دوري جدول

**دوري جدول** (Periodic Table) جنهن کي عنصرن جي دوري جدول پڻ سڏيو ويندو آهي، ڪيميائي تتن کي قطارن، جنهن کي دور (periods) چئبو آهي ۽ ڪالمن، جنهن کي گروپ (groups) چئبو آهي، ۾ ڏيکاريل ترتيب آهي. اها علم ڪيميا جي هڪ آئڪن آهي ۽ طبيعيات ۽ ٻين سائنسن ۾ وڏي پيماني تي استعمال ٿيندي آهي. اها دوري قانون جي هڪ شڪل آهي، جيڪي بيان ڪري ٿي ته جڏهن عنصرن کي انهن جي ايٽمي انگن جي ترتيب سان ترتيب ڏنو ويندو آهي، انهن جي خاصيتن جي ورهاست تقريبن واضح ٿي ويندي آهي. جدول کي چار مستطيل علائقن ۾ ورهايو ويو آهي، جنهن کي بلاڪ (block) سڏيو ويندو آهي. ساڳئي ڪالم يا گروپ جا عنصر ساڳيون ڪيميائي خاصيتون ڏيکاريندا آهن. 

عنصرن جا عمودي (vertical)، افقي (horizontal) ۽ ترچھا (diagonal) رجحان دوري جدول کي ممتاز ڪن ٿا. ڌاتو ڪردار وارا (metallic) تت هڪ گروپ هيٺ ۽ کاٻي کان     ساڄي طرف هڪ دور ۾ وڌندا آهن. غير ڌاتو ڪردار وارا (non metallic) تت دوري جدول جي گروپ ۾ مٿيان کاٻي کان هيٺيان ساڄي طرف وڌندا آهن. پهرين دوري جدول، جن کي قبوليت ملي ۽ عام طور تي قبول ڪئي وئي آهي، اها 1869ع ۾ روسي ڪيميادان، ديمتري مينڊيليف جي آهي. هن جدول کي دوري قانون جي ايٽمي مايو (Atomic masses) تي ڪيميائي خاصيتن جي انحصار جي طور تي تيار ڪيو، جيئن ته ان وقت سڀئي ڪيميائي تت معلوم نه هئا، هن جي دوري جدول ۾ خال هئا ۽ مينڊيليف ڪاميابي سان دوري قانون کي استعمال ڪندي ڪجهه غائب عنصرن جي ڪجهه خاصيتن جي اڳڪٿي ڪئي. 

19هين صدي جي آخر ۾ دوري قانون کي هڪ بنيادي دريافت طور تسليم ڪيو ويو ۽ ان جي وضاحت 20هين صدي جي شروعات ۾ ايٽمي انگن جي دريافت ۽ ڪوانٽم ميڪانيات، ٻئي ايٽم جي اندروني جوڙجڪ کي روشن ڪرڻ لاءِ ڪم ڪن ٿا، ۾ لاڳاپيل اڳڀرائي جي ڪم سان ڪئي وئي. جدول جي هڪ سڃاڻپ جديد شڪل، سال **1945**ع ۾ گلين ٽي. سيبورگ جي دريافت ته اڪٽينائيڊ حقيقت ۾ ڊي-بلاڪ (d-block) عنصرن جي بدران ايف-بلاڪ (f-block) جا عنصر آهن، سان مڪمل ٿي. دوري جدول ۽ دوري قانون هاڻي جديد علم ڪيميا جو هڪ مرڪزي ۽ لازمي حصو آهن.

