were Z = 43, 61, 72 and 75, all of which were later discovered. The sequence of atomic numbers is still used today even as new synthetic elements are being produced and studied.[82]
Second version and further development
Mendeleev's 1871 periodic table with eight groups of elements. Dashes represented elements unknown in 1871.
Eight-column form of periodic table, updated with all elements discovered to 2016
In 1871, Mendeleev published his periodic table in a new form, with groups of similar elements arranged in columns rather than in rows, and those columns numbered I to VIII corresponding with the element's oxidation state. He also gave detailed predictions for the properties of elements he had earlier noted were missing, but should exist.[83] These gaps were subsequently filled as chemists discovered additional naturally occurring elements.[84] It is often stated that the last naturally occurring element to be discovered was francium (referred to by Mendeleev as eka-caesium) in 1939.[85] However, plutonium, produced synthetically in 1940, was identified in trace quantities as a naturally occurring element in 1971.[86]
The popular[87] 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.[88][n 7] Merck and Company prepared a handout form of Deming's 18-column medium table, in 1928, which was widely circulated in American schools. By the 1930s Deming's table was appearing in handbooks and encyclopaedias of chemistry. It was also distributed for many years by the Sargent-Welch Scientific Company.[89][90][91]
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.[92]
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.[93][94][n 8]
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 1939.[95] 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.[96] 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.[97]
Different periodic tables
The long- or 32-column table
The modern periodic table is sometimes expanded into its long or 32-column form by reinstating the footnoted f-block elements into their natural position between the s- and d-blocks. Unlike the 18-column form this arrangement results in "no interruptions in the sequence of increasing atomic numbers".[98] The relationship of the f-block to the other blocks of the periodic table also becomes easier to see.[99] 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.[100] 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),[101] and the familiarity of chemists with the modern form (as introduced by Seaborg).[102]
Tables with different structures
Main article: Alternative periodic tables
Within 100 years of the appearance of Mendeleev's table in 1869 it has been estimated that around 700 different periodic table versions were published.[103] As well as numerous rectangular variations, other periodic table formats have been shaped, for example,[n 9] like a circle, cube, cylinder, building, spiral, lemniscate,[104] 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.[103]
Theodor Benfey's spiral periodic table
A popular[105] alternative structure is that of Theodor Benfey (1960). The elements are arranged in a continuous spiral, with hydrogen at the centre and the transition metals, lanthanides, and actinides occupying peninsulas.[106]
Most periodic tables are two-dimensional;[3] however, 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),[107] Wringley's Lamina System (1949),[108] Giguère's Periodic helix (1965)[109] and Dufour's Periodic Tree (1996).[110] Going one further, Stowe's Physicist's Periodic Table (1989)[111] has been described as being four-dimensional (having three spatial dimensions and one colour dimension).[112]
The various forms of periodic tables can be thought of as lying on a chemistry–physics continuum.[113] Towards the chemistry end of the continuum can be found, as an example, Rayner-Canham's "unruly"[114] Inorganic Chemist's Periodic Table (2002),[115] which emphasizes trends and patterns, and unusual chemical relationships and properties. Near the physics end of the continuum is Janet's Left-Step Periodic Table (1928). This has a structure that shows a closer connection to the order of electron-shell filling and, by association, quantum mechanics.[116] A somewhat similar approach has been taken by Alper,[117] albeit criticized by Eric Scerri as disregarding the need to display chemical and physical periodicity.[118] 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.[119] 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.[67][120]
Left-step periodic table (by Charles Janet)
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 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
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 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 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 Nh Fl Mc Lv Ts Og 119 120
f-block d-block p-block s-block
This form of periodic table is more congruent with the order in which electron shells are ideally filled according to Madelung's rule, as shown in the accompanying sequence in the left margin (read from top to bottom, left to right). In reality, the filling of electron shells is characterized by a number of irregularities.
Open questions and controversies
Placement of hydrogen and helium
Simply following electron configurations, hydrogen (electronic configuration 1s1) and helium (1s2) should be placed in groups 1 and 2, above lithium ([He]2s1) and beryllium ([He]2s2).[24] However, such placing is rarely used 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 that group, as it did share the extreme chemical inertness seen throughout the group. As the group changed its formal number, many authors continued to assign helium directly above neon, in group 18; one of the examples of such placing is the current IUPAC table.[121]
Hydrogen's chemical properties are not very close to those of the alkali metals, which occupy group 1, and on that basis hydrogen is sometimes placed elsewhere: one of the most common alternatives is in group 17;[118] one of the factors behind it is the strictly univalent predominantly non-metallic chemistry of hydrogen, and that of fluorine (the element placed on the top of group 17) is strictly univalent and non-metallic. Sometimes, to show how hydrogen has properties both corresponding to those of the alkali metals and the halogens, it may be shown in two columns simultaneously.[122] Another suggestion is above carbon in group 14: placed that way, it fits well into the trend of increasing trends of 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.[123] Finally, hydrogen is sometimes placed separately from any group; this is based on how general properties of hydrogen differ from that of any group. The other period 1 element, helium, is sometimes placed separately from any group as well.[124] The property that distinguishes helium from the rest of the noble gases (even though the extraordinary inertness of helium is extremely close to that of neon and argon[125]) is that in its closed electron shell, helium has only two electrons in the outermost electron orbital, while the rest of the noble gases have eight.
