document_title: Periodic table

question_text: how many elements in the current periodic table

question_tokens: how many elements in the current periodic table

text: Periodic table

text: Overview

text: vte

text:
Group

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

Alkali metals

Alkaline earth metals

Pnictogens

Chalcogens

Halogens

Noble gases

Period
1

Hydrogen1H7000100800000000000♠1.008

Helium2He7000400260000000000♠4.0026

2

Lithium3Li7000694000000000000♠6.94

Beryllium4Be7000901220000000000♠9.0122

Boron5B7001108100000000000♠10.81

Carbon6C7001120110000000000♠12.011

Nitrogen7N7001140070000000000♠14.007

Oxygen8O7001159990000000000♠15.999

Fluorine9F7001189980000000000♠18.998

Neon10Ne7001201800000000000♠20.180

3

Sodium11Na7001229899999999999♠22.990

Magnesium12Mg7001243050000000000♠24.305

Aluminium13Al7001269820000000000♠26.982

Silicon14Si7001280850000000000♠28.085

Phosphorus15P7001309740000000000♠30.974

Sulfur16S7001320600000000000♠32.06

Chlorine17Cl7001354500000000000♠35.45

Argon18Ar7001399480000000000♠39.948

4

Potassium19K7001390980000000000♠39.098

Calcium20Ca7001400780000000000♠40.078

Scandium21Sc7001449560000000000♠44.956

Titanium22Ti7001478670000000000♠47.867

Vanadium23V7001509420000000000♠50.942

Chromium24Cr7001519960000000000♠51.996

Manganese25Mn7001549380000000000♠54.938

Iron26Fe7001558450000000000♠55.845

Cobalt27Co7001589330000000000♠58.933

Nickel28Ni7001586930000000000♠58.693

Copper29Cu7001635460000000000♠63.546

Zinc30Zn7001653800000000000♠65.38

Gallium31Ga7001697230000000000♠69.723

Germanium32Ge7001726300000000000♠72.630

Arsenic33As7001749220000000000♠74.922

Selenium34Se7001789710000000000♠78.971

Bromine35Br7001799040000000000♠79.904

Krypton36Kr7001837980000000000♠83.798

5

Rubidium37Rb7001854680000000000♠85.468

Strontium38Sr7001876200000000000♠87.62

Yttrium39Y7001889060000000000♠88.906

Zirconium40Zr7001912240000000000♠91.224

Niobium41Nb7001929060000000000♠92.906

Molybdenum42Mo7001959500000000000♠95.95

Technetium43Tc[98]

Ruthenium44Ru7002101070000000000♠101.07

Rhodium45Rh7002102910000000000♠102.91

Palladium46Pd7002106420000000000♠106.42

Silver47Ag7002107870000000000♠107.87

Cadmium48Cd7002112410000000000♠112.41

Indium49In7002114820000000000♠114.82

Tin50Sn7002118710000000000♠118.71

Antimony51Sb7002121760000000000♠121.76

Tellurium52Te7002127600000000000♠127.60

Iodine53I7002126900000000000♠126.90

Xenon54Xe7002131289999999999♠131.29

6

Caesium55Cs7002132910000000000♠132.91

Barium56Ba7002137330000000000♠137.33

Lanthanum57La7002138910000000000♠138.91

Hafnium72Hf7002178490000000000♠178.49

Tantalum73Ta7002180950000000000♠180.95

Tungsten74W7002183840000000000♠183.84

Rhenium75Re7002186210000000000♠186.21

Osmium76Os7002190230000000000♠190.23

Iridium77Ir7002192220000000000♠192.22

Platinum78Pt7002195080000000000♠195.08

Gold79Au7002196970000000000♠196.97

Mercury80Hg7002200590000000000♠200.59

Thallium81Tl7002204380000000000♠204.38

Lead82Pb7002207200000000000♠207.2

Bismuth83Bi7002208980000000000♠208.98

Polonium84Po[209]

Astatine85At[210]

Radon86Rn[222]

7

Francium87Fr[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]

Cerium58Ce7002140120000000000♠140.12

Praseodymium59Pr7002140910000000000♠140.91

Neodymium60Nd7002144240000000000♠144.24

Promethium61Pm[145]

Samarium62Sm7002150360000000000♠150.36

Europium63Eu7002151960000000000♠151.96

Gadolinium64Gd7002157250000000000♠157.25

Terbium65Tb7002158930000000000♠158.93

Dysprosium66Dy7002162500000000000♠162.50

Holmium67Ho7002164930000000000♠164.93

Erbium68Er7002167260000000000♠167.26

Thulium69Tm7002168930000000000♠168.93

Ytterbium70Yb7002173050000000000♠173.05

Lutetium71Lu7002174970000000000♠174.97

Thorium90Th7002232040000000000♠232.04

Protactinium91Pa7002231040000000000♠231.04

Uranium92U7002238030000000000♠238.03

Neptunium93Np[237]

Plutonium94Pu[244]

Americium95Am[243]

Curium96Cm[247]

Berkelium97Bk[247]

Californium98Cf[251]

Einsteinium99Es[252]

Fermium100Fm[257]

Mendelevium101Md[258]

Nobelium102No[259]

