Ni Atomic Number

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Atomic Number of Nickel is 28. Chemical symbol for Nickel is Ni. Number of protons in Nickel is 28. Atomic weight of Nickel is 58.6934 u or g/mol. Melting point of Nickel is 1453 °C and its the boiling point is 2732 °C. Here is a list of the elements sorted by atomic number.Element nameElement symbolAtomic numberHydrogenH1HeliumHe2LithiumLi3. Symbol: Ni Atomic Number: 28 Atomic Mass: 58.6934 amu Melting Point: 1453.0 °C (1726.15 K, 2647.4 °F) Boiling Point: 2732.0 °C (3005.15 K, 4949.6 °F) Number of Protons/Electrons: 28 Number of Neutrons: 31 Classification: Transition Metal Crystal Structure: Cubic Density @ 293 K: 8.902 g/cm 3 Color: white Atomic Structure.

  1. Ni Ka Atomic Number
  2. Nickel Atom Symbol
  3. Mass Of Nickel
  4. Ni Atomic Number
An explanation of the superscripts and subscripts seen in atomic number notation. Atomic number is the number of protons, and therefore also the total positive charge, in the atomic nucleus.
The Rutherford–Bohr model of the hydrogen atom (Z = 1) or a hydrogen-like ion (Z > 1). In this model it is an essential feature that the photon energy (or frequency) of the electromagnetic radiation emitted (shown) when an electron jumps from one orbital to another be proportional to the mathematical square of atomic charge (Z2). Experimental measurement by Henry Moseley of this radiation for many elements (from Z = 13 to 92) showed the results as predicted by Bohr. Both the concept of atomic number and the Bohr model were thereby given scientific credence.

The atomic number or proton number (symbol Z) of a chemical element is the number of protons found in the nucleus of every atom of that element. The atomic number uniquely identifies a chemical element. It is identical to the charge number of the nucleus. In an uncharged atom, the atomic number is also equal to the number of electrons.

The sum of the atomic number Z and the number of neutronsN gives the mass numberA of an atom. Since protons and neutrons have approximately the same mass (and the mass of the electrons is negligible for many purposes) and the mass defect of nucleon binding is always small compared to the nucleon mass, the atomic mass of any atom, when expressed in unified atomic mass units (making a quantity called the 'relative isotopic mass'), is within 1% of the whole number A.

Atoms with the same atomic number but different neutron numbers, and hence different mass numbers, are known as isotopes. A little more than three-quarters of naturally occurring elements exist as a mixture of isotopes (see monoisotopic elements), and the average isotopic mass of an isotopic mixture for an element (called the relative atomic mass) in a defined environment on Earth, determines the element's standard atomic weight. Historically, it was these atomic weights of elements (in comparison to hydrogen) that were the quantities measurable by chemists in the 19th century.

The conventional symbol Z comes from the German word Zahl meaning number, which, before the modern synthesis of ideas from chemistry and physics, merely denoted an element's numerical place in the periodic table, whose order is approximately, but not completely, consistent with the order of the elements by atomic weights. Only after 1915, with the suggestion and evidence that this Z number was also the nuclear charge and a physical characteristic of atoms, did the word Atomzahl (and its English equivalent atomic number) come into common use in this context.


The periodic table and a natural number for each element[edit]

Russian chemist Dmitri Mendeleev, creator of the periodic table.

Loosely speaking, the existence or construction of a periodic table of elements creates an ordering of the elements, and so they can be numbered in order.

Dmitri Mendeleev claimed that he arranged his first periodic tables (first published on March 6, 1869) in order of atomic weight ('Atomgewicht').[1] However, in consideration of the elements' observed chemical properties, he changed the order slightly and placed tellurium (atomic weight 127.6) ahead of iodine (atomic weight 126.9).[1][2] This placement is consistent with the modern practice of ordering the elements by proton number, Z, but that number was not known or suspected at the time.

Niels Bohr, creator of the Bohr model.