Group 3 and its elements in periods 6 and 7
Although scandium and yttrium are always the first two elements in group 3 the identity of the next two elements is not completely settled. They are commonly lanthanum and actinium, and less often lutetium and lawrencium. The two variants originate from historical difficulties in placing the lanthanides in the periodic table, and arguments as to where the f block elements start and end.[126][n 10][n 11] It has been claimed that such arguments are proof that, "it is a mistake to break the [periodic] system into sharply delimited blocks".[128] A third variant shows the two positions below yttrium as being occupied by the lanthanides and the actinides.
Chemical and physical arguments have been made in support of lutetium and lawrencium[129][130] but the majority of authors seem unconvinced.[131] Most working chemists are not aware there is any controversy.[132] In December 2015 an IUPAC project was established to make a recommendation on the matter.[133]
Lanthanum and actinium
La and Ac below Y
Lanthanum and actinium are commonly depicted as the remaining group 3 members.[134][n 12] It has been suggested that this layout originated in the 1940s, with the appearance of periodic tables relying on the electron configurations of the elements and the notion of the differentiating electron. The configurations of caesium, barium and lanthanum are [Xe]6s1, [Xe]6s2 and [Xe]5d16s2. Lanthanum thus has a 5d differentiating electron and this establishes "it in group 3 as the first member of the d-block for period 6".[135] A consistent set of electron configurations is then seen in group 3: scandium [Ar]3d14s2, yttrium [Kr]4d15s2 and lanthanum [Xe]5d16s2. Still in period 6, ytterbium was assigned an electron configuration of [Xe]4f135d16s2 and lutetium [Xe]4f145d16s2, "resulting in a 4f differentiating electron for lutetium and firmly establishing it as the last member of the f-block for period 6".[135] 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.[135] Lanthanum, however, 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.[136]
In terms of chemical behaviour,[137] and trends going down group 3 for properties such as melting point, electronegativity and ionic radius,[138][139] scandium, yttrium, lanthanum and actinium are similar to their group 1–2 counterparts. In this variant, the number of f electrons in the most common (trivalent) ions of the f-block elements consistently matches their position in the f-block.[140] 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.[141]
Lutetium and lawrencium
Lu and Lr below Y
In other tables, lutetium and lawrencium are the remaining group 3 members.[n 13] Early techniques for chemically separating scandium, yttrium and lutetium relied on the fact that these elements occurred together in the so-called "yttrium group" whereas La and Ac occurred together in the "cerium group".[135] Accordingly, lutetium rather than lanthanum was assigned to group 3 by some chemists in the 1920s and 30s.[n 14] Several physicists in the 1950s and 60s favoured lutetium, in light of a comparison of several of its physical properties with those of lanthanum.[135] This arrangement, in which lanthanum is the first member of the f-block, is disputed by some authors since lanthanum lacks any f-electrons. However, it has been argued that this is not valid concern given other periodic table anomalies—thorium, for example, has no f-electrons yet is part of the f-block.[142] As for lawrencium, its gas phase atomic electron configuration was confirmed in 2015 as [Rn]5f147s27p1. Such a configuration represents another periodic table anomaly, regardless of whether lawrencium is located in the f-block or the d-block, as the only potentially applicable p-block position has been reserved for nihonium with its predicted configuration of [Rn]5f146d107s27p1.[143][n 15]
Chemically, scandium, yttrium and lutetium (and presumably lawrencium) behave like trivalent versions of the group 1–2 metals.[145] On the other hand, trends going down the group for properties such as melting point, electronegativity and ionic radius, are similar to those found among their group 4–8 counterparts.[135] In this variant, the number of f electrons in the gaseous forms of the f-block atoms usually matches their position in the f-block. For example, the f-electron counts for the first five f-block elements are La 0, Ce 1, Pr 3, Nd 4 and Pm 5.[135]
Lanthanides and actinides
Markers below Y
A few authors position all thirty lanthanides and actinides in the two positions below yttrium (usually via footnote markers). This variant emphasizes similarities in the chemistry of the 15 lanthanide elements (La–Lu), possibly at the expense of ambiguity as to which elements occupy the two group 3 positions below yttrium, and a 15-column wide f block (there can only be 14 elements in any row of the f block).[n 16]
Groups included in the transition metals
The definition of a transition metal, as given by IUPAC, is an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell.[146] By this definition all of the elements in groups 3–11 are transition metals. The IUPAC definition therefore excludes group 12, comprising zinc, cadmium and mercury, from the transition metals category.