Lawrencium103Lr[266]

text: 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)

text:
Ca: 7001400780000000000♠40.078 — Formal short value, rounded (no uncertainty)[5]
Po: [209] — mass number of the most stable isotope

text: Background color shows subcategory in the metal–metalloid–nonmetal trend:

text:
Metal

Metalloid

Nonmetal

Unknownchemicalproperties

Alkali metal

Alkaline earth metal

Lanthanide

Actinide

Transition metal

Post-transition metal

Reactive nonmetal

Noble gas

text: 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.[6]

text: 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]

text: As of  2016[update], 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]

text: 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]

text: Grouping methods

text: Groups

text: 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]

text: 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]

text: 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.[18]

text: 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]

text: vte

text:
IUPAC group

1

2

3a

a

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

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 metalsr

Coinage metals

Triels

Tetrels

Pnictogensr

Chalcogensr

Halogensr

Noble gasesr

Name by elementr

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

h H h

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

text: Periods

text: 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]

text: 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]

text: Blocks

text: 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.[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]

text: Metals, metalloids and nonmetals

text: 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]

text: 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]

text: 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.

text: Periodic trends and patterns

text: Electron configuration

text: The electron configuration or organisation of electrons orbiting neutral atoms shows a recurring pattern or periodicity. The electrons occupy a series of electron shells (numbered 1, 2, and so on). Each shell consists of one or more subshells (named s, p, d, f and g). As atomic number increases, electrons progressively fill these shells and subshells more or less according to the Madelung rule or energy ordering rule, as shown in the diagram. The electron configuration for neon, for example, is 1s2 2s2 2p6. With an atomic number of ten, neon has two electrons in the first shell, and eight electrons in the second shell; there are two electrons in the s subshell and six in the p subshell. In periodic table terms, the first time an electron occupies a new shell corresponds to the start of each new period, these positions being occupied by hydrogen and the alkali metals.[33][34]

text: 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]

text: Atomic radii

text: 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]

text: 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]

text: Ionization energy

text: 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.[37]

text: 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]

text: Electronegativity

text: Electronegativity 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]

text: 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]

text: Electron affinity

text: 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.[43]

text: 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]

text: 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]

text: Metallic character

text: 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.[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]

text: Linking or bridging groups

text:
Hydrogen

Helium

Lithium

Beryllium

Boron

Carbon

Nitrogen

Oxygen

Fluorine

Neon

Sodium

Magnesium

Aluminium

Silicon

Phosphorus

Sulfur

Chlorine

Argon

Potassium

Calcium

Scandium

Titanium

Vanadium

Chromium

Manganese

Iron

Cobalt

Nickel

Copper

Zinc

Gallium

Germanium

Arsenic

Selenium

Bromine

Krypton

Rubidium

Strontium

Yttrium

Zirconium

Niobium

Molybdenum

Technetium

Ruthenium

Rhodium

Palladium

Silver

Cadmium

Indium

Tin

Antimony

Tellurium

Iodine

Xenon

Caesium

Barium

Lanthanum

Cerium

Praseodymium

Neodymium

Promethium

Samarium

Europium

Gadolinium

Terbium

Dysprosium

Holmium

Erbium

Thulium

Ytterbium

Lutetium

Hafnium

Tantalum

Tungsten

Rhenium

Osmium

Iridium

Platinum

Gold

Mercury (element)

Thallium

Lead

Bismuth

Polonium

Astatine

Radon

Francium

Radium

Actinium

Thorium

Protactinium

Uranium

Neptunium

Plutonium

Americium

Curium

Berkelium

Californium

Einsteinium

Fermium

Mendelevium

Nobelium

Lawrencium

Rutherfordium

Dubnium

Seaborgium

Bohrium

Hassium

Meitnerium

Darmstadtium

Roentgenium

Copernicium

Nihonium

Flerovium

Moscovium

Livermorium

Tennessine

Oganesson

text: 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.[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]

text: History

text: First systemization attempts

text: 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]

text: 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]

text: 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]

text: 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]

text: 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.[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]

text: 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]

text: Mendeleev's table

text: Russian chemistry professor Dmitri Mendeleev and German chemist Julius Lothar Meyer independently 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]

text: 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.

text: 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.

text: 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]

text: Second version and further development

text: 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.[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]

text: 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]

text: 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.[94]

text: 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]

text: 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]

text: Different periodic tables

text: The long- or 32-column table

text: 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".[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]

text: Tables with different structures

text: Within 100 years of the appearance of Mendeleev's table in 1869, Edward G. Mazurs had collected an estimated 700 different published versions of the periodic table.[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]

text: A 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]

text:
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]

text: 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]

text: vte

text:
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

text: Open questions and controversies

text: Placement of hydrogen and helium

text: Simply following electron configurations, hydrogen (electronic configuration 1s1) and helium (1s2) should be placed in groups 1 and 2, above lithium (1s22s1) and beryllium (1s22s2).[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]

text: 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 tetravalence characteristic 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.

text: Group 3 and its elements in periods 6 and 7

text: 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.[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.

text: 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]

text: Lanthanum and actinium

text: La and Ac below Y

text: Lanthanum 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]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".[139] 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".[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]5d16s2 lanthanum, for the group 3 periodic table position below yttrium.[139] Lanthanum has the advantage of incumbency since the 5d1 electron appears for the first time in its structure whereas it appears for the third time in lutetium, having also made a brief second appearance in gadolinium.[140]

text: 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]

text: Lutetium and lawrencium

text: Lu and Lr below Y

text: 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".[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]

text: 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]

text: Lanthanides and actinides

text: Markers below Y

text:
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]

text: Groups included in the transition metals

text: 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.[150] 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.