A simple numbering based on periodic table position was never entirely satisfactory, however. Besides the case of iodine and tellurium, later several other pairs of elements (such as argon and potassium, cobalt and nickel) were known to have nearly identical or reversed atomic weights, thus requiring their placement in the periodic table to be determined by their chemical properties. However the gradual identification of more and more chemically similar lanthanide elements, whose atomic number was not obvious, led to inconsistency and uncertainty in the periodic numbering of elements at least from lutetium (element 71) onward (hafnium was not known at this time).

The Rutherford-Bohr model and van den Broek[edit]

In 1911, Ernest Rutherford gave a model of the atom in which a central nucleus held most of the atom's mass and a positive charge which, in units of the electron's charge, was to be approximately equal to half of the atom's atomic weight, expressed in numbers of hydrogen atoms. This central charge would thus be approximately half the atomic weight (though it was almost 25% different from the atomic number of gold (Z = 79, A = 197), the single element from which Rutherford made his guess). Nevertheless, in spite of Rutherford's estimation that gold had a central charge of about 100 (but was element Z = 79 on the periodic table), a month after Rutherford's paper appeared, Antonius van den Broek first formally suggested that the central charge and number of electrons in an atom was exactly equal to its place in the periodic table (also known as element number, atomic number, and symbolized Z). This proved eventually to be the case.

Moseley's 1913 experiment[edit]

Henry Moseley in his lab.

The experimental position improved dramatically after research by Henry Moseley in 1913.[3] Moseley, after discussions with Bohr who was at the same lab (and who had used Van den Broek's hypothesis in his Bohr model of the atom), decided to test Van den Broek's and Bohr's hypothesis directly, by seeing if spectral lines emitted from excited atoms fitted the Bohr theory's postulation that the frequency of the spectral lines be proportional to the square of Z.

To do this, Moseley measured the wavelengths of the innermost photon transitions (K and L lines) produced by the elements from aluminum (Z = 13) to gold (Z = 79) used as a series of movable anodic targets inside an x-ray tube.[4] The square root of the frequency of these photons (x-rays) increased from one target to the next in an arithmetic progression. This led to the conclusion (Moseley's law) that the atomic number does closely correspond (with an offset of one unit for K-lines, in Moseley's work) to the calculated electric charge of the nucleus, i.e. the element number Z. Among other things, Moseley demonstrated that the lanthanide series (from lanthanum to lutetium inclusive) must have 15 members—no fewer and no more—which was far from obvious from known chemistry at that time.

Missing elements[edit]

After Moseley's death in 1915, the atomic numbers of all known elements from hydrogen to uranium (Z = 92) were examined by his method. There were seven elements (with Z < 92) which were not found and therefore identified as still undiscovered, corresponding to atomic numbers 43, 61, 72, 75, 85, 87 and 91.[5] From 1918 to 1947, all seven of these missing elements were discovered.[6] By this time, the first four transuranium elements had also been discovered, so that the periodic table was complete with no gaps as far as curium (Z = 96).

The proton and the idea of nuclear electrons[edit]

In 1915, the reason for nuclear charge being quantized in units of Z, which were now recognized to be the same as the element number, was not understood. An old idea called Prout's hypothesis had postulated that the elements were all made of residues (or 'protyles') of the lightest element hydrogen, which in the Bohr-Rutherford model had a single electron and a nuclear charge of one. However, as early as 1907, Rutherford and Thomas Royds had shown that alpha particles, which had a charge of +2, were the nuclei of helium atoms, which had a mass four times that of hydrogen, not two times. If Prout's hypothesis were true, something had to be neutralizing some of the charge of the hydrogen nuclei present in the nuclei of heavier atoms.

In 1917, Rutherford succeeded in generating hydrogen nuclei from a nuclear reaction between alpha particles and nitrogen gas,[7] and believed he had proven Prout's law. He called the new heavy nuclear particles protons in 1920 (alternate names being proutons and protyles). It had been immediately apparent from the work of Moseley that the nuclei of heavy atoms have more than twice as much mass as would be expected from their being made of hydrogen nuclei, and thus there was required a hypothesis for the neutralization of the extra protons presumed present in all heavy nuclei. A helium nucleus was presumed to be composed of four protons plus two 'nuclear electrons' (electrons bound inside the nucleus) to cancel two of the charges. At the other end of the periodic table, a nucleus of gold with a mass 197 times that of hydrogen was thought to contain 118 nuclear electrons in the nucleus to give it a residual charge of +79, consistent with its atomic number.