دوري جدول جو ارتقا، سائنس جي ترقي سان گڏ جاري آهي. فطرت ۾، صرف ايٽمي نمبر **94** تائين جا عنصر موجود آهن. اڳتي وڌڻ لاءِ، ليبارٽري ۾ نوان عنصر ترڪيب ڪرڻ ضروري آهي. سال **2010**ع تائين، پهرين **118** عنصر جي سڃاڻپ ٿي وئي، جڏهن غير موجود عنصرن کي ليبارٽري ۾ ترڪيب ڪيو ويو ۽ جدول جي پهريون ست قطارون (periods) مڪمل ٿي ويون، جڏهن ته، وڌيڪ ڳري عنصرن جي ڪيميائي خاصيتن جي تصديق ڪرڻ لاءِ، اڃا تائين ضرورت آهي، جئين ته انهن جون خاصيتون، دوري جدول ۾ انهن جي جاء سان ملنديون آهن. وڌيڪ ڳري عنصرن جي نيون دريافتون جدول کي انهن ستن قطارن کان اڳتي وڌائينديون. جئين ته اهو اڃا تائين معلوم ناهي ته ڪيترا وڌيڪ عنصر جو وجود ممڪن آهي، پر وڌيڪ، نظرياتي حساب اها اندازو ڏين ٿا ته عنصرن جو هي نامعلوم علائقو جدول جي ڄاتل سڃاتل حصي جي نمونن جي پيروي نه ڪندو. ڪجهه سائنسي بحث پڻ جاري آهي ته ڇا ڪجهه عنصر موجود جدول ۾ صحيح جاء تي آهن. دوري قانون جي ڪيتريون متبادل جدول موجود آهن ۽ بحث آهي ته ڇا موجود شڪل دوري جدول جي هڪ بهترين شڪل آهي.

## جوڙجڪ<span class="anchor" id="Detailed table"></span>

**نوٽ**: بارڊر عنصر جي قدرتي موجودگي کي ڏيکاري ٿي.

_________________**<small>شروع کان (</small><small>ابتدائي)</small>****______________________**...............................**زوال جي پيداوار** **.............................**------------------------**<small>مصنوعي</small> <small>طور</small>**

--------------------------
------------------------------**[[معياري ايٽمي مايو|معياري ايٽمي مايو (A<sub>r</sub>)]]:**

*** [[ڪيلشيئم|ڪيلشيم]] (Ca): 40.078 3 a.m.u<br>* [[پولونيئم|پولونيم]] (Po): [209] (سڀ کان وڌيڪ مستحڪم آئسوٽوپ جو ماس نمبر (A)**

هر ڪيميائي عنصر جو هڪ منفرد ايٽمي نمبر "Z"  (جرمن: Zahl، مطلب، انگ) هوندو آهي، جيڪو ان جي مرڪز (nucleus) پروٽانن جي تعداد جي نمائندگي ڪري ٿو. تنهن ڪري هر هڪ الڳ ايٽمي نمبر ايٽم جي هڪ طبقي سان ملندو آهي: انهن طبقن کي ڪيميائي عنصر سڏيو ويندو آهي. ڪيميائي عنصر اهي آهن جيڪي دوري جدول ۾ درجه بندي سان ترتيب ڏنل آهن. هائيڊروجن اهو عنصر آهي جنهن جو ايٽمي نمبر "1" آهي؛ اھڙي طرح هيليم جو "2" ۽ لٿيم جو "3" آهي. انهن مان هر هڪ نالا هڪ يا ٻن اکرن واري ڪيميائي علامت سان وڌيڪ مختصر ڪري سگهجي ٿو؛ هائيڊروجن، هيليم ۽ ليٿيم لاءِ ترتيب وار H، He، ۽ Li آهن. نيوٽران ايٽم جي ڪيميائي سڃاڻپ کي متاثر نه ڪندا آهن، پر ان جي وزن کي متاثر ڪندا آهن. ايٽم جن ۾ پروٽانن جي تعداد ساڳي هوندي آهي، پر نيوٽران جي تعداد مختلف هوندي آهي، انهن کي هڪ ئي ڪيميائي عنصر جا آئسوٽوپ سڏيو ويندو آهي. قدرتي طور تي پيدا ٿيندڙ عنصر عام طور تي مختلف آئسوٽوپن جي ميلاپ جي طور تي موجود آهن؛ ڇاڪاڻ ته هر آئسوٽوپ عام طور تي هڪ خاصيت جي ڪثرت سان ٿئي ٿو، قدرتي طور تي پيدا ٿيندڙ عنصرن جا چڱي طرح بيان ڪيل ايٽمي وزن هوندا آهن، جيڪي ان عنصر جي قدرتي طور تي پيدا ٿيندڙ ايٽم جي سراسري ماين جي طور تي بيان ڪيا ويندا آهن. سڀني عنصرن ۾ ڪيترائي آئسوٽوپس هوندا آهن، مختلف قسمن ۾ پروٽانن جو تعداد ساڳيو هوندو آهي پر نيوٽرانن جو تعداد مختلف هوندو آهي. مثال طور، ڪاربان ۾ ٽي قدرتي طور تي پيدا ٿيندڙ آئسوٽوپس هوندا آهن: ان جي سڀني ايٽمن ۾ ڇهه پروٽان هوندا آهن ۽ لڳ ڀڳ ۾ ڇهه نيوٽران هوندا آهن، پر هڪ سيڪڙي ۾ ست ۽ هڪ تمام ننڍڙي حصي ۾ اٺ به هوندا آهن. آئسوٽوپس ڪڏهن به دوري جدول ۾ الڳ سان نه ڏيکاريا ويندا آهن؛ ۽ هميشه هڪ تت طور سراسري وزن سان ڏيکاريا ويندا آھن، اهي هميشه هڪ عنصر جي تحت گڏ ڪيا ويندا آهن، جهڙوڪ ڪاربان (C6)12.001

