B C D No Idea 2 If The Electronegativities Of H And O: Fill & Download for Free

Then solve old questions from different years,you will get an idea of the question pattern. You too think how a question setter sets a question for the students.Let us take an example. Suppose you know“Oxidation number of H in hydride is -1.”You have to make a multiple question from above fact.Q.1.What is the oxidation number of H in CuH?a) -1 b)+1 c) O d)-2Ans: will be b) +1 ……..(not -1)You know in hydrides the o.n. of H is -1 but in above question you find the o.n. of H is +1 i.e. an exceptional case.Cu is more electronegative than H,which is written in electrochemical series.So from each fact you try to find multiple questions by using :: how/when/exceptional cases if any.All the best.

What is Inductive Effect?When an electron-releasing or an electron-withdrawing species is introduced to a chain of atoms (generally a carbon chain), the corresponding negative or positive charge is relayed through the carbon chain by the atoms belonging to it. This causes a permanent dipole to arise in the molecule and is referred to as the inductive effect.An illustration describing the inductive effect that arises in a chloroethane molecule due to the more electronegative chlorine atom is provided above.Also ReadIntroduction to Organic ChemistryPurification of Organic CompoundsLassaigne’s testVictor Meyer’s MethodInductive Effect on Acidity and BasicityUsing the inductive effect, we can predict the acidity and basicity of compounds. It can be said as a generalisation the electron withdrawing groups (EWG) increase the acidity of a compound and electron donating group decrease the acidity of a compound.This is because, if we take the conjugate base of the acid, that is, RCOO-, if R is electron withdrawing, then the conjugate base is stabilised via delocalisation of the formed negative charge.If R had been electron donating, then the conjugate base would be destabilised because of inter-electronic repulsions.Inductive Effect on Acidity and Basicity of CompoundsThus it can be said that, +I groups decrease acidity (or increase basicity) and –I groups increase acidity (or decrease basicity) of compounds.For Example, formic acid ( HCOOH) is more acidic than acetic acid (CH3COOH)due to the +I inductive effect of the methyl group attached to the carboxylic acid group.Note: If Ka of acid is high, it is a strong acid, but if PKa of acid is high, it is said to be a weak acid [pka = -log(ka)] Same logic applies to bases.Consider, the acidity of mono-, di- and trichloroacetic acid.It can be said that the presence of three Cl atoms make oxygen highly electron deficient and thereby, polarising the O-H bond the most. Therefore the acidity order for the above compounds would be, III > II > I.Types of Inductive EffectNegative inductive effect or -I effectPositive inductive effect +I effect-I Effect (Negative Inductive Effect)When an electronegative atom, such as a halogen, is introduced to a chain of atoms (generally carbon atoms), the resulting unequal sharing of electrons generates a positive charge which is transmitted through the chain.This causes a permanent dipole to arise in the molecule wherein the electronegative atom holds a negative charge and the corresponding effect is called the electron withdrawing inductive effect, or the -I effect.+I Effect (Positive Inductive Effect)When a chemical species with the tendency to release or donate electrons, such as an alkyl group, is introduced to a carbon chain, the charge is relayed through the chain and this effect is called the Positive Inductive Effect or the +I EffectInductive Effect on Stability of MoleculesThe charge on a given atom and the charge on a group bonded to the atom play a strong part when determining the stability of the resulting molecule as per the inductive effect.An example of this can be observed when a group displaying the -I effect is bonded to a positively charged atom and the positive charge on the resulting molecule is amplified, reducing its stability.On the other hand, when a negatively charged atom is introduced to a group displaying a -I effect, the charge disparity is somewhat quenched and the resulting molecule would be stable as per the inductive effect.Also,When a group displaying the -I effect is bonded to a molecule, the electron density of the resulting molecule effectively reduces, making it more likely to accept electrons and thereby increasing the acidity of the molecule.When a +I group attaches itself to a molecule, there is an increase in the electron density of the molecule. This increases the basicity of the molecule since it is now more capable of donating electrons.Applications of Inductive EffectIllustration 1:Give the stability of the following canonical formsThe structure I and III have more covalent bonds and are more stable than II and IV. Between I and III, I is more stable because the negative charge is on an electronegative element.Between II and IV, II is more stable because of the same reason as said above.The order is I > III > II > IVIllustration 2:We know that EWG increases acidity and EDG decreases acidity.