Periodic Classification of Elements

Matter is something that has some mass and occupies some space is what we have learnt in our schools. Subsequently, matter consists of molecules, which are made up of compounds of various combinations of atoms. Mankind has always been curious to study about the properties and behaviour of matter and identified that there are different types of atoms that ultimately make up the matter. We all know we refer to this type of atoms by the term ‘elements’. This article discusses the periodic classification of elements & how it changed with time.

It was recent that IUPAC named 4 more elements and added them to The Modern periodic table – which we will discuss upon later in this article. Science has developed so much that from just 31 elements identified in 1800, we have discovered and named 117 elements at present. Right from the initial days, scientists did feel a need to classify the elements so that their properties and behaviour can be studied collectively, thus simplifying an otherwise tedious process of learning about each element one at a time.

History of Periodic Classification of Elements

One of the first attempts to classify elements dates back to 1829; when only 31 elements were known. Johann Dobereiner, a German chemist had identified groups of three elements, each showing similar behaviour. These groups were known as Dobereiner’s Triads. Few of these trials were:

1.Lithium (Li), Sodium(Na), Potassium(K)

2.Calcium(Ca), Strontium(Sr), Barium(Ba)

3.Chlorine(Cl), Bromine(Br), Iodine(I)

The Law of Triads for this classification observed that the atomic weight of the middle element in each triad is roughly the average of the other two and the properties of the middle element are the intermediate of the first and last element. This classification failed because all the elements known couldn’t be made into triads and the observations were largely considered coincidental.

A French geologist, A.E.B de Chancurtois came up with a cylindrical table of elements in 1862, in which elements were arranged in the increasing order of their atomic weights. This model too didn’t earn much recognition.

Next notable classification was produced in 1865 by John Alexander Newlands, an English chemist. He profounded the Law of octaves, which states that when elements are arranged in the increasing order of their atomic weights, every eighth element showed properties similar to the first element. The work of Newlands too threw light on periodic repetition of properties as de Chancurtois’, but could be applied to elements up to Calcium (Ca) only.

The independent works of Russian chemist Dmitri Mendeleev and German chemist Lothar Mayer in the latter half of 1800s led to the formulation of Periodic Law, which states that properties of elements are a periodic function of their atomic weights. Both Mendeleev and Lothar Meyer arranged the elements in the form of rows and columns, keeping similar elements in the same column. But Mendeleev studied the compounds formed by the elements of a group and even went on to change the order atomic weights in few cases; so that elements in each group exhibit similar properties. He even left gaps in his table and predicted the properties of elements that would fill the gaps.

The Modern periodic table, developed by an English physicist Henry Mosely, is based on the the studies through X-ray diffraction, which revealed that atomic number– an indicator of the no. of protons in the nucleus, which is unique for any element –  is the more fundamental criteria than atomic mass. He stated the modern periodic law as – properties of elements are a periodic function of their atomic numbers.  Keeping atomic number as the base also enables us to predict how many more elements are yet to be found out in between the two known elements in the Modern periodic table. Later when noble gases were invented, they could be easily accommodated in the table.

Periodic Classification of Elements – Present form of periodic table (Long form)

Based on the modern periodic law, the elements are arranged into rows and columns based on their atomic number. There are 18 columns, called groups and 7 rows, called periods. We know that electrons are distributed in the various shells around the nucleus. The period represents the outermost shell to which electron is added.

The groups represent elements with the same number of electrons in their outer-most shell, called the valence shell.

Hydrogen (atomic number-1) shows similarity in properties to group-1 and group-17 elements, hence is positioned separately on top of the table.

Due to large number of elements in 6^th and 7^th period, 14 elements from both are kept outside the main table, called Lanthanoids and Actinoids respectively.

Valency

The reactivity of elements are determined based on the rule of Octet rule – which states that elements lose or gain their outermost electrons (valence electrons) to achieve noble gas configuration which is 8 electrons in their valence shell. So elements having 1,2 or 3 elements in their valence shell can achieve stable noble gas configuration by losing electrons and those having 5,6 or 7 can make up 8 electrons by gaining electrons. This number electrons lost or gained is referred to as valency of the element and is an indicator of the reactivity of the element.

Trends in the periodic table

As said earlier, properties of elements are a periodic function of their atomic numbers. The arrangement of elements in the table form allows us to identify certain trends in their physical and chemical properties along the groups and periods.