text: 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.[151] Other commentators, such as Jensen,[152] 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.[152]

text: 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.[153] 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]

text: Elements with unknown chemical properties

text: Although all elements up to oganesson have been discovered, of the elements above hassium (element 108), only copernicium (element 112), nihonium (element 113), and flerovium (element 114) have known chemical properties, and only for copernicium is there enough evidence for a conclusive categorisation at present. The other elements may behave differently from what would be predicted by extrapolation, due to relativistic effects; for example, flerovium has been predicted to possibly exhibit some noble-gas-like properties, even though it is currently placed in the carbon group.[154] The current experimental evidence still leaves open the question of whether flerovium behaves more like a metal or a noble gas.[155]

text: Further periodic table extensions

text:
Ununennium

Unbinilium

Unbiunium

Unquadquadium

Unquadpentium

Unquadhexium

Unquadseptium

Unquadoctium

Unquadennium

Unpentnilium

Unpentunium

Unpentbium

Unpenttrium

Unpentquadium

Unpentpentium

Unpenthexium

Unpentseptium

Unpentoctium

Unpentennium

Unhexnilium

Unhexunium

Unhexbium

Unhextrium

Unhexquadium

Unhexpentium

Unhexhexium

Unhexseptium

Unhexoctium

Unhexennium

Unseptnilium

Unseptunium

Unseptbium

Unbibium

Unbitrium

Unbiquadium

Unbipentium

Unbihexium

Unbiseptium

Unbioctium

Unbiennium

Untrinilium

Untriunium

Untribium

Untritrium

Untriquadium

Untripentium

Untrihexium

Untriseptium

Untrioctium

Untriennium

Unquadnilium

Unquadunium

Unquadbium

Unquadtrium

text: 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.[157] 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.[158]

text: Element with the highest possible atomic number

text: The number of possible elements is not known. A very early suggestion made by Elliot Adams in 1911, and based on the arrangement of elements in each horizontal periodic table row, was that elements of atomic weight greater than circa 256 (which would equate to between elements 99 and 100 in modern-day terms) did not exist.[159] A higher—more recent—estimate is that the periodic table may end soon after the island of stability,[160] 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.[161] Other predictions of an end to the periodic table include at element 128 by John Emsley,[3] at element 137 by Richard Feynman,[162] and at element 155 by Albert Khazan.[3][n 17]

text: 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.[163] Hence the non-relativistic Bohr model is inaccurate when applied to such an element.

text: 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.[164] 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.[165] This does not happen if the innermost orbital is filled, so that element 173 is not necessarily the end of the periodic table.[162]

text: Optimal form

text: The many different forms of periodic table have prompted the question of whether there is an optimal or definitive form of periodic table.[166] 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]


document_title: List of National Rugby League records

question_text: who has won the most nrl premierships as a player

root

text: List of National Rugby League records

text: Team records[edit]

text: Premierships[edit]

text:
Titles
Club
Seasons

21
 South Sydney Rabbitohs
1908, 1909, 1914, 1918, 1925, 1926, 1927, 1928, 1929, 1931, 1932, 1950, 1951, 1953, 1954, 1955, 1967, 1968, 1970, 1971, 2014

15
 St. George Dragons
1941, 1949, 1956, 1957, 1958, 1959, 1960, 1961, 1962, 1963, 1964, 1965, 1966, 1977, 1979

13
 Sydney Roosters
1911, 1912, 1913, 1923, 1935, 1936, 1937, 1940, 1945, 1974, 1975, 2002, 2013

11
 Balmain Tigers
1915, 1916, 1917, 1919, 1920, 1924, 1939, 1944, 1946, 1947, 1969

8
 Canterbury-Bankstown Bulldogs
1938, 1942, 1980, 1984, 1985, 1988, 1995, 2004

8
 Manly-Warringah Sea Eagles
1972, 1973, 1976, 1978, 1987, 1996, 2008, 2011

6
 Brisbane Broncos
1992, 1993, 1997SL, 1998, 2000, 2006

4
 Western Suburbs Magpies
1930, 1934, 1948, 1952

4
 Parramatta Eels
1981, 1982, 1983, 1986

3
 Newtown Jets
1910, 1933, 1943

3
 Canberra Raiders
1989, 1990, 1994

2
 North Sydney Bears
1921, 1922

2
 Newcastle Knights
1997ARL, 2001

2
 Penrith Panthers
1991, 2003

2
 Melbourne Storm
1999, 2007, 2009, 2012

1
 Wests Tigers
2005

1
 St. George Illawarra Dragons
2010

1
 North Queensland Cowboys
2015

1
 Cronulla-Sutherland Sharks
2016

tag: table

text: Most consecutive[edit]

text: 11 -  St. George (1956 to 1966)

tag: ul

text: Runners Up[edit]