-Me group is a +I group whereas -OMe is an +R group, so –OMe decreases the acidity more strongly than -MeTherefore the order is, d>c>e>a>bSince ka is directly proportional to acidity, the answer is a→ t, b →p, c→ s, d→ q, e→ r.Illustration 3:Solution: NaNH2 is a base, therefore, the most acidic proton of the substrate would react to form a conjugate base. The idea here is to find out the most acidic proton.There are totally four protons, -COOH, -OH, nitro-substituted –OH, and alkyne proton.Since two moles of the base are used, two protons would react.The order of acidity of the protons are-COOH>-OH (Nitro substituted)>-OH> acetylenic protonSo the product would be,Illustration 4: The order of acidity of the following compounds, isSolution: To find out the acidity of the compounds, remove the proton and check the stability of the conjugate base formed.CB of structures I and II are stabilised by intramolecular hydrogen bonding ( I more than II).Between the meta and para isomer, meta would be more acidic due to –I effect of oxygen.Therefore, the order is, I>II>III>IVIllustration 5:Solution: The most basic among the four is I. This is because structures II and IV are aromatic. Between I and III, I is more basic due to the presence of an oxygen atom in III, which decreases basicity by –I effect.Between II and IV, II would be more basic because, In IV, the lone pair on nitrogen is delocalised to make the compound aromatic. The non-availability of the lone pair for donation makes IV the least basic.Therefore, the order is, I > III > II > IV.Source: https://byjus.com

The periodic table, or periodic table of elements, is a tabular arrangement of the chemical elements, ordered by their atomic number, electron configuration, and recurring chemical properties, whose structure shows periodic trends. Generally, within one row (period) the elements are metals to the left, and non-metals to the right, with the elements having similar chemical behaviours placed in the same column. Table rows are commonly called periods and columns are called groups. Six groups have accepted names as well as assigned numbers: for example, group 17 elements are the halogens; and group 18 are the noble gases. Also displayed are four simple rectangular areas or blocks associated with the filling of different atomic orbitals.The organization of the periodic table can be used to derive relationships between the various element properties, but also the predicted chemical properties and behaviours of undiscovered or newly synthesized elements. Russian chemist Dmitri Mendeleev was the first to publish a recognizable periodic table in 1869, developed mainly to illustrate periodic trends of the then-known elements. He also predicted some properties of unidentified elements that were expected to fill gaps within the table. Most of his forecasts proved to be correct. Mendeleev's idea has been slowly expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour. The modern periodic table now provides a useful framework for analyzing chemical reactions, and continues to be widely used in chemistry, nuclear physics and other sciences.All the elements from atomic numbers 1 (hydrogen) through 118 (oganesson) have been either discovered or synthesized, completing the first seven rows of the periodic table.[1][2]The first 98 elements exist in nature, although some are found only in trace amounts and others were synthesized in laboratories before being found in nature.[n 1]Elements 99 to 118 have only been synthesized in laboratories or nuclear reactors.[3]The synthesis of elements having higher atomic numbers is currently being pursued: these elements would begin an eighth row, and theoretical work has been done to suggest possible candidates for this extension. Numerous synthetic radionuclides of naturally occurring elements have also been produced in laboratories.Part of a series on thePeriodic tablePeriodic table forms[show]Periodic table history[show]Sets of elementsBy periodic table structure[show]By metallic classification[show]By other characteristics[show]ElementsList of chemical elements[show]Properties of elements[show]Data pages for elements[show]BookCategoryChemistry PortalvteContents1Overview2Grouping methods2.1Groups2.2Periods2.3Blocks2.4Metals, metalloids and nonmetals3Periodic trends and patterns3.1Electron configuration3.2Atomic radii3.3Ionization energy3.4Electronegativity3.5Electron affinity3.6Metallic character3.7Linking or bridging groups4History4.1First systemization attempts4.2Mendeleev's table4.3Second version and further development5Different periodic tables5.1The long- or 32-column table5.2Tables with different structures6Open questions and controversies6.1Placement of hydrogen and helium6.2Group 3 and its elements in periods 6 and 76.2.1Lanthanum and actinium6.2.2Lutetium and lawrencium6.2.3Lanthanides and actinides6.3Groups included in the transition metals6.4Elements with unknown chemical properties6.5Further periodic table extensions6.6Element with the highest possible atomic number6.7Optimal form7See also8Notes9References9.1Bibliography10Further reading11External linksOverviewFor a more detailed periodic table, see Periodic table (detailed cells).hidevtePeriodic tableGroup123456789101112131415161718Alkali metalsAlkaline earth metalsPnicto­gensChal­co­gensHalo­gensNoble