Some chemists treat the categories "d-block elements" and "transition metals" interchangeably, thereby including groups 3–12 among the transition metals. In this instance the group 12 elements are treated as a special case of transition metal in which the d electrons are not ordinarily involved in chemical bonding. The 2007 report of mercury(IV) fluoride (HgF4), a compound in which mercury would use its d electrons for bonding, has prompted some commentators to suggest that mercury can be regarded as a transition metal.[147] Other commentators, such as Jensen,[148] 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.[148]
Still other chemists further exclude the group 3 elements from the definition of a transition metal. They do so on the basis that the group 3 elements do not form any ions having a partially occupied d shell and do not therefore exhibit any properties characteristic of transition metal chemistry.[149] In this case, only groups 4–11 are regarded as transition metals. However, though the group 3 elements show few of the characteristic chemical properties of the transition metals, they do show some of their characteristic physical properties (on account of the presence in each atom of a single d electron).[51]
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 the 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.[150] More recent experiments have suggested, however, that flerovium behaves chemically like lead, as expected from its periodic table position.[151]
Further periodic table extensions
Main article: Extended periodic table
Periodic table with eight rows, extended to element 172[152]
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.[153] More recently, physicists such as Pekka Pyykkö have theorized that these additional elements do not follow the Madelung rule, which predicts how electron shells are filled and thus affects the appearance of the present periodic table. There are currently several competing theoretical models for the placement of the elements of atomic number less than or equal to 172.[154]
Element with the highest possible atomic number
The number of possible elements is not known. A very early suggestion made by Elliot Adams in 1911, and based on the arrangement of elements in each horizontal periodic table row, was that elements of atomic weight greater than 256± (which would equate to between elements 99 and 100 in modern-day terms) did not exist.[155] A higher—more recent—estimate is that the periodic table may end soon after the island of stability,[156] which is expected to centre around element 126, as the extension of the periodic and nuclides tables is restricted by proton and neutron drip lines.[157] Other predictions of an end to the periodic table include at element 128 by John Emsley,[3] at element 137 by Richard Feynman,[158] and at element 155 by Albert Khazan.[3][n 17]
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.[159] Hence the non-relativistic Bohr model is inaccurate when applied to such an element.
Relativistic Dirac equation
The relativistic Dirac equation has problems for elements with more than 137 protons. For such elements, the wave function of the Dirac ground state is oscillatory rather than bound, and there is no gap between the positive and negative energy spectra, as in the Klein paradox.[160] 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;[161] however, this does not happen if the innermost orbital is filled, so that element 173 is not necessarily the end of the periodic table.[158]
Optimal form
The many different forms of periodic table have prompted the question of whether there is an optimal or definitive form of periodic table. The answer to this question is thought to depend on whether the chemical periodicity seen to occur among the elements has an underlying truth, effectively hard-wired into the universe, or if any such periodicity is instead the product of subjective human interpretation, contingent upon the circumstances, beliefs and predilections of human observers. An objective basis for chemical periodicity would settle the questions about the location of hydrogen and helium, and the composition of group 3. Such an underlying truth, if it exists, is thought to have not yet been discovered. In its absence, the many different forms of periodic table can be regarded as variations on the theme of chemical periodicity, each of which explores and emphasizes different aspects, properties, perspectives and relationships of and among the elements.[n 18]
LutetiumThoriumProtactiniumNeptuniumAmericiumCuriumBerkeliumCaliforniumEinsteiniumFermiumMendeleviumNobeliumLawrenciumGrapheneWaterChemical element
This article uses material from the Wikipedia article "Periodic table", which is released under the Creative Commons Attribution-Share-Alike License 3.0. There is a list of all authors in Wikipedia


I was over taking pregnancy tests. The last one I took was Christmas day because all I wanted was to find out when I had family around. So when it was negative (again,) I gave up on that family dream. When the doctors asked to verify my blood for pregnancy before issuing any medicine, I left before they had a chance to tell me "no." It felt like the only thing worse than hearing it from a stick, was to hear it from a medical professional.
So, I got it my car, cried, drove home where my sister's entire family was visiting, picked them up, and drove to Disney World. About half way to the park is when I got a phone call from those pesky medical professionals, with my family in my car, making all of my Christmas wishes to heaven come true.
I could glamorize how amazing the first trimester of growing a human inside of your body feels… but there was not an ounce of glam. Every scent in Disney World became my worst enemy, every song made me cry, and every kid made me anxious. The second I got that phone call, I began researching what the cans and can'ts are, and how the heck I was going to get through seven more months of my face being over the bathroom throne.
I have tried to start off my mornings with a hot cup of herbal tea. The first teas I tried were just ginger, but the ginger AND lemon tastes the best to me! They make me feel at ease and refreshed. It's also been amazing for nausea and sinus congestion. The mints I found in nearly every pinterest post, but I can only take a few a day before they're ineffective with hunger nausea.
My sensitive nose didn't stop at just "bad smells," it was offended by anything that didn't smell clean. Some of our home cleaning products were too harsh smelling, but this J.R. Watkins room freshener in Coconut is gentle and natural smelling enough to use in my bedroom. One of our large dogs is sharing the bed with me, and although she is adorable, she is the cause of a lot of discomfort.
Luckily I stopped using perfumes and started using this fragrance oil from Whole Foods (thanks to the recommendation of Shay from Pretty Little Liars.) The scent she recommended was "Persian Garden," which thankfully, I haven't stopped loving.