text:
Tally
Club
Seasons

15
 Sydney Roosters
1908, 1919, 1921, 1928, 1931, 1934, 1938, 1941, 1960, 1972, 1980, 2000, 2003, 2004, 2010

13
 South Sydney Rabbitohs
1910, 1916, 1917, 1920, 1923, 1924, 1935, 1937, 1939, 1949, 1952, 1965, 1969

12
 St. George Dragons
1927, 1930, 1933, 1942, 1946, 1953, 1971, 1975, 1985, 1992, 1993, 1996

11
 Manly-Warringah Sea Eagles
1951, 1957, 1959, 1968, 1970, 1982, 1983, 1995, 1997ARL, 2007, 2013

10
 Canterbury-Bankstown Bulldogs
1940, 1947, 1967, 1974, 1979, 1986, 1994, 1998, 2012, 2014

9
 Balmain Tigers
1909, 1936, 1945, 1948, 1956, 1964, 1966, 1988, 1989

8
 Western Suburbs Magpies
1918, 1925, 1932, 1950, 1958, 1961, 1962, 1963

7
 Newtown Jets
1913, 1914, 1929, 1944, 1954, 1955, 1981

5
 Parramatta Eels
1976, 1977, 1984, 2001, 2009

4
 Glebe
1911, 1912, 1915, 1922

3
 Cronulla-Sutherland Sharks
1973, 1978, 1997SL

3
 Melbourne Storm
2006, 2008, 2016

2
 Canberra Raiders
1987, 1991

2
 New Zealand Warriors
2002, 2011

1
 University
1926

1
 North Sydney Bears
1943

1
 Penrith Panthers
1990

1
 St. George Illawarra Dragons
1999

1
 North Queensland Cowboys
2005

1
 Brisbane Broncos
2015

text: Most consecutive[edit]

text: 3 -  Western Suburbs Magpies (1961 to 1963)

text: Minor Premierships[edit]

text:
Tally
Club
Seasons

19
 Sydney Roosters
1912, 1913, 1923, 1931, 1934, 1935, 1936, 1937, 1940, 1941, 1945, 1974, 1975, 1980, 1981, 2004, 2013, 2014, 2015

17
 South Sydney Rabbitohs
1908, 1909, 1914, 1918, 1925, 1926, 1927, 1929, 1932, 1949, 1950, 1951, 1953, 1968, 1969, 1970, 1989

15
 St. George Dragons
1928, 1946, 1956, 1957, 1958, 1959, 1960, 1962, 1963, 1964, 1965, 1966, 1967, 1979, 1985

9
 Manly-Warringah Sea Eagles
1971, 1972, 1973, 1976, 1983, 1987, 1995, 1996, 1997ARL

7
 Balmain Tigers
1915, 1916, 1917, 1919, 1920, 1924, 1939

7
 Canterbury-Bankstown Bulldogs
1938, 1942, 1947, 1984, 1993, 1994, 2012

6
 Newtown Jets
1910, 1933, 1943, 1944, 1954, 1955

5
 Western Suburbs Magpies
1930, 1948, 1951, 1961, 1978

5
 Parramatta Eels
1977, 1982, 1986, 2001, 2005

4
 Brisbane Broncos
1992, 1997SL, 1998, 2000

3
 Melbourne Storm
2006, 2007, 2008, 2011, 2016, 2017

2
 North Sydney Bears
1921, 1922

2
 Cronulla-Sutherland Sharks
1988, 1999

2
 Penrith Panthers
1991, 2003

2
 St. George Illawarra Dragons
2009, 2010

1
 Glebe
1911

1
 Canberra Raiders
1990

1
 New Zealand Warriors
2002

text: Most consecutive[edit]

text: 6 -  St. George (1962 to 1967)

text: Wooden spoons[edit]

text:
Tally
Club
Seasons

17
 Western Suburbs Magpies
1909, 1910, 1912, 1913, 1916, 1933, 1940, 1942, 1953, 1955, 1971, 1983, 1984, 1987, 1988, 1998, 1999

13
 Parramatta Eels
1947, 1952, 1954, 1956, 1957, 1958, 1959, 1960, 1961, 1970, 1972, 2012, 2013

10
 University
1921, 1923, 1927, 1929, 1930, 1931, 1934, 1935, 1936, 1937

9
 North Sydney Bears
1915, 1917, 1919, 1932, 1941, 1948, 1950, 1951, 1979

8
 Newtown Jets
1924, 1925, 1928, 1939, 1968, 1976, 1977, 1978

8
 South Sydney Rabbitohs
1945, 1946, 1962, 1975, 1990, 2003, 2004, 2006

5
 Canterbury-Bankstown Bulldogs
1943, 1944, 1964, 2002, 2008

5
 Sydney Roosters
1949, 1963, 1965, 1966, 2009

4
 Balmain Tigers
1911, 1974, 1981, 1994

4
 Penrith Panthers
1973, 1980, 2001, 2007

4
 Newcastle Knights
2005, 2015, 2016, 2017

3
 Annandale
1914, 1918, 1920

3
 St. George Dragons
1922, 1926, 1938

3
 Illawarra Steelers
1985, 1986, 1989

3
 Gold Coast Seagulls
1991, 1992, 1993

3
 North Queensland Cowboys
1995, 1997SL, 2000

3
 Cronulla-Sutherland Sharks
1967, 1969, 2014

2
 South Queensland Crushers
1996, 1997ARL

1
 Cumberland
1908

1
 Canberra Raiders
1982

1
 Melbourne Storm
2010

1
 Gold Coast Titans
2011

text: Most consecutive[edit]

text: 6 -  Parramatta (1956 to 1961)

tag: p

text: Team wins, losses and draws[edit]

text: Matches played[edit]