gasesPeriod1Hydro­gen1H​1.008He­lium2He​4.00262Lith­ium3Li​6.94Beryl­lium4Be​9.0122Boron5B​10.81Carbon6C​12.011Nitro­gen7N​14.007Oxy­gen8O​15.999Fluor­ine9F​18.998Neon10Ne​20.1803So­dium11Na​22.990Magne­sium12Mg​24.305Alumin­ium13Al​26.982Sili­con14Si​28.085Phos­phorus15P​30.974Sulfur16S​32.06Chlor­ine17Cl​35.45Argon18Ar​39.9484Potas­sium19K​39.098Cal­cium20Ca​40.078Scan­dium21Sc​44.956Tita­nium22Ti​47.867Vana­dium23V​50.942Chrom­ium24Cr​51.996Manga­nese25Mn​54.938Iron26Fe​55.845Cobalt27Co​58.933Nickel28Ni​58.693Copper29Cu​63.546Zinc30Zn​65.38Gallium31Ga​69.723Germa­nium32Ge​72.630Arsenic33As​74.922Sele­nium34Se​78.971Bromine35Br​79.904Kryp­ton36Kr​83.7985Rubid­ium37Rb​85.468Stront­ium38Sr​87.62Yttrium39Y​88.906Zirco­nium40Zr​91.224Nio­bium41Nb​92.906Molyb­denum42Mo​95.95Tech­netium43Tc​[98]Ruthe­nium44Ru​101.07Rho­dium45Rh​102.91Pallad­ium46Pd​106.42Silver47Ag​107.87Cad­mium48Cd​112.41Indium49In​114.82Tin50Sn​118.71Anti­mony51Sb​121.76Tellur­ium52Te​127.60Iodine53I​126.90Xenon54Xe​131.296Cae­sium55Cs​132.91Ba­rium56Ba​137.33Lan­thanum57La​138.91Haf­nium72Hf​178.49Tanta­lum73Ta​180.95Tung­sten74W​183.84Rhe­nium75Re​186.21Os­mium76Os​190.23Iridium77Ir​192.22Plat­inum78Pt​195.08Gold79Au​196.97Mer­cury80Hg​200.59Thallium81Tl​204.38Lead82Pb​207.2Bis­muth83Bi​208.98Polo­nium84Po​[209]Asta­tine85At​[210]Radon86Rn​[222]7Fran­cium87Fr​[223]Ra­dium88Ra​[226]Actin­ium89Ac​[227]Ruther­fordium104Rf​[267]Dub­nium105Db​[268]Sea­borgium106Sg​[269]Bohr­ium107Bh​[270]Has­sium108Hs​[270]Meit­nerium109Mt​[278]Darm­stadtium110Ds​[281]Roent­genium111Rg​[282]Coper­nicium112Cn​[285]Nihon­ium113Nh​[286]Flerov­ium114Fl​[289]Moscov­ium115Mc​[290]Liver­morium116Lv​[293]Tenness­ine117Ts​[294]Oga­nesson118Og​[294]Cerium58Ce​140.12Praseo­dymium59Pr​140.91Neo­dymium60Nd​144.24Prome­thium61Pm​[145]Sama­rium62Sm​150.36Europ­ium63Eu​151.96Gadolin­ium64Gd​157.25Ter­bium65Tb​158.93Dyspro­sium66Dy​162.50Hol­mium67Ho​164.93Erbium68Er​167.26Thulium69Tm​168.93Ytter­bium70Yb​173.05Lute­tium71Lu​174.97Thor­ium90Th​232.04Protac­tinium91Pa​231.04Ura­nium92U​238.03Neptu­nium93Np​[237]Pluto­nium94Pu​[244]Ameri­cium95Am​[243]Curium96Cm​[247]Berkel­ium97Bk​[247]Califor­nium98Cf​[251]Einstei­nium99Es​[252]Fer­mium100Fm​[257]Mende­levium101Md​[258]Nobel­ium102No​[259]Lawren­cium103Lr​[266]1 (red)=Gas 3 (black)=Solid 80 (green)=Liquid 109 (gray)=Unknown Color of the atomic number shows state of matter (at 0 °C and 1 atm)Primordial From decay Synthetic Border shows natural occurrence of the elementStandard atomic weight (Ar)[4]Ca: 40.078 — Formal short value, rounded (no uncertainty)[5]Po: [209] — mass number of the most stable isotopeBackground color shows subcategory in the metal–metalloid–nonmetal trend:MetalMetalloidNonmetalUnknownchemicalpropertiesAlkali metalAlkaline earth metalLan­thanideActinideTransition metalPost-​transition metalReactive nonmetalNoble gasEach chemical element has a unique atomic number (Z) representing the number of protons in its nucleus.[n 2]Most elements have differing numbers of neutronsamong 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.[6]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.[7]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.[8][9][10][11]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).[12]Grouping methodsGroupsMain 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.[13]Consequently, elements in the same group tend to have a shared chemistry and exhibit a clear trend in properties with increasing atomic number.[14]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.[15][16][17]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).[18]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.[19]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),[18]since they display fewer similarities and/or vertical trends.Elements in the same group tend to show patterns in atomic radius, ionization energy, and electronegativity. From top to bottom in a group, the atomic radii of the elements increase. Since there are more filled energy levels, valence electrons are found farther from the nucleus. From the top, each successive element has a lower ionization energy because it is easier to remove an electron since the atoms are less tightly bound. Similarly, a group has a top-to-bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus.[20]There are exceptions to these trends: for example, in group 11, electronegativity increases farther down the group.[21]hidevteGroups in the Periodic tableIUPAC group123aa456789101112131415161718Mendeleev(I–VIII)IIIIIIIVVVIVIIVIIIIIIIIIIVVVIVIIbCAS(US, A-B-A)IAIIAIIIBIVBVBVIBVIIBVIIIBIBIIBIIIAIVAVAVIAVIIAVIIIAold IUPAC(Europe, A-B)IAIIAIIIAIVAVAVIAVIIAVIIIIBIIBIIIBIVBVBVIBVIIB0Trivial nameAlkali metalsAlkaline earth metalsrCoin­age metalsTrielsTetrelsPnicto­gensrChal­co­gensrHalo­gensrNoble gasesrName by elementrLith­ium groupBeryl­lium groupScan­dium groupTitan­ium groupVana­dium groupChro­mium groupMan­ga­nese groupIron groupCo­balt groupNickel groupCop­per groupZinc groupBoron groupCar­bon groupNitro­gen groupOxy­gen groupFluor­ine groupHelium orNeon groupPeriod 1H hHePeriod 2LiBeBCNOFNePeriod 3NaMgAlSiPSClArPeriod 4KCaScTiVCrMnFeCoNiCuZnGaGeAsSeBrKrPeriod 5RbSrYZrNbMoTcRuRhPdAgCdInSnSbTeIXePeriod 6CsBaLaCe–LuHfTaWReOsIrPtAuHgTlPbBiPoAtRnPeriod 7FrRaAcTh–LrRfDbSgBhHsMtDsRgCnNhFlMcLvTsOga 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, La and Ac, or (2) Sc, Y, Lu and Lr.[22]b Group 18, the noble gases, were not discovered at the time of Mendeleev's original table. Later (1902), Mendeleev accepted the evidence for their existence, and they could be placed in a new "group 0", consistently and without breaking the periodic table principle.r Group name as recommended by IUPAC.h Hydrogen (H), while placed in group 1, is not considered to be part of the alkali metals.PeriodsMain 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.[23]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.[24]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.[20]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.[25]BlocksMain article: Block (periodic table)Left to right: s-, f-, d-, p-block in the periodic tableSpecific 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.