As if I wasn't already terrified of germs, learning about the pregnacy added an abundance of soaps and sanitizers in our home, cars, and pockets. My favorite travel sized ones have been the aromatherapy Stress Relief (eucalyptus spearmint) and Firework Fizz from Bath & Body Works.
I am not afraid or truly concerned about stretch marks, but I have heard of how itchy and dry your tummy can feel the second half of pregnancy. When I know change is coming, I try to change my habits immediately, instead of waiting to discipline myself later (which is also why i started sleeping with a preggo pillow on my left side six weeks early.) The Burt's Bee's Belly Butter is a natural fragrance free formula containing Cocoa, Shea, and Jojoba Butters. After drying off from a shower, I apply it to all the places on my body that won't stop growing.
I have had water bottles in my favorites before! Water is my favorite since it is liquid gold traveling to places that don't have clean water sources. Not only does it save the earth having a water bottle, but you can save a lot of cash! Airports, movie theaters, and events all sell a single water bottle for basically the price of a case. We have the Zero Water filter in our fridge, and my new favorite bottle is the 32oz. Miracol bottle.
I was going to share my favorite foods, but lets be real… no one wants to see or hear about food at this stage!
I hope that this was the slightest bit helpful to you or a loved one! I am excited to eventually share about remaining a vegetarian through this pregnancy, trimester favorites, the journey of basically going through it alone with my husband deployed,and all the amazing collaborations coming up!
For all maternity related posts:click here.


text:
Ich suche » Farrell and what to look forward to this off-seaso
JobKambio » Forum » Ich suche » Farrell and what to look forward to this off-seaso
#1 | Farrell and what to look forward to this off-seaso 09.09.2019 10:05
TORONTO - When Montreal Canadiens goaltender Carey Price looks at the Toronto Maple Leafs, he sees a big, physically strong team. "Theyre in your face," Price said. "Theyre a playoff team." The Leafs may be built for playoff hockey, but getting there could be a challenge down the stretch, especially after losing 4-3 to the Canadiens on Saturday night at Air Canada Centre. Montreal moved five points up on Toronto in the Atlantic Division standings and in the process made it far less likely that this thrilling, back-and-forth affair could be a first-round playoff preview. "I think thats what everybody would want to see, and Im sure its going to happen sooner or later," said Habs forward Rene Bourque, who had a goal and an assist and was not yet born the last time these teams met in the playoffs back in 1979. As the Habs won for the fourth time in five games, the Leafs (36-28-8) dropped their fourth in a row as part of a troubling late-season swoon in the absence of injured goalie Jonathan Bernier. Toronto still occupies the first wild-card spot in the Eastern Conference but leads Detroit by just one point and Columbus by two and has played two more games that the Red Wings and Blue Jackets. The Leafs could fall out of playoff position by the end of Sunday. James Reimer, who allowed four goals on 37 shots, including a short-side game-winner by Tomas Plekanec, said he and his teammates have a "healthy sense of urgency" with 10 games remaining. "I think we know we played well tonight and I think we made some mistakes but I think they made some mistakes too," Reimer said. "Really, it was kind of one bad bounce that decided the game. I think we can hold our heads high on this one and go into tomorrow (at the New Jersey Devils) feeling good about ourselves." The Habs are feeling good about themselves after winning a track meet of a hockey game that featured three goals in the first period — by Montreals Max Pacioretty, Bourque and captain Brian Gionta and Torontos Joffrey Lupul and Tyler Bozak. Montreal coach Michel Therrien liked the way his team dictated the play early. But it was bouncing back in the third period after Nazem Kadri tied the score for the Leafs that impressed Price, who finished with 33 saves. "That was definitely a test of character," Price said. "When a team scores in the third period to tie it up when theyre at home, youre on the road, they grab a lot of momentum. ... Being able to grab the lead and then hold it with a good team effort like that is, I think, rewarding." The Habs reward if these winning ways keep up is either second or third place in the Atlantic Division, crucial spots that would mean avoiding Mondays opponent, the Boston Bruins, or the Pittsburgh Penguins in the first round. Now five points back of Montreal and six behind the Tampa Bay Lightning, who won in Toronto on Wednesday night, the Leafs will need an uphill climb to avoid a wild-card spot and a tough matchup. But now theres some legitimate concern, given recent woes, that a playoff spot altogether might be in danger. This loss, which winger Mason Raymond said "stings" and captain Dion Phanuef called "disappointing," didnt help that cause. "You have to turn the page," Phaneuf said. "Theres no looking back on today, tomorrow. Tomorrows a new day, weve got to pull ourselves out of it. Its this group thats going to get us out of this bind. Weve been close, but close isnt good enough right now. We know that were going to get out of it, weve just got to find a way." The Leafs and Habs (29-26-7) were close because this was such a back-and-forth game with plenty of chances at either end. The turning point came 9:14 into the third when Toronto winger James van Riemsdyk was called for goaltender interference when he