text:
Rank
Team
1st Match
Matches
Won
Drawn
Lost
PF
PA
 % Won

1
 Melbourne Storm
14 March 1998
473
296
6
171
11,232
8,251
63.21

2
 Brisbane Broncos
6 March 1988
705
435
14
257
16,258
12,160
62.62

3
 St. George Dragons
23 April 1921
1,545
910
56
579
27,452
21,729
60.71

4
 Manly-Warringah Sea Eagles
12 April 1947
1,513
876
36
601
29,426
23,934
59.08

5
 Glebe
20 April 1908
297
163
6
128
3,966
3,365
55.89

6
 Sydney Roosters
20 April 1908
2,126
1,121
69
936
37,426
33,315
54.35

7
 Canterbury-Bankstown Bulldogs
25 April 1935
1,749
910
53
786
30,347
29,773
53.54

8
 Balmain Tigers
20 April 1908
1,705
871
68
766
26,688
25,609
53.08

9
 South Sydney Rabbitohs
20 April 1908
2,054
1,063
46
945
34,892
34,947
52.87

10
 St. George Illawarra Dragons
6 March 1999
435
225
6
204
9,428
8,387
52.41

11
 Canberra Raiders
27 February 1982
843
425
9
409
18,137
17,220
50.94

12
 Newcastle Knights
5 March 1988
683
340
15
328
14,156
14,033
50.87

13
 Cronulla-Sutherland Sharks
2 April 1967
1,170
554
23
593
20,922
21,384
48.33

14
 New Zealand Warriors
10 March 1995
514
238
7
269
10,868
10,395
46.98

15
 Newtown Jets
20 April 1908
1,305
583
59
663
18,836
20,231
46.93

16
 Parramatta Eels
12 April 1947
1,560
707
40
813
26,801
28,556
46.60

17
 Newcastle (1908–1909)
20 April 1908
20
9
0
11
295
229
45.00

18
 Western Suburbs Magpies
20 April 1908
1,691
734
49
908
25,244
28,961
44.86

19
 Wests Tigers
6 February 2000
397
174
3
220
8,346
9,400
44.20

20
 Gold Coast Titans
18 March 2007
220
97
0
123
4,110
5,023
44.09

21
 Penrith Panthers
2 April 1967
1,152
482
28
642
20,716
23,317
43.05

22
 North Sydney Bears
20 April 1908
1,665
678
71
916
25,133
27,837
42.85

23
 North Queensland Cowboys
11 March 1995
517
213
6
298
10,573
12,257
41.77

24
 Illawarra Steelers
28 February 1982
396
153
13
230
6,614
7,313
40.28

25
 Western Reds
12 March 1995
61
24
1
36
995
1,425
40.16

26
 Northern Eagles
6 February 2000
76
30
1
45
1,582
2,118
40.13

27
 Hunter Mariners
2 March 1997
18
7
0
11
350
363
38.89

28
 Adelaide Rams
1 March 1997
42
13
1
28
696
1,017
32.15

29
 Gold Coast Chargers
5 March 1988
246
53
9
184
3,285
5,854
23.37

30
 South Queensland Crushers
11 March 1995
65
13
1
51
844
1,628
20.77

31
 University
8 May 1920
242
47
5
190
2,392
5,308
20.45

32
 Annandale
30 April 1910
153
25
6
122
1,065
2,583
18.30

33
 Cumberland
9 May 1908
8
1
0
7
38
191
12.50

text: Source: Rugby League Tables & Statistics. Last updated: 18 January 2016.

text: Most consecutive wins[edit]

text:
Rank
Wins
Team
Period

1
19 matches
 Eastern Suburbs
from Manly-Warringah, 13 April 1975 to Newtown, 23 August 1975

2
17 matches
 Canterbury-Bankstown Bulldogs
from St. George Illawarra, 31 March 2002 to North Queensland, 3 August 2002

=3
16 matches
 South Sydney
from Newtown, 30 May 1908 to Glebe, 17 July 1909

=3
16 matches
 Eastern Suburbs
from North Sydney, 15 June 1912 to Glebe, 28 June 1913

=5
15 matches
 South Sydney
from Balmain, 2 May 1925 to North Sydney, 8 May 1926

=5
15 matches
 Manly-Warringah Sea Eagles
from South Sydney, 12 March 1995 to Parramatta, 9 July 1995

=5
15 matches
 Melbourne Storm
from Penrith Panthers, 4 August 2012 to New Zealand Warriors, 25 April 2013

=5
15 matches
 Cronulla-Sutherland Sharks
from Melbourne Storm, 28 March 2016 to Newcastle Knights, 24 July 2016

Source: Rugby League Tables & Statistics. Last updated: 1 August 2016.

text: Most consecutive losses[edit]

text:
Rank
Losses
Team
Period

1
42 matches
 University
from Balmain, 28 April 1934 to Eastern Suburbs, 22 August 1936

2
25 matches
 Eastern Suburbs
from North Sydney, 25 July 1965 to Canterbury-Bankstown, 9 April 1967

=3
22 matches
 South Sydney
from St. George, 23 June 1945 to St. George, 12 April 1947

=3
22 matches
 Western Suburbs Magpies
from Eastern Suburbs Roosters, 25 April 1984 to St. George Dragons, 14 April 1985

5
20 matches
 Newtown Jets
from Penrith Panthers, 27 March 1977 to Manly-Warringah Sea Eagles, 14 August 1977

6
19 matches
 Newcastle Knights
from Brisbane Broncos, 16 April 2016 to New Zealand Warriors, 5 March 2017

Source: Rugby League Tables & Statistics. Last updated: 11 March 2017.

text: Result records[edit]

text: Greatest win margins[edit]

text:
Rank
Margin
Winning Team
Defeated Team
Score
Venue
Date

1
85
 St. George
 Canterbury-Bankstown
91–6
Earl Park
11 May 1935

2
80
 Eastern Suburbs
 Canterbury-Bankstown
87–7
Sydney Sports Ground
18 May 1935

3
70
 Parramatta Eels
 Cronulla-Sutherland Sharks
74–4
Parramatta Stadium
23 August 2003