[26][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.[27]Metals, metalloids and nonmetalsMetals, metalloids, nonmetals, and elements with unknown chemical properties in the periodic table. Sources disagree on the classification of some of these elements.According to their shared physical and chemical properties, the elements can be classified into the major categories of metals, metalloids and nonmetals. Metals are generally shiny, highly conducting solids that form alloys with one another and salt-like ionic compounds with nonmetals (other than noble gases). A majority of nonmetals are coloured or colourless insulating gases; nonmetals that form compounds with other nonmetals feature covalent bonding. In between metals and nonmetals are metalloids, which have intermediate or mixed properties.[28]Metal and nonmetals can be further classified into subcategories that show a gradation from metallic to non-metallic properties, when going left to right in the rows. The metals may be subdivided into the highly reactive alkali metals, through the less reactive alkaline earth metals, lanthanides and actinides, via the archetypal transition metals, and ending in the physically and chemically weak post-transition metals. Nonmetals may be simply subdivided into the polyatomic nonmetals, being nearer to the metalloids and show some incipient metallic character; the essentially nonmetallic diatomic nonmetals, nonmetallic and the almost completely inert, monatomic noble gases. Specialized groupings such as refractory metals and noble metals, are examples of subsets of transition metals, also known[29]and occasionally denoted.[30]Placing elements into categories and subcategories based just on shared properties is imperfect. There is a large disparity of properties within each category with notable overlaps at the boundaries, as is the case with most classification schemes.[31]Beryllium, for example, is classified as an alkaline earth metal although its amphoteric chemistry and tendency to mostly form covalent compounds are both attributes of a chemically weak or post-transition metal. Radon is classified as a nonmetallic noble gas yet has some cationic chemistry that is characteristic of metals. Other classification schemes are possible such as the division of the elements into mineralogical occurrence categories, or crystalline structures. Categorizing the elements in this fashion dates back to at least 1869 when Hinrichs[32]wrote that simple boundary lines could be placed on the periodic table to show elements having shared properties, such as metals, nonmetals, or gaseous elements.Periodic trends and patternsMain article: Periodic trendsElectron configurationMain article: Electronic configurationApproximate order in which shells and subshells are arranged by increasing energy according to the Madelung ruleThe electron configuration or organisation of electrons orbiting neutral atoms shows a recurring pattern or periodicity. The electrons occupy a series of electron shells (numbered 1, 2, and so on). Each shell consists of one or more subshells (named s, p, d, f and g). As atomic number increases, electrons progressively fill these shells and subshells more or less according to the Madelung rule or energy ordering rule, as shown in the diagram. The electron configuration for neon, for example, is 1s22s22p6. With an atomic number of ten, neon has two electrons in the first shell, and eight electrons in the second shell; there are two electrons in the s subshell and six in the p subshell. In periodic table terms, the first time an electron occupies a new shell corresponds to the start of each new period, these positions being occupied by hydrogen and the alkali metals.[33][34]Periodic table trends (arrows show an increase)Since the properties of an element are mostly determined by its electron configuration, the properties of the elements likewise show recurring patterns or periodic behaviour, some examples of which are shown in the diagrams below for atomic radii, ionization energy and electron affinity. It is this periodicity of properties, manifestations of which were noticed well before the underlying theory was developed, that led to the establishment of the periodic law (the properties of the elements recur at varying intervals) and the formulation of the first periodic tables.[33][34]Atomic radiiMain article: Atomic radiusAtomic 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.[35]The electrons in the 4f-subshell, which is progressively filled across the lanthanide series, are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii that are smaller than would be expected and that are almost identical to the atomic radii of the elements immediately above them.[36]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.[37]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.[36]Ionization energyIonization energy: each period begins at a minimum for the alkali metals, and ends at a maximum for the noble gasesMain article: Ionization energyThe 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.[37]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.[37]ElectronegativityMain article: ElectronegativityGraph showing increasing electronegativity with growing number of selected groupsElectronegativity is the tendency of an atom to attract a shared pair of electrons.[38]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.[39]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.