steam-rolled Price. Habs defenceman Andrei Markov appeared to make contact with van Riemsdyk before he hit Price, but the goalie went to the ice and believed he was interfered with. "Its contact to my head, so I thought it was a penalty, personally," Price said. Leafs coach Randy Carlyle, who conceded he didnt see a replay of the incident, was more worried about the impact of the penalty. Plekanec scored just as van Riemsdyk was being released from the box at 11:14, sneaking a shot in the tiny space between Reimer and the right post. "We clawed back into the hockey game and then we take a penalty early in the third and they score," Carlyle said. "It was the difference in the hockey game, and the margin of error now in these games is so close that one bounce or one mis-play or one unfortunate mistake cost us points." Mistakes — be it turnovers by Phil Kessel, David Clarkson and Kadri or soft goals allowed by Reimer — hurt the Leafs dearly. Another slow start wasnt ideal, either. "Obviously we were behind the 8-ball a little bit early on giving up two goals," Raymond said of allowing two goals in the games first seven minutes. "I like the way we battled back, but turnovers killed us a bit and thats tough." The Habs benefited because they pounced on so many mistakes the Leafs committed. In front of an early-spring crowd of 19,789 thats hungry for playoff hockey, Montreal executed like a team ready for that next step. "Both teams need those points. We were ready to play," Therrien said. "At this time of the year, you need to be in a playoff mindset. And our mindsets like that." NOTES — Montreals streak of consecutive penalties killed was snapped at 25 on Kadris power-play goal at 2:49 of the third. The last time the Habs surrendered a goal on the power play was March 6 at the Phoenix Coyotes. ... Leafs centre Dave Bolland played just 9 minutes 1 second in his return to the lineup. Bolland missed the previous 56 games after suffering a severed tendon in the back of his ankle Nov. 2. ... Habs forward Lars Eller suffered a lower-body injury, Therrien said, and did not play in the third period. Michael Bournival was called up, and he will meet the team in Boston. .... Bernier, who has now missed four straight games with a groin injury, will miss his fifth in a row Sunday when the Leafs visit the Devils. Carlyle said the 25-year-old will not make the trip. Carlo Pinsoglio Jersey . The Senators return from a lengthy layoff caused by Wednesdays attack on Parliament Hill to host the New Jersey Devils on Saturday night. Alex Sandro Juventus Jersey . The incident occurred at 6:28 of the first period in Anaheims 6-3 home win over Dallas on Sunday. Garbutt left the penalty box and skated at Penner before leaving his skates to deliver a check. http://www.juventusfcpro.com/Kids-Blaise-Matuidi-Jersey/ . LOUIS -- The St. Federico Bernardeschi Jersey . This week, topics cover the World Series champion Red Sox, John Farrell and what to look forward to this off-season. Andrea Barzagli Juventus Jersey . has left the San Jose Sharks to become the Boston Bruins director of player personnel.Brad Jacobs helped Canada hold on for a slim 7-5 win over Great Britain in a key round-robin match in the mens Olympic curling tournament. Great Britains David Murdoch gave his rink a 2-1 lead after three ends, before the Canadians (5-2) picked up three points in the fourth. Murdoch (5-2) replied with two in the fifth, but Jacobs earned a point in the seventh end for a 5-4 lead..dddddddddddd Murdoch tied it 5-5 in the eighth before Canada took the lead again in the ninth. Britain had an opportunity to tie, but overcurled on their final shot of the match, giving Canada an extra point for the 7-5 victory. Canada and Great Britain are now tied for third in the round robin standings behind China and Sweden, who have identical 6-1 records. ' ' '
2-1 win over the Minnesota Wild on Saturday » « first round -- in the 2014 draft.

redpajama_set_name: RedPajamaCommonCrawl


National agricultural retailer Mole Valley Farmers wanted to create a series of videos to launch their own brand product, Deltamole. Mole Valley wanted to use real customers in the video to advocate, narrate and demonstrate the product being applied especially for the how to video because they wanted to show the product in it's true light.
After extensive research, concept reviews and storyboarding we created 3 videos, a demonstration video and two web adverts. All three where filmed on location, filming was tricky because we halve to get the shots live whist milking the cows. The Mole Valley Farmers' customer was fantastic and relaxed into the filming despite being watched by his team.
The video production aided the successful launch of Deltmole.
Customer Services can direct customers to the how to video for instruction on use.
Due to the success of the video Mole Valley have since been back for more.


text:
/**
 * Created by diegob on 26.09.15.
 */

(function() {

  /**
   * @constructor
   */
  diga.UserSelectController = function($mdDialog, users) {
    this.$mdDialog_ = $mdDialog;
    this.users = users;
  };

  diga.UserSelectController.prototype.assignUser = function(user) {
    this.$mdDialog_.hide(user)
  }

  angular.module('diga').controller('UserSelectController', ['$mdDialog', 'users', diga.UserSelectController]);
})();

redpajama_set_name: RedPajamaGithub


text:
Frederick H. Dyer, Compendium of the War of the Rebellion: Regimental Histories
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Browsing named entities in a specific section of Frederick H. Dyer, Compendium of the War of the Rebellion: Regimental Histories. Search the whole document.
Found 2,988 total hits in 710 results.