Let us consider atomic size, one of the important physical property of elements. As we go down a group, the number of shells containing electrons increases, significantly increasing the atomic size. The trend across a group is governed by the effective nuclear charge. As we go right in a group, the effective nuclear charge increases, pulling the electrons closer to the nucleus and reducing the atomic size. This effect is lesser in magnitude as compared to the trend down the table.

The metallic and non-metallic character of elements can be largely attributed to the number of valence electrons and valence. Metallic character is characterised by the ease of elements lose electrons to attain octet. Elements having lesser than 4 electrons in valence shell are hence metallic, which means metals are found towards the left of the table. As we go down a group, the valence electrons are much farther away from the nucleus, which means they can be released easier, increasing the metallic character. As for non-metals, they are found towards the right of the table; and are generally electron acceptors so as to complete their octet. Down a group, accepting an electron becomes a tougher job, decreasing non-metallic character.

The general trends in the properties are summarised as shown below:

The trends discussed above are the general observed trends. There are a lot of other factors like ionisation enthalpy, electronegativity, electron gain enthalpy; as seen in the figure above. governing the properties of elements, which you will learn in higher classes. The combinations of these properties give rise to exceptional behaviour and anomalies, thus making the learning of elements beyond generalisation yet interesting. Moseley’s modern periodic table can be considered the best available to us now, but as extensive scientific research reveals more understanding of matter every day, the quest for a perfect periodic classification of elements still continues.

Filled orbitals and number of elements in different periods

Types of elements

Salient Features of Modern Periodic Table

There are eighteen vertical columns known as groups in the modern periodic table which are arranged from left to right and seven horizontal rows which are known as periods.

• Elements are arranged on the basis of increasing order of atomic numbers.
• The horizontal rows in the periodic table are called Periods and the vertical columns are known as the Groups.
• The elements in the Modern Periodic Table are arranged in 7 Periods and 18 Groups.
• The elements in the Modern Periodic Table are classified into four categories viz. Representative Elements, Transition Elements, Inner Transition Elements, and Noble Gases.

Physical properties

S-block elements have their sub-shells filled with electrons. Alkali metals have one electron in their outermost s-sub shell. Let’s take a look at the physical properties of Alkali metals.

1. Atomic radii: As we know atomic size decreases as we move from left to right in the periodic table, s-block metals – alkali metals to be precise- have the largest atomic radii in a particular group. Also, the atomic and ionic sizes increase as we move down the group.
2. Physical State: All the elements exist as silvery white, soft, light metals and are considered metals due to very low ionization energies. These elements can be compressed into sheets or drawn into wires. When freshly cut, they’re lustrous but the lustre tarnishes easily with exposure to air.
3. Density: The density increases as we move from Li to Fr but there’s an anomaly in case of sodium and potassium where K is lighter than Na. Li is the lightest among all the metals. Li, Na, and K are lighter when compared to water.
4. Ionization Enthalpy: s-block metals have low values of ionization enthalpies as it is easier to remove their valence electrons to attain stability. Ionization enthalpy decreases as we move down a group as the outer electrons feel a lesser pull from the nucleus. This is predominant in the last members of each group, where the shielding of nuclear charge by inner electrons means the outer electrons are very loosely attached to the atom (the effect is called as the Screening effect), further decreasing its ionization enthalpy
5. Melting & Boiling Points: The presence of large atomic radius makes these elements bind very weakly at the time of forming crystals. These elements have very low melting and boiling points because of a weak inter-atomic bonding.
6. Colour: When we heat most of the s-block elements or their salts, the outermost electrons get excited and jump to higher energy levels. When they come back to their ground state, colours characteristic to each metal is imparted to their flames. This is rather used in the laboratories to identify the action in salts using a test called the `flame test’. This can’t be applied for Magnesium and Beryllium as their outer electrons are too bound to get thermally excited.
7. Hydration Enthalpy: Alkaline earth metals and Lithium generally form hydrated salts. The extent of hydration decreases with size and hence decreases down the group.