[21]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.[21]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.[40]Electron affinityMain article: Electron affinityDependence of electron affinity on atomic number.[41]Values generally increase across each period, culminating with the halogens before decreasing precipitously with the noble gases. Examples of localized peaks seen in hydrogen, the alkali metals and the group 11 elements are caused by a tendency to complete the s-shell (with the 6s shell of gold being further stabilized by relativistic effects and the presence of a filled 4f sub shell). Examples of localized troughs seen in the alkaline earth metals, and nitrogen, phosphorus, manganese and rhenium are caused by filled s-shells, or half-filled p- or d-shells.[42]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. Chlorinemost 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.[43]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.[43]A trend of decreasing electron affinity going down groups would be expected. The additional electron will be entering an orbital farther away from the nucleus. As such this electron would be less attracted to the nucleus and would release less energy when added. In going down a group, around one-third of elements are anomalous, with heavier elements having higher electron affinities than their next lighter congenors. Largely, this is due to the poor shielding by d and f electrons. A uniform decrease in electron affinity only applies to group 1 atoms.[44]Metallic characterThe lower the values of ionization energy, electronegativity and electron affinity, the more metalliccharacter the element has. Conversely, nonmetallic character increases with higher values of these properties.[45]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,[46]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.[47][48]Linking or bridging groupsSc, Y, La, Ac, Lu, Lr, Cu, Ag, Au, Zn, Cd, Hg, He, Ne, Ar, Kr, Xe, RnHydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAluminiumSiliconPhosphorusSulfurChlorineArgonPotassiumCalciumScandiumTitaniumVanadiumChromiumManganeseIronCobaltNickelCopperZincGalliumGermaniumArsenicSeleniumBromineKryptonRubidiumStrontiumYttriumZirconiumNiobiumMolybdenumTechnetiumRutheniumRhodiumPalladiumSilverCadmiumIndiumTinAntimonyTelluriumIodineXenonCaesiumBariumLanthanumCeriumPraseodymiumNeodymiumPromethiumSamariumEuropiumGadoliniumTerbiumDysprosiumHolmiumErbiumThuliumYtterbiumLutetiumHafniumTantalumTungstenRheniumOsmiumIridiumPlatinumGoldMercury (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumUraniumNeptuniumPlutoniumAmericiumCuriumBerkeliumCaliforniumEinsteiniumFermiumMendeleviumNobeliumLawrenciumRutherfordiumDubniumSeaborgiumBohriumHassiumMeitneriumDarmstadtiumRoentgeniumCoperniciumNihoniumFleroviumMoscoviumLivermoriumTennessineOganesson32-column periodic table showing, from left to right, the location of group 3; lutetium and lawrencium; groups 11–12; and the noble gasesFrom 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.[49]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[50]or, more generally, s block metals[51][52]but have some of the physical properties of d block transition metals.[53]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.[54][55]Lawrencium, as an analogue of lutetium, would presumably 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.[58]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.[59][60]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.[49]HistoryMain article: History of the periodic tableFirst systemization attemptsThe discovery of the elements mapped to significant periodic table development dates (pre-, per- and post-)In 1789, Antoine Lavoisier published a list of 33 chemical elements, grouping them into gases, metals, nonmetals, and earths.[61]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;[62]this became known as the Law of Triads.[63]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.[62]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.[64]This concept eventually became known as valency; different elements bond with different numbers of atoms.[65]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.[66]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.[67]Concurrently, English chemist William Odling published an arrangement of 57 elements, ordered on the basis of their atomic weights. With some irregularities and gaps, he noticed what appeared to be a periodicity of atomic weights among the elements and that this accorded with "their usually received groupings".[68]Odling alluded to the idea of a periodic law but did not pursue it.[69]He subsequently proposed (in 1870) a valence-based classification of the elements.[70]Newlands' periodic table, as presented to the Chemical Society in 1866, and based on the law of octavesEnglish 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.[71][72]This so termed Law of Octaves was ridiculed by Newlands' contemporaries, and the Chemical Society refused to publish his work.[73]Newlands was nonetheless able to draft a table of the elements and used it to predict the existence of missing elements, such as germanium.[74]The Chemical Society only acknowledged the significance of his discoveries five years after they credited Mendeleev.[75]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.