**ذيلي مدار:**
Each chemical element has a unique atomic number (*Z* for "Zahl", German for "number") representing the number of protons in its nucleus. Each distinct atomic number therefore corresponds to a class of atom: these classes are called the chemical elements. The chemical elements are what the periodic table classifies and organizes. Hydrogen is the element with atomic number 1; helium, atomic number 2; lithium, atomic number 3; and so on. Each of these names can be further abbreviated by a one- or two-letter chemical symbol; those for hydrogen, helium, and lithium are respectively H, He, and Li. Neutrons do not affect the atom's chemical identity, but do affect its weight. Atoms with the same number of protons but different numbers of neutrons are called isotopes of the same chemical element. Naturally occurring elements usually occur as mixes of different isotopes; since each isotope usually occurs with a characteristic abundance, naturally occurring elements have well-defined atomic weights, defined as the average mass of a naturally occurring atom of that element.
All elements have multiple isotopes, variants with the same number of protons but different numbers of neutrons. 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. When atomic mass is shown, it is usually the weighted average of naturally occurring isotopes; but if no isotopes occur naturally in significant quantities, the mass of the most stable isotope usually appears, often in parentheses.
### گروپن جا نالا ۽ نمبر
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). The f-block groups are ignored in this numbering. Groups can also be named by their first element, e.g. the "scandium group" for group 3. Previously, groups were known by Roman numerals. In the United States, 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 for groups 1 through 7, and "B" was used for groups 11 through 17. 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 (International Union of Pure and Applied Chemistry) naming system (1–18) was put into use, and the old group names (I–VIII) were deprecated.

### پريزنٽيشن جون قسمون
Both forms represent the same periodic table. The form with the f-block included in the main body is sometimes called the 32-column or long form; the form with the f-block cut out the 18-column or medium-long form. The 32-column form has the advantage of showing all elements in their correct sequence, but it has the disadvantage of requiring more space. The form chosen is an editorial choice, and does not imply any change of scientific claim or statement. For example, when discussing the composition of group 3, the options can be shown equally (unprejudiced) in both forms.