4
68
 Canberra Raiders
 Parramatta Eels
68–0
Bruce Stadium
22 August 1993

5
67
 South Sydney
 Western Suburbs
67–0
Royal Agricultural Society Grounds
21 July 1910

Source: Rugby League Tables & Statistics. Last updated: 28 May 2007.

text: Most points in a game[edit]

text:
Rank
Points
Winning Team
Defeated Team
Score
Venue
Date

1
102
 Newcastle Knights
 Canberra Raiders
70–32
Canberra Stadium
18 March 2006

2
97
 St. George
 Canterbury-Bankstown
91–6
Earl Park
11 May 1935

3
96
 Penrith Panthers
 Northern Eagles
68–28
Brookvale Oval
8 September 2002

=4
94
 Eastern Suburbs
 Canterbury-Bankstown
87–7
Sydney Sports Ground
18 May 1935

=4
94
 Sydney Roosters
 North Queensland Cowboys
64–30
Sydney Football Stadium
9 June 2007

Source: Rugby League Tables & Statistics. Last updated: 28 May 2007.

text: Fewest points in a game[edit]

text:
Rank
Points
Winning Team
Defeated Team
Score
Venue
Date

1
0
 Newtown Jets drew with  Canterbury-Bankstown Bulldogs
–
0–0
Henson Park
28 May 1982

2
1
 Newtown
 St. George
1–0
Sydney Cricket Ground
12 May 1973

=3
2
 Newtown
 Eastern Suburbs
2–0
Sydney Cricket Ground
16 May 1914

=3
2
 Balmain
 North Sydney
2–0
North Sydney Oval
25 April 1931

=3
2
 South Sydney
 Parramatta
2–0
Sydney Sports Ground
18 July 1965

Source: Rugby League Tables & Statistics. Last updated: 8 February 2007.

text: Most Points Scored by One Team in a Game[edit]

text:
Rank
Points
Winning Team
Defeated Team
Score
Venue
Date

1
91
 St. George
 Canterbury-Bankstown
91–6
Earl Park
11 May 1935

2
87
 Eastern Suburbs
 Canterbury-Bankstown
87–7
Sydney Sports Ground
18 May 1935

=3
74
 Parramatta Eels
 Cronulla-Sutherland Sharks
74–4
Parramatta Stadium
23 August 2003

=3
74
 Canberra Raiders
 Penrith Panthers
74–12
Canberra Stadium
10 August 2008

5
72
 Penrith Panthers
 Manly-Warringah Sea Eagles
72–12
Penrith Stadium
7 August 2004

text: Highest Scores by a Losing Side[edit]

text:
Points
Winning Team
Defeated Team
Score
Venue
Date

36
 Canterbury-Bankstown Bulldogs
 Hunter Mariners
48–36
Belmore Sports Ground
12 May 1997

36
 Wests Tigers
 Canterbury-Bankstown Bulldogs
37–36
Telstra Stadium
27 March 2005

36
 Gold Coast Titans
 Newcastle Knights
38–36
Energy Australia Stadium
9 May 2010

36
 Gold Coast Titans
 Melbourne Storm
38–36
Suncorp Stadium
13 May 2017

34
 St George Illawarra Dragons
 Manly-Warringah Sea Eagles
36–34
OKI Jubilee Stadium
29 August 2004

34
 Canterbury-Bankstown Bulldogs
 Canberra Raiders
41–34
GIO Stadium
24 May 2015

34
 Parramatta Eels
 Brisbane Broncos
52–34
Suncorp Stadium
24 August 2017

text: Highest Drawn Games[edit]

text:
Score
Teams
Venue
Date

34-34
 Parramatta vs.  North Sydney
Parramatta Oval
25 June 1949

34–34
 Illawarra Steelers vs.  Manly-Warringah Sea Eagles
WIN Stadium
4 May 1997

34–34
 South Sydney Rabbitohs vs.  Brisbane Broncos
Aussie Stadium
28 August 2004

32–32
 Penrith Panthers vs.  New Zealand Warriors
CUA Stadium
1 August 2009

31–31
 Sydney Roosters vs.  New Zealand Warriors
Sydney Football Stadium
5 August 2007

Source: Rugby League Tables & Statistics. Last updated: 8 February 2007.

text: Best Start to a season[edit]

text:
Wins
Team
Season

15
 Manly-Warringah Sea Eagles
1995 ARL season

12
 South Sydney
1925 NSWRFL season

11
 Western Suburbs
1934 NSWRFL season

11
 Balmain
1966 NSWRFL season

10
 St. George
1959 NSWRFL season

10
 Sydney City Roosters
1996 ARL season

text: Undefeated in a season[edit]

text:
Year
Team

1915
 Balmain

1921
 North Sydney

1925
 South Sydney

1936
 Eastern Suburbs

1937
 Eastern Suburbs

1959
 St. George

text: Failed To Win a Match in a season[edit]

text:
Year
Team

1918
 Annandale

1920
 Annandale

1921
 University

1935
 University

1937
 University

1946
 South Sydney

1966
 Eastern Suburbs

text: Most Points Scored by a Team in a season[edit]

text: 839 points by the  Parramatta Eels in 2001.

text: Most Tries and Goals by a Team in a Grand Final[edit]

text: 8 tries, 7 goals by the  Eastern Suburbs Roosters in 1975.

text: Largest winning margin in a Grand Final[edit]

text: 40 points to nil by the  Manly Warringah Sea Eagles over the  Melbourne Storm in 2008.

text: Fewest Points Scored by a Team in a season[edit]

text: 38 points by  Cumberland in 1908.