[76][77]Mendeleev's tableDmitri Mendeleev, watercolour by Ilya RepinMendeleev's periodic table from his book An Attempt Towards a Chemical Conception of the EtherA version of Mendeleev's 1869 periodic table: An experiment on a system of elements based on their atomic weights and chemical similarities. This early arrangement presents the periods vertically, and the groups horizontally.Russian chemistry professor Dmitri Mendeleev and German chemist Julius Lothar Meyerindependently published their periodic tables in 1869 and 1870, respectively.[78]Mendeleev's table was his first published version; that of Meyer was an expanded version of his (Meyer's) table of 1864.[79]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.[80]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.[81]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.[82]The second decision was to occasionally ignore the order suggested by the atomic weights and switch adjacent elements, such as tellurium and iodine, to better classify them into chemical families.Mendeleev published in 1869, using atomic weight to organize the elements, information determinable to fair precision in his time. Atomic weight worked well enough to allow Mendeleev to accurately predict the properties of missing elements.Following the discovery, in 1911, by Ernest Rutherford of the atomic nucleus, it was proposed that the integer count of the nuclear charge is identical to the sequential place of each element in the periodic table. In 1913, Henry Moseley using X-ray spectroscopy confirmed this proposal experimentally. Moseley determined the value of the nuclear charge of each element, and showed that Mendeleev's ordering actually places the elements in sequential order by nuclear charge.[83]Nuclear charge is identical to proton count, and determines the value of the atomic number (Z) of each element. Using atomic number gives a definitive, integer-based sequence for the elements. Moseley predicted, in 1913, that the only elements still missing between aluminium (Z=13) and gold (Z=79) were Z = 43, 61, 72, and 75, all of which were later discovered. The atomic number is the absolute definition of an element, and gives a factual basis for the ordering of the periodic table. The periodic table is used to predict the properties of new synthetic elements before they are produced and studied.[84]Second version and further developmentMendeleev'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 2016In 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.[85]These gaps were subsequently filled as chemists discovered additional naturally occurring elements.[86]It is often stated that the last naturally occurring element to be discovered was francium (referred to by Mendeleev as eka-caesium) in 1939.[87]Plutonium, produced synthetically in 1940, was identified in trace quantities as a naturally occurring element in 1971.[88]The popular[89]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.[90][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.[91][92][93]With the development of modern quantum mechanical theories of electronconfigurations 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.[94]Glenn T. Seaborg, in 1945, suggested a new periodic table showing the actinides as belonging to a second f-block series.In 1945, Glenn Seaborg, an American scientist, made the suggestion that the actinide elements, like the lanthanides, were filling an f sub-level. Before this time the actinides were thought to be forming a fourth d-block row. Seaborg's colleagues advised him not to publish such a radical suggestion as it would most likely ruin his career. As Seaborg considered he did not then have a career to bring into disrepute, he published anyway. Seaborg's suggestion was found to be correct and he subsequently went on to win the 1951 Nobel Prize in chemistry for his work in synthesizing actinide elements.[95][96][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.[97]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.[98]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.[99]Different periodic tablesThe long- or 32-column tableThe periodic table in 32-column formatThe 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".[100]The relationship of the f-block to the other blocks of the periodic table also becomes easier to see.[101]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.[102]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),[103]and the familiarity of chemists with the modern form (as introduced by Seaborg).[104]Tables with different structuresMain article: Alternative periodic tablesWithin 100 years of the appearance of Mendeleev's table in 1869, Edward G. Mazurs had collected an estimated 700 different published versions of the periodic table.[105][106][107]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,[108]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.[107]Theodor Benfey's spiral periodic tableA popular[109]alternative structure is that of Otto Theodor Benfey (1960). The elements are arranged in a continuous spiral, with hydrogen at the centre and the transition metals, lanthanides, and actinides occupying peninsulas.[110]Most periodic tables are two-dimensional;[3]three-dimensional tables are known to as far back as at least 1862 (pre-dating Mendeleev's two-dimensional table of 1869). More recent examples include Courtines' Periodic Classification (1925),[111]Wringley's Lamina System (1949),[112]Giguère's Periodic helix (1965)[113]and Dufour's Periodic Tree (1996).[114]Going one further, Stowe's Physicist's Periodic Table (1989)[115]has been described as being four-dimensional (having three spatial dimensions and one colour dimension).