Periodic tables usually at least show the elements' symbols; many also provide supplementary information about the elements, either via colour-coding or as data in the cells. The above table shows the names and atomic numbers of the elements, and also their blocks, natural occurrences and standard atomic weights. For the short-lived elements without standard atomic weights, the mass number of the most stable known isotope is used instead. Other tables may include properties such as state of matter, melting and boiling points, densities, as well as provide different classifications of the elements.
### اليڪٽراني تشڪيل

The periodic table is a graphic description of the periodic law, which states that the properties and atomic structures of the chemical elements are a periodic function of their atomic number. Elements are placed in the periodic table according to their electron configurations, the periodic recurrences of which explain the trends in properties across the periodic table.

An electron can be thought of as inhabiting an atomic orbital, which characterizes the probability it can be found in any particular region around the atom. Their energies are quantised, which is to say that they can only take discrete values. Furthermore, electrons obey the Pauli exclusion principle: different electrons must always be in different states. This allows classification of the possible states an electron can take in various energy levels known as shells, divided into individual subshells, which each contain one or more orbitals. Each orbital can contain up to two electrons: they are distinguished by a quantity known as spin, conventionally labelled "up" or "down". In a cold atom (one in its ground state), electrons arrange themselves in such a way that the total energy they have is minimized by occupying the lowest-energy orbitals available. Only the outermost electrons (so-called valence electrons) have enough energy to break free of the nucleus and participate in chemical reactions with other atoms. The others are called core electrons.

Elements are known with up to the first seven shells occupied. The first shell contains only one orbital, a spherical s orbital. As it is in the first shell, this is called the 1s orbital. This can hold up to two electrons. The second shell similarly contains a 2s orbital, and it also contains three dumbbell-shaped 2p orbitals, and can thus fill up to eight electrons (2×1 + 2×3 = 8). The third shell contains one 3s orbital, three 3p orbitals, and five 3d orbitals, and thus has a capacity of 2×1 + 2×3 + 2×5 = 18. The fourth shell contains one 4s orbital, three 4p orbitals, five 4d orbitals, and seven 4f orbitals, thus leading to a capacity of 2×1 + 2×3 + 2×5 + 2×7 = 32. Higher shells contain more types of orbitals that continue the pattern, but such types of orbitals are not filled in the ground states of known elements. The subshell types are characterized by the quantum numbers. Four numbers describe an orbital in an atom completely: the principal quantum number *n*, the azimuthal quantum number ℓ (the orbital type), the orbital magnetic quantum number *m*ℓ, and the spin magnetic quantum number *m<sub>s</sub>*.

#### ذيلي مدارن لاء آفبائو جو اصول

The sequence in which the subshells are filled is given in most cases by the Aufbau principle, also known as the Madelung or Klechkovsky rule (after Erwin Madelung and Vsevolod Klechkovsky respectively). This rule was first observed empirically by Madelung, and Klechkovsky and later authors gave it theoretical justification. The shells overlap in energies, and the Madelung rule specifies the sequence of filling according to:
:1s ≪ 2s 1, where the superscript indicates the number of electrons in the subshell. Helium adds a second electron, which also goes into 1s, completely filling the first shell and giving the configuration 1s2.

Starting from the third element, lithium, the first shell is full, so its third electron occupies a 2s orbital, giving a 1s2 2s1 configuration. The 2s electron is lithium's only valence electron, as the 1s subshell is now too tightly bound to the nucleus to participate in chemical bonding to other atoms: such a shell is called a "core shell". The 1s subshell is a core shell for all elements from lithium onward. The 2s subshell is completed by the next element beryllium (1s2 2s2). The following elements then proceed to fill the 2p subshell. Boron (1s2 2s2 2p1) puts its new electron in a 2p orbital; carbon (1s2 2s2 2p2) fills a second 2p orbital; and with nitrogen (1s2 2s2 2p3) all three 2p orbitals become singly occupied. This is consistent with Hund's rule, which states that atoms usually prefer to singly occupy each orbital of the same type before filling them with the second electron. Oxygen (1s2 2s2 2p4), fluorine (1s2 2s2 2p5), and neon (1s2 2s2 2p6) then complete the already singly filled 2p orbitals; the last of these fills the second shell completely.