The discovery of the neutron makes Z the proton number[edit]

All consideration of nuclear electrons ended with James Chadwick's discovery of the neutron in 1932. An atom of gold now was seen as containing 118 neutrons rather than 118 nuclear electrons, and its positive charge now was realized to come entirely from a content of 79 protons. After 1932, therefore, an element's atomic number Z was also realized to be identical to the proton number of its nuclei.

The symbol of Z[edit]

The conventional symbol Z possibly comes from the German word Atomzahl (atomic number).[8] However, prior to 1915, the word Zahl (simply number) was used for an element's assigned number in the periodic table.

Chemical properties[edit]

Each element has a specific set of chemical properties as a consequence of the number of electrons present in the neutral atom, which is Z (the atomic number). The configuration of these electrons follows from the principles of quantum mechanics. The number of electrons in each element's electron shells, particularly the outermost valence shell, is the primary factor in determining its chemical bonding behavior. Hence, it is the atomic number alone that determines the chemical properties of an element; and it is for this reason that an element can be defined as consisting of any mixture of atoms with a given atomic number.

New elements[edit]

The quest for new elements is usually described using atomic numbers. As of 2021, all elements with atomic numbers 1 to 118 have been observed. Synthesis of new elements is accomplished by bombarding target atoms of heavy elements with ions, such that the sum of the atomic numbers of the target and ion elements equals the atomic number of the element being created. In general, the half-life of a nuclide becomes shorter as atomic number increases, though undiscovered nuclides with certain 'magic' numbers of protons and neutrons may have relatively longer half-lives and comprise an island of stability.

See also[edit]

Look up atomic number in Wiktionary, the free dictionary.


  1. ^ abThe Periodic Table of Elements, American Institute of Physics
  2. ^The Development of the Periodic Table, Royal Society of Chemistry
  3. ^Ordering the Elements in the Periodic Table, Royal Chemical Society
  4. ^Moseley, H.G.J. (1913). 'XCIII.The high-frequency spectra of the elements'. Philosophical Magazine. Series 6. 26 (156): 1024. doi:10.1080/14786441308635052. Archived from the original on 22 January 2010.
  5. ^Eric Scerri, A tale of seven elements, (Oxford University Press 2013) ISBN978-0-19-539131-2, p.47
  6. ^Scerri chaps. 3–9 (one chapter per element)
  7. ^Ernest Rutherford, New Zealand history online. (19 October 1937). Retrieved on 2011-01-26.
  8. ^Origin of symbol Z.
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Atomic Number:

Atomic number (z) is the number of protons in an atom.

Since for a neutral atom, the number of electrons is equal to number of protons, so Atomic Number is also equal to the number of electrons in neutral the atom.

Atomic number = Number of protons = Number of Electrons (for neutral atom)

Example: The atomic number of Iron (Fe) is 26; Fe atom contains 26 protons and 26 electrons.

Mass Number:

The mass number (A) is defined as the sum of the number of protons and neutrons present in the nucleus of an atom. It is also referred as Atomic Mass.

Example: Mass number of Nitrogen atom is 14 then it contains 7 protons and 7 neutrons.

Mass number = Number of Protons + Number of Neutrons

Notation of element :

Atomic number Element Mass number= Z XA

Example: Sodium atom has notation as 11Na23 .

Question: An atom X has mass number 40 and atomic number 18. Find out the number of protons, number of electrons and number of neutrons present in the atom X?


Number of protons:

Number of protons = Atomic number = 18

Number of electrons:

In an atom number of electrons = Number of protons = 18

Number of Neutrons = Mass number – Atomic number = 40-18 = 22

Electronic Configuration :

The arrangement of electrons in the shells is known as electronic configuration.