text: Most Points Conceded by a Team in a season[edit]

text: 944 points by the  Western Suburbs Magpies in 1999.

text: Fewest points conceded by a team in a season[edit]

text: 41 points by  South Sydney in 1909.

text: Keeping opposition scoreless: most games in a season[edit]

text: 6 matches by the  Sydney Roosters in 2013.

text: Biggest comebacks[edit]

text:
26 Points:  North Queensland Cowboys came from 26–0 down at halftime to win 36–28 against the  Penrith Panthers on 29 May 1998 at Penrith Stadium
26 Points:  Penrith Panthers came from 32–6 down after 53 minutes to draw 32–32 against the  New Zealand Warriors on 1 August 2009 at CUA Stadium
24 Points:  Wests Tigers came from 24–0 down after 26 minutes to win 36–32 against the  Newcastle Knights on 11 August 2001 at Campbelltown Stadium
24 Points:  St. George Illawarra Dragons came from 34–10 down after 57 minutes to win 36–34 against the  Manly-Warringah Sea Eagles on 29 August 2004 at Oki Jubilee Stadium
24 Points:  South Sydney Rabbitohs came from 28–4 down after 52 minutes to win 29–28 against the  North Queensland Cowboys on 28 June 2008 at Dairy Farmers Stadium
24 Points:  Cronulla-Sutherland Sharks came from 24–0 down after 31 minutes to win 30–28 against the  Sydney Roosters on 5 July 2014 at Allianz Stadium
24 Points:  North Queensland Cowboys came from 30-6 down after 58 minutes to win 36-30 against the  Parramatta Eels on 8 June 2015 at Pirtek Stadium
23 Points:  Penrith Panthers came from 31–8 down after 53 minutes to win 32–31 against the  Wests Tigers on 4 June 2000 at Penrith Stadium.
22 Points:  Melbourne Storm came from 22-0 down after 20 minutes to win 36-32 against the  Cronulla Sharks on 16 March 2003 at Southern Cross Group Stadium.
22 Points:  Parramatta Eels came from 22–0 down at halftime to win 34–28 against the  Penrith Panthers on 17 July 2010 at CUA Stadium.
22 Points:  Cronulla-Sutherland Sharks came from 22-0 down after 43 minutes to win 24-22 against the  Brisbane Broncos on 28 June 2014 at Suncorp Stadium.
22 Points:  Canberra Raiders came from 22-0 down after 28 minutes to win 30-22 against the  Wests Tigers on 19 April 2015 at Leichhardt Oval.
22 Points:  Canberra Raiders came from 22-0 down after 39 minutes to win 29-25 against the  Newcastle Knights on 3 July 2016 at GIO Stadium Canberra.
22 Points:  Penrith Panthers came from 28-6 down after 46 minutes to win 36-28 against the  New Zealand Warriors on 13 May 2017 at Pepper Stadium.

text: Individual records[edit]

text: Note: Figures in boldface are currently playing in the NRL.

text: Most games (NRL's '300 Club')[edit]

text:
Games
Player
Clubs
Years

356
Cameron Smith
 Melbourne
2002–

355
Darren Lockyer
 Brisbane
1995–2011

350
Terry Lamb
 Western Suburbs,  Canterbury-Bankstown
1980–1996

349
Steve Menzies
 Manly-Warringah,  Northern Eagles
1993–2008

347
Corey Parker
 Brisbane
2001–2016

336
Brad Fittler
 Penrith,  Sydney Roosters
1989–2004

332
Cliff Lyons
 North Sydney,  Manly-Warringah
1985–1999

330
Nathan Hindmarsh
 Parramatta
1998–2012

328
Andrew Ettingshausen
 Cronulla-Sutherland
1983–2000

325
Geoff Gerard
 Parramatta,  Manly-Warringah,  Penrith
1974–1989

321
Cooper Cronk
 Melbourne
2004–

318
Jason Croker
 Canberra
1991–2006

317
Hazem El Masri
 Canterbury-Bankstown
1996–2009

317
Chris Heighington
 Wests Tigers,  Cronulla-Sutherland
2003–

315
Paul Langmack
 Canterbury-Bankstown,  Western Suburbs,  Sydney Roosters
1983–1999

315
Luke Priddis
 Canberra,  Brisbane,  Penrith,  St. George Illawarra
1997–2010

313
Steve Price
 Canterbury-Bankstown,  New Zealand Warriors
1994–2009

313
Brent Kite
 St. George Illawarra,  Manly-Warringah,  Penrith
2002–2015

312
Ruben Wiki
 Canberra,  New Zealand Warriors
1993–2008

309
Petero Civoniceva
 Brisbane,  Penrith
1998–2012

307
Brett Kimmorley
 Newcastle,  Hunter,  Melbourne,  Northern Eagles,  Cronulla-Sutherland,  Canterbury-Bankstown
1995–2010

307
Luke Lewis
 Penrith,  Cronulla-Sutherland
2001–

305
Ryan Hoffman
 Melbourne,  New Zealand Warriors
2003–

304
Paul Gallen
 Cronulla-Sutherland
2001–

303
Anthony Watmough
 Northern Eagles,  Manly-Warringah,  Parramatta
2002–2015

302
Anthony Minichiello
 Sydney Roosters
2000–2014

301
Luke Ricketson
 Sydney Roosters
1991–2005

300
Scott Prince
 North Queensland,  Brisbane,  Wests Tigers,  Gold Coast Titans
1998–2013

300
John Morris
 Newcastle,  Parramatta,  Wests Tigers,  Cronulla-Sutherland
2001–2014

Source:http://www.nrl.com/cameron-smiths-300th-game-a-rugby-league-icon/tabid/10874/newsid/88068/default.aspx The NRL'S 300 club. As of 10 September 2017.