Starting from element 11, sodium, the second shell is full, making the second shell a core shell for this and all heavier elements. The eleventh electron begins the filling of the third shell by occupying a 3s orbital, giving a configuration of 1s2 2s2 2p6 3s1 for sodium. This configuration is abbreviated [Ne] 3s1, where [Ne] represents neon's configuration. Magnesium ([Ne] 3s2) finishes this 3s orbital, and the following six elements aluminium, silicon, phosphorus, sulfur, chlorine, and argon fill the three 3p orbitals ([Ne] 3s2 3p1 through [Ne] 3s2 3p6). This creates an analogous series in which the outer shell structures of sodium through argon are analogous to those of lithium through neon, and is the basis for the periodicity of chemical properties that the periodic table illustrates: at regular but changing intervals of atomic numbers, the properties of the chemical elements approximately repeat.

The first 18 elements can thus be arranged as the start of a periodic table. Elements in the same column have the same number of valence electrons and have analogous valence electron configurations: these columns are called groups. The single exception is helium, which has two valence electrons like beryllium and magnesium, but is typically placed in the column of neon and argon to emphasise that its outer shell is full. (Some contemporary authors question even this single exception, preferring to consistently follow the valence configurations and place helium over beryllium.) There are eight columns in this periodic table fragment, corresponding to at most eight outer-shell electrons. A period begins when a new shell starts filling. Finally, the colouring illustrates the blocks: the elements in the s-block (coloured red) are filling s&nbsp;orbitals, while those in the p-block (coloured yellow) are filling p&nbsp;orbitals.

Starting the next row, for potassium and calcium the 4s subshell is the lowest in energy, and therefore they fill it. Potassium adds one electron to the 4s shell ([Ar] 4s1), and calcium then completes it ([Ar] 4s2). However, starting from scandium ([Ar] 3d1 4s2) the 3d subshell becomes the next highest in energy. The 4s and 3d subshells have approximately the same energy and they compete for filling the electrons, and so the occupation is not quite consistently filling the 3d orbitals one at a time. The precise energy ordering of 3d and 4s changes along the row, and also changes depending on how many electrons are removed from the atom. For example, due to the repulsion between the 3d electrons and the 4s ones, at chromium the 4s energy level becomes slightly higher than 3d, and so it becomes more profitable for a chromium atom to have a [Ar] 3d5 4s1 configuration than an [Ar] 3d4 4s2 one. A similar anomaly occurs at copper, whose atom has a [Ar] 3d10 4s1 configuration rather than the expected [Ar] 3d9 4s2. These are violations of the Madelung rule. Such anomalies, however, do not have any chemical significance: most chemistry is not about isolated gaseous atoms, and the various configurations are so close in energy to each other that the presence of a nearby atom can shift the balance. Therefore, the periodic table ignores them and considers only idealized configurations.

At zinc ([Ar] 3d10 4s2), the 3d orbitals are completely filled with a total of ten electrons. Next come the 4p orbitals, completing the row, which are filled progressively by gallium ([Ar] 3d10 4s2 4p1) through krypton ([Ar] 3d10 4s2 4p6), in a manner analogous to the previous p-block elements. From gallium onwards, the 3d orbitals form part of the electronic core, and no longer participate in chemistry. The s- and p-block elements, which fill their outer shells, are called main-group elements; the d-block elements (coloured blue below), which fill an inner shell, are called transition elements (or transition metals, since they are all metals).

The next 18 elements fill the 5s orbitals (rubidium and strontium), then 4d (yttrium through cadmium, again with a few anomalies along the way), and then 5p (indium through xenon). Again, from indium onward the 4d orbitals are in the core. Hence the fifth row has the same structure as the fourth.