March 13th (search for this): chapter 27
sion, 18th Army Corps, Army of the James, to July, 1864. 1st Brigade, 1st Division, 18th Army Corps, to December, 1864. 1st Brigade, 3rd Division, 24th Army Corps, Dept. of Virginia, to June, 1865. Service. Duty near Fort Albany, Defenses of Washington, till December 4, 1862. March to Falmouth, Va., December 5-9. Battle of Fredericksburg December 12-15. Burnside's Second Campaign Mud March January 20-24, 1863. Moved to Newport News, Va., February 9, thence to Suffolk March 13. Siege of Suffolk April 12-May 4. Providence Church Road, Nansemond River, May 3. Reconnoissance across the Nansemond May 4. Moved to Portsmouth May 13, thence to Yorktown. Dix's Peninsula Campaign June 24-July 7. Expedition from White House to South Anna River July 1-7. Moved to Portsmouth July 8-14; thence to Julian Creek July 30, and duty there till March 19, 1864. Moved to Yorktown March 19. Butler's operations on south side of James River and against Petersburg
ppi, to June, 1865. Service. Duty at Franklin, La., till January, 1864. Reenlisted January 4, 1864. Ordered to New Orleans, La., and duty there till March. Red River Campaign March 10-May 22. Advance from Franklin to Alexandria March 14-26. Action at Natchitoches March 31. Crump's Hill April 2. Wilson's Farm April 7. Bayou de Paul Carroll's Mill April 8. Battle of Pleasant Hill April 9. Natchitoches April 19 and 22. Monett's Bluff and Cane River Crossing Ap62. Movement to Falmouth, Va., October 27-November 19. Battle of Fredericksburg, Va., December 12-15. Burnside's Second Campaign ( Mud March ) January 20-24, 1863. Moved to Newport News, Va., February 9, thence to Norfolk and Suffolk March 14. Siege of Suffolk April 12-May 4. Battery Huger, Hill's Point, April 19. Reconnoissance across Nansemond River May 4. Moved to Portsmouth May 13, thence to Yorktown, Va. Dix's Peninsula Campaign June 24-July 7. Expedition from Whi
ps, Army of the Potomac, to June, 1865. Service. Duty at Munson's Hill, Defenses of Washington, D. C., till March, 1862. Advance on Manassas, Va., March 10-15. Camp at Upton's Hill till April 9. Advance on Falmouth, Va., April 9-18. Occupation of Fredericksburg April 18 and duty there till May 25. McDowell's a-11. At Camp California, near Alexandria, Va., till March 10, 1862. Scout to Burke's Station January 17, 1862 (Co. A ). Advance on Manassas, Va., March 10-15. Reconnoissance to Gainesville March 20, and to Rappahannock Station March 28-29. Warrenton Junction March 28. Moved to the Virginia Peninsula April 4. anchester and mustered in December 23, 1861. Left State for Boston, Mass., January 24, 1862; thence sailed for Ship Island, Miss., February 15, arriving there March 15. Attached to Butler's New Orleans Expedition to March, 1862. 1st Brigade, Dept. of the Gulf, to November, 1862. Independent Command, Dept. of the Gulf,
ass Region of Kentucky till June. Moved to Vicksburg, Miss., June 3-14. Siege of Vicksburg, Miss., June 14-July 4. Advance on Jackson, Miss., July 4-10. Siege of Jackson, Miss., July 10-17. At Milldale, Miss., till August 10. Moved to Covington, Ky., August 10-21; thence to Camp Nelson, Ky., August 25. Duty guarding railroad between Cincinnati, Ohio, and Camp Nelson, Ky., till January, 1864. Moved to Camp Burnside January 15. March to Knoxville, Tenn., February 19-March 17. March across Cumberland Mountains to Camp Burnside and Nicholasville, Ky., March 21-31. Moved to Annapolis, Md., April 2-5. Campaign from the Rapidan to the James River, Va., May 3-June 15. Battles of the Wilderness May 5-7; Spottsylvania May 8-12; Spottsylvania Court House May 12-21. Assault on the Salient at Spottsylvania Court House May 12. North Anna River May 23-26. On line of the Pamunkey May 26-28. Totopotomoy May 28-31. Cold Harbor June 1-12. Bethesda
to Barnett's Ferry September 27-28. Moved to Aikens Landing October 1. Duty in trenches before Richmond till March 3, 1865. Battle of Fair Oaks October 27-28, 1864. Moved to Fort Monroe, Va., March 4-5; thence to White House Landing March 18 to establish a depot for General Sheridan's Cavalry, and duty there till March 24. March to lines north of the James March 24-28. Occupation of Richmond April 3. Duty there and at Manchester till July. Provost duty in District of Norpaign November 4-December 23. Campbell Station November 16. Siege of Knoxville November 17-December 4. Repulse of Longstreet's assault on Fort Saunders November 29. Duty in East Tennessee till March, 1864. Moved to Annapolis, Md., March 18-April 7. Campaign from the Rapidan to the James River, Va., May 3-June 15. Battles of the Wilderness May 5-7. Spottsylvania May 8-12. Spottsylvania Court House May 12-21. Assault on the Salient, Spottsylvania C. H., May 12. Nor