Electrons are present in fixed energy levels called Shells. These Shells are also called Orbits. These orbits or Shells are represented by the letters K, L, M, N,… or the numbers n = 1, 2, 3, 4,5,.….

Note: Shell and Orbits are not identical ; They are different ideas. Here for simplification, Students are advised to consider the sense same. Shell more correctly indicates a Sphere (a three dimensional object) while Orbits are Circular paths (A two dimensional object). If Bohr’s Model is considered to be valid; Shells and Orbits are similar for indication purpose.

Rules for accommodating electrons in various shells (Bohr-Bury Rules)

  1. The maximum number of electrons that can be accommodated in any energy level of the atom is given by a formula 2n2 where n is the number of that Orbit number/ Shell Number / Energy level.
Orbit Number Shell NameMaximum number of Electrons

(Using 2n2)

1K Shell2
2L Shell8
3M Shell18
4N Shell32

2. After a series of experiments and a detailed study by scientists like Louis de Broglie, Schrödinger, Somerfield and others proved that shells or energy levels have Sub Shells within them. Electrons are present in these sub shells which constitute a Shell.

Types of Sub Shell : Sub shells are s, p, d and f types.

Every sub-shell can accommodate a fixed number of electrons.

“s” sub shell can hold a maximum of 2 electrons.

“p” sub shell can hold a maximum of 6 electrons.

“d” sub shell can hold a maximum of 10 electrons.

“f” sub shell can hold a maximum of 14 electrons.

Arrangements of Shell and Sub – Shells :

Shell NameSub – Shell name Number of Electrons

in that sub shell

K shell (1st Shell)1s2
L Shell (2nd Shell)2s




M Shell (3rd Shell)3s






Energy Order of different Sub-Shells :


3. Electrons are filled up in an atom from Lower Energy Level to Higher Energy Level. This rule is called Aufbau Principle.

Electronic Configuration of Sodium atom:

Sodium has atomic number 11 and mass number 23.

The nucleus of sodium has 11 protons and 12 neutrons and it is surrounded by 11 electrons.

Now Its electronic configuration is 1s2, 2s2, 2p6, 3s1

Note: you can write electronic configuration easily, follow steps :

1. Write the sub-shell energy sequence: 1s<2s<2p<3s<3p<4s<3d<4p<5s

2. Check how many electrons are present in the atom

3. Fill these electrons in the sequence till all electrons are not exhausted; remember sub-shell can accommodate maximum number of electrons as s sub-shell = 2 ; p sub-shell =6, d sub-shell = 10

The electronic configuration of sodium: 1s2, 2s2, 2p6, 3s1 can also be written as 2, 8, 1 (indicating 1st shell has 2 electrons, 2nd shell has 8 electrons, 3rd shell has 1 electron).

K shell or I shell = 2 electrons

L– Shell or II shell = 8 electrons

M– Shell or III shell = 1 electron

∴ Electronic configuration of sodium atom = 2, 8, 1

Similarly the electronic configuration of other atoms are :

Element Atomic Number Electronic Configuration

Type 1

Electronic Configuration

Type 2

Hydrogen (H)11s11
Lithium (Li)31s2, 2s12,1
Carbon (C)61s2, 2s2, 2p22,4
Aluminium (Al)131s2, 2s2, 2p6,3s2 ,3p12,8,3
Scandium (Sc)211s2, 2s2, 2p6, 3s2, 3p6,4s2,3d12,8,9,2
Vanadium (V)231s2, 2s2, 2p6, 3s2, 3p6,4s2,3d32,8,11,2
Iron (Fe)261s2, 2s2, 2p6, 3s2, 3p6,4s2,3d62,8,14,2

Note the changes in the electronic configuration of Sc, V, Fe .

Q. Now try to write the electronic configuration of

(a) Boron (B) atomic number = 5

(b) Florine (F) atomic number = 9

(c) Titanium (Ti) atomic number = 22

(d) Nickel (Ni) atomic number = 28

Valence electrons:

The electrons, which are present in the outermost shell of an atom are called valence electrons.