The sixth row of the table likewise starts with two s-block elements: caesium and barium. After this, the first f-block elements (coloured green below) begin to appear, starting with lanthanum. These are sometimes termed inner transition elements. As there are now not only 4f but also 5d and 6s subshells at similar energies, competition occurs once again with many irregular configurations; this resulted in some dispute about where exactly the f-block is supposed to begin, but most who study the matter agree that it starts at lanthanum in accordance with the Aufbau principle. Even though lanthanum does not itself fill the 4f subshell as a single atom, because of repulsion between electrons, its 4f orbitals are low enough in energy to participate in chemistry. At ytterbium, the seven 4f orbitals are completely filled with fourteen electrons; thereafter, a series of ten transition elements (lutetium through mercury) follows, and finally six main-group elements (thallium through radon) complete the period. From lutetium onwards the 4f orbitals are in the core, and from thallium onwards so are the 5d orbitals.

The seventh row is analogous to the sixth row: 7s fills (francium and radium), then 5f (actinium to nobelium), then 6d (lawrencium to copernicium), and finally 7p (nihonium to oganesson). Starting from lawrencium the 5f orbitals are in the core, and probably the 6d orbitals join the core starting from nihonium. Again there are a few anomalies along the way: for example, as single atoms neither actinium nor thorium actually fills the 5f subshell, and lawrencium does not fill the 6d shell, but all these subshells can still become filled in chemical environments. For a very long time, the seventh row was incomplete as most of its elements do not occur in nature. The missing elements beyond uranium started to be synthesized in the laboratory in 1940, when neptunium was made. (However, the first element to be discovered by synthesis rather than in nature was technetium in 1937.) The row was completed with the synthesis of tennessine in 2010 (the last element oganesson had already been made in 2002), and the last elements in this seventh row were given names in 2016.



This completes the modern periodic table, with all seven rows completely filled to capacity.

### Electron configuration table
The following table shows the electron configuration of a neutral gas-phase atom of each element. Different configurations can be favoured in different chemical environments. The main-group elements have entirely regular electron configurations; the transition and inner transition elements show twenty irregularities due to the aforementioned competition between subshells close in energy level. For the last ten elements (109–118), experimental data is lacking and therefore calculated configurations have been shown instead. Completely filled subshells have been greyed out.

## تبديليون
## دوري رجحان
## عنصرن جي درجه بندي
## تاريخ
## ستين دور کان اڳ مستقبل ۾ واڌارو
## متبادل دوري جدولون
## پڻ ڏسو
* نيوڪليوسنٿيسس 
## لنڪس
**دوري جدول** (Periodic Table) جنهن کي عنصرن جي دوري جدول پڻ چيو ويندو آهي، سا ڪيميائي عنصرن جي قطارن (Periods) ۽ ڪالمن (Groups) ۾ هڪ جدولي پيشڪش آهي. اهو ڪيميا جو هڪ آئڪن آهي ۽ وڏي پيماني تي فزڪس ۽ ٻين سائنسن ۾ استعمال ٿيندو آهي. قطار ۾ عنصر انهن جي ايٽمي انگن جي ترتيب سان ترتيب ڏنل آهن ۽ ڪالمن ۾ انهن جي خاصيتون ٻيهر ورجائي ظاهر ٿئين تيون. ٽيبل کي چار لڳ ڀڳ مستطيل علائقن ۾ ورهايو ويو آهي، جنهن کي "بلاڪ" سڏجي ٿو. ساڳئي گروپ ۾ عناصر ساڳيون ڪيميائي خاصيتون ڏيکاريندا آهن. ھي جدول ائٽمي نمبر، اليڪٽراني جوڙجڪ ۽ ورجندڙ ڪيميائي خاصيتن جي ڄاڻ ڏيندي آهي.