Group-13 Elements

• Group-13 elements belong to the first group of p-block elements and include Boron (B), Aluminium (Al), Gallium (Ga), Indium (In) and Thallium (Tl).
• Boron is the only metalloid in the group while the rest of the elements are metals. Gallium is liquid at temperatures over 30^oC.
• They contain three electrons in the outer most shell i.e. a filled s-orbital and one electron in the p-orbital. (ns²np¹)
• Atomic Radius and density increases down the group with Thallium (Tl) having the largest atomic radius and highest density. (**Atomic radius of gallium is less than Aluminium due to poor shielding effect of d-orbital electrons.)
• First Ionization Energy decreases down the group. (except Thallium)
• Boiling Point decreases down the group.
• Electronegativity decreases from Boron to Aluminium, but then increases marginally. This discrepancy is due to aberrant trends in atomic size.
• Sum of first three ionization enthalpies decreases considerably from boron to aluminium, but then increases due to poor shielding effect of intervening d- and f- orbital electrons.
• The relative stability of the +1 oxidation state increases from Aluminium to Thallium.
• These are electron deficient elements, hence act as good Lewis acids. The tendency to behave as Lewis acids decreases down the group.
• Oxide of boron is acidic, oxides of Aluminium and gallium are amphoteric, while oxides of indium and thallium are basic in nature.
• Boron and Aluminium react at high temperatures with oxygen and nitrogen to form oxides and nitrides respectively.
• Boron is nonreactive at moderate temperatures to acids and alkalis, whereas Aluminium reacts with both acids and bases.
• Boron, in the form of borax (Na₂B₄O₇.10H₂O), is used in household cleaning products, as a pH buffer for gel electrophoresis, as flux for soldering and in laboratory tests for transition metals due to the characteristic colours of their metaborates when exposed to flame.
• Orthoboric acid is used in the manufacture of fibreglass, LCD display screen glass and as a dry lubricant for carom boards.
• Aluminium is used in the manufacture of aircraft bodies, food packaging materials in the form of aluminium foil, pipes, tubes, coins and electrical wires (due to high electrical conductivity).
• Lithium Aluminium Hydride (LiAlH₄) and Sodium borohydride (NaBH₄) are potent reducing agents used in organic chemistry.

Group-14 Elements

• Group-14 elements include Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn) and Lead (Pb).
• Carbon is a non-metal, silicon and germanium are metalloids, tin and lead are metals.
• They contain four electrons in the outer most shell i.e. a filled s-orbital and two electrons in the p-orbital. (ns²np²).
• There is a considerable increase in covalent radius from carbon to silicon, but a small increase from silicon to lead. This is due to poor shielding effect of completely filled d and f orbitals.
• First ionization enthalpy is higher than the corresponding members of group 13. Ionisation enthalpy generally decreases down the group. Discrepancies in the trend are also due to poor shielding effect of d and f orbital electrons.
• Electronegativity from silicon to lead is almost same, but higher than the corresponding group 13 members.
• Boiling point decreases down the group from silicon to lead.
• Melting and Boiling points are higher than the corresponding group 13 members.
• A tendency to show +2 oxidation state increases from germanium to lead. Stability of +4 oxidation state decreases down the group.
• Lead is stable in +2 oxidation state and acts as an oxidising agent in +4 oxidation state.
• Presence of d-orbital in elements apart from carbon increases their tendency to form complexes by accepting electron pairs from donors.
• Two types of oxides are formed, on heating with oxygen- monoxides and dioxides. Dioxides tend to be more acidic than monoxides.
• Acidic nature of dioxides decreases down the group; tin and lead dioxide is amphoteric.
• Stability of dihalides increases down the group, whereas stability of tetrahalides decreases down the group.
• Carbon has a unique property of forming long chains by linking with other carbon atoms. This is called catenation. Tendency of catenation decreases down the group (Lead does not show catenation.)
• Carbon also has unique ability to form pπ– pπ multiple bonds with itself and with other atoms of small size and high electronegativity.
• Due to catenation and formation of pπ– pπ multiple bonds, allotropic forms of carbon, namely graphite, diamond and fullerenes exist.
• Carbon, in form of graphite fibres, is used to make tennis racquets and aircrafts. Graphite is used as an electrode in batteries.
• Activated charcoal is used in water purification and removal of poisonous gases.
• Diamond is used in jewellery and in cutting equipment.
• Silicon is used in electronics and manufacture of cement.
• Germanium is used in synthesis of polyethylene terephthalate.
• Tin is used to make alloys and solder, used for making connections in electrical circuits.
• Lead is used in radiation shields.