Example: Sodium

Atomic number of sodium is 11

Ni Ka Atomic Number

Electronic configuration of sodium is 2, 8, 1

In sodium 3rd shell is the outermost shell (valence shell). In this shell it has 1 electron. Hence the number of valence electrons present in sodium is 1.

The chemical properties of an atom are dependent on these valence electrons. Since they are ones that are participate in a chemical reaction.

  • Elements having valence electrons 1, 2 or 3 are called metals.

Exception: Hydrogen has 1 valence electron but it is not consider as a metal.

  • These elements lose electrons easily and form a positively charged ion called cation.

Example: Na – e → Na+

  • Elements having valence electrons 4, 5, 6 or 7 are called as Non-metals
  • These elements gain electrons and forms a negatively charged ion called anion.

Example: F + e → F

  • Elements with same number of valence electrons have similar chemical properties; whereas the elements with different valence electrons have different chemical properties.


The combining capacity of the atoms to form molecules either with same or different elements is defined as Valency.

Valency can be considered as number of Valence electrons; but in some cases it may have different meaning . It is due to the combining property of atoms to form stable molecules.

The valency of element is either equal to the number of valency electron is it atom or equal to number of electron required to complete to octet (eight electrons) in valency shell.

For example :

  • Atom contains less than four electrons in its outermost shell, the Valency of an atom is equal to the number of electrons present in the valence shell. More correctly it is called Electrovalency.

Example: Sodium has one electron in its outermost shell, so the Valency of sodium is 1.

Calcium has two electrons in its outermost shell, so the valency of calcium is 2.

Aluminum has three electrons in its outermost shell, so the valency of aluminum is 3.

  • If the outer shell has more than four electrons, the valency = 8 – the number of electrons in the outer shell. More correctly it is called Covalency.
  • Covalency :Number of electrons shared by one atom of an element to achieve inert gas configuration.

For Chlorine (atomic number = 17) electronic configuration is 1s2, 2s2, 2p6, 3s2, 3p5 or 2,8,7.

Its outermost shell has 7 electrons , therefore its valency can be either 7 or 8-7=1.

1 valency indicates Cl combining capacity is 1 or it can form 1 chemical bond in its molecules.

Octet Rule:

An element with 8 electrons in an outermost-shell is said to possess a complete Octet.

Nickel Atom Symbol

Atoms or ions with octet configuration are stable.

Example: Noble gases possess complete Octet configuration ; they are nonreactive and stable (Except Helium, which possess 2 electrons in its outermost shell).

Atomic mass: The total number of protons and neutrons present in one atom of an element is called atomic mass.

Helium has two protons and two neutrons.

Mass number of He = 2 + 2 = 4

Mass of He atom 2 + 2 = 4u (unified mass)

Isotopes: Atoms of the same element with the same atomic number but different atomic masses are called Isotopes.

Examples: 1H1, 1H2, 1H3 are the isotopes of hydrogen,6 C12 ; 6C14 are isotopes of carbon.

The chemical properties of all the isotopes of an element are the same. This is due to the presence of same number of electrons.

Isobars: Atoms of different elements with different atomic numbers but have the same mass number are called Isobars.

Example: 18 Ar40, 19 K40 and 20Ca40.

Reasons for chemical reactivity of an atom: The chemical activity of an atom depends on the number valence electrons. An atom with complete octet configuration is chemically inert and it does not participate in chemical reactions.

Example: Noble gases.

The atoms of element with incomplete octet configuration are chemically active. These elements combine with others to attain stable electronic configuration (octet configuration).

In simple words, atoms combine together so that they acquire 8 electrons in their outermost shell or valence shell.

Mass Of Nickel

Example 1: Hydrogen has one electron and it requires one more electron to attain the nearest inert gas configuration. To achieve this each hydrogen atom contributes an electron and form hydrogen molecule.

Ni Atomic Number

Example 2: Chlorine atom has 7 electrons in its valence shell and it requires 1 more electron to attain the stable octet configuration. To achieve this, each chlorine atom shares an electron with each other and form chlorine molecule.