وقتي ٽيبل هڪ ترتيب ڏنل ترتيب آهي ڪيميائي عناصر کي قطارن (Periods) ۽ ڪالمن (Groups). اهو ڪيميا جو هڪ آئڪن آهي ۽ وڏي پيماني تي فزڪس ۽ ٻين سائنسن ۾ استعمال ٿيندو آهي. اهو دور جي قانون جي هڪ تصوير آهي، جنهن ۾ چيو ويو آهي ته جڏهن عناصر انهن جي ايٽمي انگن جي ترتيب سان ترتيب ڏنل آهن، انهن جي ملڪيت جي تقريبن ٻيهر ورجائي ظاهر ٿئي ٿي.
 ٽيبل کي چار لڳ ڀڳ مستطيل علائقن ۾ ورهايو ويو آهي، جنهن کي "بلاڪ" سڏجي ٿو. ساڳئي گروپ ۾ عناصر ساڳيون ڪيميائي خاصيتون ڏيکاريندا آهن. عمودي، افقي ۽ ويڪرائي رجحانات دورانياتي جدول جي خصوصيت ڪن ٿا. دھاتي ڪردار هڪ گروپ جي هيٺان ۽ هڪ دور ۾ ساڄي کان کاٻي طرف وڌندو آهي. نان ميٽالڪ ڪردار وڌندو آهي دوراني جدول جي هيٺان کاٻي پاسي کان مٿي ساڄي طرف. پھرين دوري جدول جيڪا عام طور تي قبول ڪئي وئي، جيڪا 1869ع ۾ روسي ڪيمسٽ دمتري مينڊيليف جي ھئي؛ هن ايٽمي ماس تي ڪيميائي ملڪيتن جي انحصار جي طور تي دورياتي قانون ٺاهيو. جيئن ته ان وقت سڀ عنصر سڃاتل نه هئا، ان ڪري هن جي دور جي جدول ۾ خال هئا، ۽ مينڊيليف ڪاميابيءَ سان دورياتي قانون کي استعمال ڪيو ته جيئن ڪجهه غائب عنصرن جي ملڪيتن جي اڳڪٿي ڪئي وڃي. 19 صدي عيسويء جي آخر ۾ دورياتي قانون کي هڪ بنيادي دريافت طور تسليم ڪيو ويو. ان جي وضاحت 20 صدي جي شروعات ۾ ڪئي وئي، ايٽمي انگن جي دريافت ۽ ڪوانٽم ميڪانڪس ۾ لاڳاپيل اڳڀرائي واري ڪم سان، ٻئي نظريا ايٽم جي اندروني ڍانچي کي روشن ڪرڻ لاءِ ڪم ڪن ٿا. 1945ع ۾ گلين ٽي. سيبورگ جي دريافت سان جدول جي جديد شڪل کي تسليم ڪيو ويو ته ايڪٽائنائيڊ حقيقت ۾ ڊي-بلاڪ عنصرن جي بجاءِ ايف-بلاڪ هئا. دوري جدول ۽ قانون هاڻي جديد ڪيميا جو مرڪزي ۽ لازمي حصو آهن.

### پريزينٽيشن

Both forms represent the same periodic table. The form with the f-block included in the main body is sometimes called the 32-column or long form; the form with the f-block cut out the 18-column or medium-long form. The 32-column form has the advantage of showing all elements in their correct sequence, but it has the disadvantage of requiring more space. The form chosen is an editorial choice, and does not imply any change of scientific claim or statement. For example, when discussing the composition of group 3, the options can be shown equally (unprejudiced) in both forms.

Periodic tables usually at least show the elements' symbols; many also provide supplementary information about the elements, either via colour-coding or as data in the cells. The above table shows the names and atomic numbers of the elements, and also their blocks, natural occurrences and standard atomic weights. For the short-lived elements without standard atomic weights, the mass number of the most stable known isotope is used instead. Other tables may include properties such as state of matter, melting and boiling points, densities, as well as provide different classifications of the elements.

## حوالا

زمرو:ڪيميا
زمرو:دوري جدول
زمرو:ڪيميائي درجه بندي