[116]The various forms of periodic tables can be thought of as lying on a chemistry–physics continuum.[117]Towards the chemistry end of the continuum can be found, as an example, Rayner-Canham's "unruly"[118]Inorganic Chemist's Periodic Table (2002),[119]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.[120]A somewhat similar approach has been taken by Alper,[121]albeit criticized by Eric Scerri as disregarding the need to display chemical and physical periodicity.[122]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.[123]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.[69][124]hidevteLeft-step periodic table (by Charles Janet)f1f2f3f4f5f6f7f8f9f10f11f12f13f14d1d2d3d4d5d6d7d8d9d10p1p2p3p4p5p6s1s21sHHe2sLiBe2p 3sBCNOFNeNaMg3p 4sAlSiPSClArKCa3d 4p 5sScTiVCrMnFeCoNiCuZnGaGeAsSeBrKrRbSr4d 5p 6sYZrNbMoTcRuRhPdAgCdInSnSbTeIXeCsBa4f 5d 6p 7sLaCePrNdPmSmEuGdTbDyHoErTmYbLuHfTaWReOsIrPtAuHgTlPbBiPoAtRnFrRa5f 6d 7p 8sAcThPaUNpPuAmCmBkCfEsFmMdNoLrRfDbSgBhHsMtDsRgCnNhFlMcLvTsOg119120f-blockd-blockp-blocks-blockThis form of periodic table is more congruent with the order in which electron shells are ideally filled according to the Madelung rule, as shown in the accompanying sequence in the left margin (read from top to bottom, left to right). In reality, the filling of electron shells is characterized by a number of irregularities.Open questions and controversiesPlacement of hydrogen and heliumSimply following electron configurations, hydrogen (electronic configuration 1s1) and helium (1s2) should be placed in groups 1 and 2, above lithium (1s22s1) and beryllium (1s22s2).[26]While such a placement is common for hydrogen, it is rarely used for helium outside of the context of electron configurations: When the noble gases (then called "inert gases") were first discovered around 1900, they were known as "group 0", reflecting no chemical reactivity of these elements known at that point, and helium was placed on the top of that group, as it did share the extreme chemical inertness seen throughout the group. As the group changed its formal number, many authors continued to assign helium directly above neon, in group 18; one of the examples of such placing is the current IUPAC table.[125]Hydrogen's chemical properties are not very close to those of the alkali metals, which occupy group 1. On this basis it is sometimes placed elsewhere. A common alternative is at the top of group 17[122]given hydrogen's strictly univalent and largely non-metallic chemistry, and the strictly univalent and non-metallic chemistry of fluorine (the element otherwise at the top of group 17). Sometimes, to show hydrogen has properties corresponding to both those of the alkali metals and the halogens, it is shown at the top of the two columns simultaneously.[126]Another suggestion is above carbon in group 14: placed that way, it fits well into the trends of increasing ionization potential values and electron affinity values, and is not too far from the electronegativity trend, even though hydrogen cannot show the tetravalencecharacteristic of the heavier group 14 elements.[127]Finally, hydrogen is sometimes placed separately from any group; this is based on its general properties being different from those of the elements in any other group. The other period 1 element, helium, is sometimes placed separately from any group as well.[128]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)[129]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 7Although 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.[130][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".[132]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[133][134]but the majority of authors seem unconvinced.[135]Most working chemists are not aware there is any controversy.[136]In December 2015 an IUPAC project was established to make a recommendation on the matter.[137]Lanthanum and actiniumLa and Ac below YLanthanum and actinium are commonly depicted as the remaining group 3 members.[138][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]6s2and [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".[139]A consistent set of electron configurations is then seen in group 3: scandium [Ar]3d14s2, yttrium [Kr]4d15s2and lanthanum [Xe]5d16s2. Still in period 6, ytterbium was assigned an electron configuration of [Xe]4f135d16s2and 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".[139]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]5d16s2lanthanum, for the group 3 periodic table position below yttrium.[139]Lanthanum has the advantage of incumbency since the 5d1electron 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.[140]In terms of chemical behaviour,[141]and trends going down group 3 for properties such as melting point, electronegativity and ionic radius,[142][143]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.[144]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.[145]Lutetium and lawrenciumLu and Lr below YIn 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".[139]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.[139]This arrangement, in which lanthanum is the first member of the f-block, is disputed by some authors since lanthanum lacks any f-electrons. It has been argued that this is not valid concern given other periodic table anomalies—thorium, for example, has no f-electrons yet is part of the f-block.[146]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.[147][n 15]Chemically, scandium, yttrium and lutetium (and presumably lawrencium) behave like trivalent versions of the group 1–2 metals.