Julien Creek July 30. Duty there till March 19, 1864. Ballahock, on Bear Quarter Road, and Deep Creek, February 29-March 1, 1864. Moved to Great Bridge March 19, thence to Yorktown April 19. Butler's operations on south side of the James River against Petersburg and Richmond May 4-28. Port Walthal Junction May 7. te House to South Anna River July 1-7. Moved to Portsmouth July 8-14; thence to Julian Creek July 30, and duty there till March 19, 1864. Moved to Yorktown March 19. Butler's operations on south side of James River and against Petersburg and Richmond May 4-28. Port Walthal Junction, Chester Station, May 6-7. Swift Ce Defenses of Bermuda Hundred December 18-30. Duty at City Point till March 19, 1865. Reported to Gen. Parke, Commanding 9th Army Corps, before Petersburg, March 19. Repulse of attack on Fort Stedman March 25. Duty at Fort Stedman till April 2. Appomattox Campaign March 28-April 9. Assault on and fall of Petersbu
on Head, S. C., till April, 1862. Affair Hunting Island. March, 1862. Reconnoissance up Savannah River to Elba Island March 7-11. Expedition to Bluffton March 20-24. Occupation of Edisto Island April 5. Affair at Watts' Court April 10. Reconnoissance of Seabrook Island April 14. Advance on Jehossie Island Apriria, Va., till March 10, 1862. Scout to Burke's Station January 17, 1862 (Co. A ). Advance on Manassas, Va., March 10-15. Reconnoissance to Gainesville March 20, and to Rappahannock Station March 28-29. Warrenton Junction March 28. Moved to the Virginia Peninsula April 4. Siege of Yorktown, Va., April 5-May 4. rper's Ferry, W. Va., February 3, thence moved to Cumberland, Md., and return to Washington February 25. Ordered to New Orleans, La., and sailed from New York March 20. Duty at Camp Parapet, Carrollton, Jefferson City and along Lake Pontchartrain till June. Ordered to Morganza, La., June 7. Movement to Fortress Monroe,
March 21st (search for this): chapter 27
2nd Expedition to Fort Fisher January 7-15, 1865. Assault and capture of Fort Fisher January 15. Sugar Loaf Battery February 11. Fort Anderson February 18. Capture of Wilmington February 22. Advance on Kinston and Goldsboro March 6-21. Guard Railroad from Little Washington to Goldsboro till August. Mustered out August 23, 1865. Regiment lost during service 3 Officers and 82 Enlisted men killed and mortally wounded and 5 Officers and 194 Enlisted men by disease. Total 2 railroad between Cincinnati, Ohio, and Camp Nelson, Ky., till January, 1864. Moved to Camp Burnside January 15. March to Knoxville, Tenn., February 19-March 17. March across Cumberland Mountains to Camp Burnside and Nicholasville, Ky., March 21-31. Moved to Annapolis, Md., April 2-5. Campaign from the Rapidan to the James River, Va., May 3-June 15. Battles of the Wilderness May 5-7; Spottsylvania May 8-12; Spottsylvania Court House May 12-21. Assault on the Salient at Spott
8, 1864. Moved to Fort Monroe, Va., March 4-5; thence to White House Landing March 18 to establish a depot for General Sheridan's Cavalry, and duty there till March 24. March to lines north of the James March 24-28. Occupation of Richmond April 3. Duty there and at Manchester till July. Provost duty in District of NMarch 24-28. Occupation of Richmond April 3. Duty there and at Manchester till July. Provost duty in District of Northern Neck, Dept. of Virginia, till December. Mustered out December 19, 1865. Regiment lost during service 15 Officers and 163 Enlisted men killed and mortally wounded and 6 Officers and 166 Enlisted men by disease. Total 350. 3rd New Hampshire Regiment Infantry. Organized at Concord and mustered in August 23, 1861.d, S. C., till April, 1862. Affair Hunting Island. March, 1862. Reconnoissance up Savannah River to Elba Island March 7-11. Expedition to Bluffton March 20-24. Occupation of Edisto Island April 5. Affair at Watts' Court April 10. Reconnoissance of Seabrook Island April 14. Advance on Jehossie Island April 17.
Veterans mustered out October 12, 1864. Reconnoissance to Hatcher's Run December 9-10. Dabney's Mills, Hatcher's Run, February 5-7, 1865. Watkins' House March 25. Appomattox Campaign March 28-April 9. On line of Hatcher's and Gravelly Runs March 29-30. Hatcher's Run or Boydton Road March 31. White Oak Road Marksburg, Va., December 12-15. Burnside's Second Campaign, Mud March, January 20-24, 1863. Moved to Newport News, Va., February 11; thence to Lexington, Ky., March 25-31. Duty in the Blue Grass Region of Kentucky till June. Moved to Vicksburg, Miss., June 3-14. Siege of Vicksburg, Miss., June 14-July 4. Advance on Duty at City Point till March 19, 1865. Reported to Gen. Parke, Commanding 9th Army Corps, before Petersburg, March 19. Repulse of attack on Fort Stedman March 25. Duty at Fort Stedman till April 2. Appomattox Campaign March 28-April 9. Assault on and fall of Petersburg April 2. Occupation of Petersburg April 3


text:
Keeping Tobacco Sacred
Tobacco, before it became the most enslaved and abused plant in the world, was a sacred gift, a medicine, a prayer of the American Indian people.