text: At one club[edit]

text:
Games
Player
Clubs
Years

356
Cameron Smith
 Melbourne
2002–

355
Darren Lockyer
 Brisbane
1995–2011

347
Corey Parker
 Brisbane
2001–2016

330
Nathan Hindmarsh
 Parramatta
1998–2012

328
Andrew Ettingshausen
 Cronulla-Sutherland
1983–2000

321
Cooper Cronk
 Melbourne
2004–

318
Jason Croker
 Canberra
1991–2006

317
Hazem El Masri
 Canterbury-Bankstown
1996–2009

309
Cliff Lyons
 Manly-Warringah
1986–1999

304
Paul Gallen
 Cronulla-Sutherland
2001–

302
Anthony Minichiello
 Sydney Roosters
2000–2014

301
Luke Ricketson
 Sydney Roosters
1991–2005

Source: Eels NRL Player Stats. Retrieved 3 February 2012.

text: Most tries[edit]

text: In a career[edit]

text:
Rank
Tries
Player
Clubs

1
212
Ken Irvine
 North Sydney,  Manly-Warringah

=2
180
Steve Menzies
 Manly-Warringah,  Northern Eagles

=2
180
Billy Slater
 Melbourne

4
165
Andrew Ettingshausen
 Cronulla-Sutherland

5
164
Terry Lamb
 Western Suburbs,  Canterbury-Bankstown

6
163
Brett Stewart
 Manly-Warringah

=7
159
Matt Sing
 Penrith,  Sydney Roosters,  North Queensland

=7
159
Hazem El Masri
 Canterbury-Bankstown

9
154
Nathan Merritt
 South Sydney,  Cronulla-Sutherland

=10
152
Harold Horder
 South Sydney,  North Sydney

=10
152
Manu Vatuvei
 New Zealand Warriors

Source: NRL Stats. Last updated: 9 September 2017.

text: In a season[edit]

text:
Tries
Player
Club
Season

38
Dave Brown[1]
 Eastern Suburbs
1935 NSWRFL season

34
Ray Preston
 Newtown
1954 NSWRFL season

29
Les Brennan
 South Sydney
1954 NSWRFL season

28
Bobby Lulham
 Balmain
1947 NSWRFL season

28
Johnny Graves
 South Sydney
1951 NSWRFL season

text: In a game[edit]

text: 8 - Frank Burge,  Glebe vs.  University, Round 7, 1920

text: Most points[edit]

text: In a career[edit]

text:
Rank
Points
Player
Clubs
Period

1
2,418
Hazem El Masri
 Canterbury-Bankstown
316 games (1996–2009)

2
2,176
Andrew Johns
 Newcastle
249 games (1993–2007)

3
2,163
Cameron Smith
 Melbourne
356 games (2002–)

4
2,107
Jason Taylor
 Western Suburbs,  North Sydney,  Northern Eagles,  Parramatta
276 games (1990–2001)

5
2,056
Johnathan Thurston
 Canterbury-Bankstown,  North Queensland
299 games (2002–)

6
2,034
Daryl Halligan
 North Sydney,  Canterbury-Bankstown
230 games (1991–2000)

7
1,971
Mick Cronin
 Parramatta
216 games (1977–1986)

8
1,917
Graham Eadie
 Manly-Warringah
233 games (1971–1983)

9
1,841
Eric Simms
 South Sydney
206 games (1965–1975)

10
1,793
Luke Burt
 Parramatta
264 games (1999–2012)

Source: Rugby League Tables & Statistics. Last updated: 9 September 2017.

text: In a season[edit]

text:
Points
Player
Club
Year

342 (16t, 139g)
Hazem El Masri
 Canterbury-Bankstown
2004 NRL season

308 (15t, 124g)
Brett Hodgson
 Wests Tigers
2005 NRL season

296 (17t, 114g)
Hazem El Masri
 Canterbury-Bankstown
2006 NRL season

296 (18t, 112g)
Jarrod Croker
 Canberra
2016 NRL season

294 (9t, 129g)
Hazem El Masri
 Canterbury-Bankstown
2003 NRL season

text: In a game[edit]

text: 45 - Dave Brown,[1]  Eastern Suburbs vs.  Canterbury-Bankstown, Round 6, 1935

text: Oldest Players[edit]

text:
Age
Player
Club
Year

40 years and 5 days
Billy Wilson
 North Sydney
1967 NSWRFL season

39 years and 311 days
Ted Courtney
 Western Suburbs
1924 NSWRFL season

38 years and 149 days
Roy Kirkaldy
 Canterbury-Bankstown
1948 NSWRFL season

38 years and 28 days
Henry Porter
 Canterbury-Bankstown
1948 NSWRFL season

38 years and 25 days
Sandy Pearce
 Eastern Suburbs
1921 NSWRFL season

text: Youngest Players[edit]

text:
Age
Player
Club
Year

16 years and 85 days
Ray Stehr
 Eastern Suburbs
1929 NSWRFL season

16 years and 220 days
Jack Arnold
 Western Suburbs
1936 NSWRFL season

16 years and 238 days
Jordan Rankin
 Gold Coast Titans
2008 NRL season

16 years and 258 days
Frank Burge
 Glebe
1911 NSWRFL season

16 years and 303 days
Adam Ritson
 Cronulla-Sutherland
1993 NSWRL season