[149]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.[139]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.[139]Lanthanides and actinidesMarkers below YA few authors position all thirty lanthanides and actinides in the two positions below yttrium (usually via footnote markers). This variant, which is stated in the 2005 Red Book to be the IUPAC-agreed version as of 2005 (a number of later versions exist, and the last update is from 1st Dec. 2018),[150][n 16]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 17]Groups included in the transition metalsThe 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.[152]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.[153]Other commentators, such as Jensen,[154]have argued that the formation of a compound like HgF4can 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.[154]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.[155]In this case, only groups 4–11 are regarded as transition metals. Though the group 3 elements show few of the characteristic chemical properties of the transition metals, they do show some of their characteristic physical properties (on account of the presence in each atom of a single d electron).[53]Elements with unknown chemical propertiesAlthough all elements up to oganesson have been discovered, of the elements above hassium (element 108), only copernicium (element 112), nihonium (element 113), and flerovium (element 114) have known chemical properties, and only for copernicium is there enough evidence for a conclusive categorisation at present. The other elements may behave differently from what would be predicted by extrapolation, due to relativistic effects; for example, flerovium has been predicted to possibly exhibit some noble-gas-like properties, even though it is currently placed in the carbon group.[156]The current experimental evidence still leaves open the question of whether flerovium behaves more like a metal or a noble gas.[157]Further periodic table extensionsMain article: Extended periodic tableHydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAluminiumSiliconPhosphorusSulfurChlorineArgonPotassiumCalciumScandiumTitaniumVanadiumChromiumManganeseIronCobaltNickelCopperZincGalliumGermaniumArsenicSeleniumBromineKryptonRubidiumStrontiumYttriumZirconiumNiobiumMolybdenumTechnetiumRutheniumRhodiumPalladiumSilverCadmiumIndiumTinAntimonyTelluriumIodineXenonCaesiumBariumLanthanumCeriumPraseodymiumNeodymiumPromethiumSamariumEuropiumGadoliniumTerbiumDysprosiumHolmiumErbiumThuliumYtterbiumLutetiumHafniumTantalumTungstenRheniumOsmiumIridiumPlatinumGoldMercury (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumUraniumNeptuniumPlutoniumAmericiumCuriumBerkeliumCaliforniumEinsteiniumFermiumMendeleviumNobeliumLawrenciumRutherfordiumDubniumSeaborgiumBohriumHassiumMeitneriumDarmstadtiumRoentgeniumCoperniciumNihoniumFleroviumMoscoviumLivermoriumTennessineOganessonUnunenniumUnbiniliumUnbiuniumUnquadquadiumUnquadpentiumUnquadhexiumUnquadseptiumUnquadoctiumUnquadenniumUnpentniliumUnpentuniumUnpentbiumUnpenttriumUnpentquadiumUnpentpentiumUnpenthexiumUnpentseptiumUnpentoctiumUnpentenniumUnhexniliumUnhexuniumUnhexbiumUnhextriumUnhexquadiumUnhexpentiumUnhexhexiumUnhexseptiumUnhexoctiumUnhexenniumUnseptniliumUnseptuniumUnseptbiumUnbibiumUnbitriumUnbiquadiumUnbipentiumUnbihexiumUnbiseptiumUnbioctiumUnbienniumUntriniliumUntriuniumUntribiumUntritriumUntriquadiumUntripentiumUntrihexiumUntriseptiumUntrioctiumUntrienniumUnquadniliumUnquaduniumUnquadbiumUnquadtriumPeriodic table with eight rows, extended to element 172[158]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.[159]More recently, physicists such as Pekka Pyykkö have theorized that these additional elements do not follow the Madelung rule, which predicts how electron shells are filled and thus affects the appearance of the present periodic table. There are currently several competing theoretical models for the placement of the elements of atomic number less than or equal to 172. In all of these it is element 172, rather than element 168, that emerges as the next noble gas after oganesson, although these must be regarded as speculative as no complete calculations have been done beyond element 122.[160]Element with the highest possible atomic numberThe number of possible elements is not known. A very early suggestion made by Elliot Adams in 1911, and based on the arrangement of elements in each horizontal periodic table row, was that elements of atomic weight greater than circa 256 (which would equate to between elements 99 and 100 in modern-day terms) did not exist.[161]A higher—more recent—estimate is that the periodic table may end soon after the island of stability,[162]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.[163]Other predictions of an end to the periodic table include at element 128 by John Emsley,[3]at element 137 by Richard Feynman,[164]and at element 155 by Albert Khazan.[3][n 18]Bohr modelThe 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.[165]Hence the non-relativistic Bohr model is inaccurate when applied to such an element.Relativistic Dirac equationThe 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.[166]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.[167]This does not happen if the innermost orbital is filled, so that element 173 is not necessarily the end of the periodic table.[164]Optimal formThe many different forms of periodic table have prompted the question of whether there is an optimal or definitive form of periodic table.[168]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.[