Today, many tribal religious and spiritual/cultural leaders and teachers continue to honor the sacred traditions and values of tobacco. Though as a young man I had studied the stories and been to many ceremonies of the Sacred Pipe and Tobacco, it was first from an Elder of Taos Pueblo that I felt emotionally and spiritually the real power of tobacco. I first met him when I was just starting out traveling with the Sacred Pipe. This was around 1976 or so.
Strawberry Jim and the Sacred Tobacco
He was very welcoming of the sacred pipe into his home. On my very first visit to his house I stayed for nearly a month. It was right in the middle of winter and soon after I arrived there was a gigantic snow fall, several feet. This stranded me in the house. After a few days I started digging through the supplies in my old Chevy. I came up with a jar of homemade strawberry preserves. At the time I just called him "Grandfather" and it was an interesting surprise to see how much he relished these strawberry preserves. I think that jar of strawberries made me a lifelong friend on the spot. In the days that followed we had a warm and lasting friendship. Every day we had together a small and private ceremony with the sacred pipe. He really relished the meditation, as did I. We would sit in his small but outstanding kitchen, in his old, old adobe home and smoke together. It was only later that I learned everyone locally called him Strawberry Jim. He used to grow really good strawberries and bring them into town on his horse-drawn cart back many years ago.
He loved to share the smoke of the Sacred Pipe and his reverence and sweetness taught me the full emotional and spiritual meaning of sacred smoking. Even after all these years and so many ceremonies, Strawberry Jim smoked with great power, dignity and beauty.
When he smoked, the smoke itself would travel from his lips in a tiny, tight column, thin as a pencil and straight. It would come out of his mouth and climb straight up into the room like a snake, like a column of pure white light. It was very beautiful and something I only saw when Strawberry Jim smoked and prayed with the sacred pipe.
Jim spoke almost no english. He did speak Spanish very well, as well as his own Tewa language. I never found this to be a problem. Though I do not speak Spanish or his language we always managed to understand each other perfectly well. For many years, until his death, I came to visit him every year. This became a very important and much anticipated journey for me. When he was very ill and shortly before he died I gathered in his small bedroom with all the members of his family. There in the dark room, lighted by the old pueblo fireplace, I made his farewell song and pressed for the last time the long stem of the sacred pipe to his ancient lips.
Strwaberry Jim was gracious. he was beautiful, calm and very well spoken. He was held in high esteem by every person who ever came in contact with him. Having his friendship was one of the great honors of my life. His gentleness and smiling face remains an endless teaching for me about what a real indian immersed in his sacred path really is all about.
It was interesting to have this experience there at Taos Pueblo. Traditionally it a handful of tribes in the north that follow the Sacred Pipe. The Sioux tribes, the Ahnishinabe, Cree and Algonquian and so forth. So much has been lost. Many tribes in the southeast had complex sacred smoking ceremonies that have all now been lost.
Tobacco as we know it today was the religious property of the Creek Indians, at the southern end of the Mississippi River. My people are at the north mouth of the same river. In a short version of the coming of sacred tobacco: a couple had gone into the forest to make love. They had a great time in a big meadow filled with a deep leafy plant they had never seen before. When they returned to the village after their day of making love, it turns out they had been gone many years and everyone thought they had died and disappeared into the deep old forest. They lead the elders to the meadow and this plant began its long ceremonial and sacred use among these fantastic old Creek Indian people. That tobacco plant is now called "Virginia Gold Leaf" and it is the finest tobacco in the world, in the history of the world. It was stolen from these tribal elders. They never gave their permission for its commercial enslavement and remains their legal property to this day. They have never been paid.
Smokin all alone the strange assortment of additives and the genetically altered tobacco of the modern tobacco industry is certainly dangerous and well should it be. Tobacco is one of the most powerful plants in the history of the world.
The old indians understood that the secret to its power was in smoking together. Smoking the sacred pipe it becomes a symbol and a ceremony for passing life through the body and the soul at the same time. We like to say, "in a visible manner" as we believe the smoke has a special power to be seen and understood by God, by the Mystery Life, by the Spirits of the Mother Earth.
Smoking together is everything. People do this with a marijuana joint. Getting stoned together with a joint is much more closely connected to the old sacred ceremonies than sucking on a Marlboro in a bar in the middle of the night.
Indian teachers know that there is a mysterious power that happens when a group of people with open hearts and well informed spirits gather together and make a ceremony of sharing the smoke. It is a power that works to bring something magical and sacred from the people when they do this together. This sacred experience of tobacco does not work when you smoke alone.
Try smoking your cigarette with your friends, pass it around. You will see the energy of what you are doing will change instantly. Where you were before isolated inside a compulsive habit, now something else entirely is happening. Try it.
American Indian Ritul Pantelleria Sacred Pipe tobacco
I Am A Wa-beeno : Part Two
I Am A Wa-beeno : Part One
Alignment (Shadows of Sacred Stones)


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