Group-15 Elements

• Group-15 elements include Nitrogen (N), Phosphorous (P), Arsenic (As), Antimony (Sb) and Bismuth (Bi).
• Nitrogen and Phosphorous are non-metals, arsenic and antimony are metalloids while bismuth is a metal.
• They contain five electrons in the outer most shell i.e. a filled s-orbital and three electrons in the p-orbital. (ns²np³). The p-orbital is half-filled, making it more stable.
• Covalent and ionic radii increase in size down the group. The lack of regular trend is due to poor shielding effect of d and f orbitals.
• Ionisation enthalpy decreases down the group due to increase in atomic size, and is higher than the corresponding group 15 members due to extra stability of half-filled p orbitals.
• Electronegativity decreases down the group.
• Metallic character increases down the group.
• Boiling point increases up to arsenic and then decreases up to bismuth.
• Allotropy is seen in all elements except nitrogen.
• Stability of -3 and +5 oxidation state decreases down the group.
• Nitrogen can form multiple pπ– pπ bonds, whereas phosphorous and arsenic can form dπ– pπ bonds.
• Oxides formed are of two types with different oxidation states of the element – +3 and +5.
• Oxides of higher oxidation state are more acidic and acidic character of these oxides decreases down the group.
• All elements react with metals and possess -3 oxidations states in their respective compounds.
• Nitrogen is used as a coolant, to make fertilizers, plastics and explosives. It is also used in cryopreservation, pharmaceuticals and X-ray detectors.
• Phosphorous is used in fertilizers along with nitrogen, in making matches to light fires, softening hard water, and military applications like smoke bombs.
• Arsenic is used in semiconductors in the form of gallium arsenide. Arsenic-lead alloys are used to manufacture bullets.
• Antimony is used as a dopant in semiconductors.
• Bismuth is used as a catalyst to make acrylic fibres and in cosmetics.

Group-16 Elements

• Group-16 elements include Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po).
• Oxygen and Sulphur are non-metals, selenium and tellurium are metalloids, whereas polonium is a metal.
• They contain six electrons in the outer most shell i.e. a filled s-orbital and four electrons in the p-orbital. (ns²np⁴).
• Atomic radii, ionic radii, metallic character, melting and boiling points increase down the group.
• Ionization enthalpy decreases down the group due to increase in size, but is less than the corresponding elements of Group-15. (Extra stable Half-filled p-orbitals)
• Electronegativity decreases down the group.
• All elements exhibit allotropy.
• Stability of -2 and +6 oxidation states decreases down the group. Oxygen can show only -2 oxidation state, except in case of OF₂ (+2 oxidation state).
• Stability of +4 oxidation state increases down the group due to inert pair effect.
• Acidic character of hydrides and their reducing power increases down the group. (Except H₂O).
• Reducing property of oxides decreases down the group and their oxides are acidic in nature.
• These can form either dihalides or tetrahalides or hexahalides. Only hexafluorides are stable.
• Oxygen is used as liquid fuel in rockets, in oxy-acetylene flames used for welding, in smelting iron into steel and as a breathing gas by organisms on the planet.
• Sulphur, in the form of sulphuric acid, is used in the manufacture of fertilizers, paints, dye and storage batteries, in petroleum refining, cleansing metals and as a laboratory reagent. Sulphur is also used in the vulcanization of rubber.
• Selenium is used in manganese electrolysis and in alloys with bismuth.
• Tellurium is used in solar panels.

Group-17 Elements

• Group-17 elements include Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At).
• They contain seven electrons in the outer most shell i.e. a filled s-orbital and five electrons in the p-orbital. (ns²np⁵).
• They have the smallest atomic radii in their respective periods. Atomic, Ionic radii, melting and boiling points increase down the group, as expected.
• Ionisation enthalpy, electronegativity, oxidising power, affinity to hydrogen, ionic character of metal halides, all decrease down the group.
• Bond dissociation enthalpy decreases down the group, however, due to large electron-electron repulsion in fluorine, its bond dissociation enthalpy is close to that of chlorine.
• Oxidation states exhibited are -1, +1, +3, +5 and +7. Fluorine exists in -1 oxidation state only.
• The acidic strength of hydrides and stability of oxides increases down the group (I>Br>Cl).
• Fluorine is used to fluorinate uranium and polonium, used as nuclear fuel in reactors. It is also used in PET imaging and PTFE synthesis.
• Chlorine is used as a disinfectant, in bleaching wood pulp for paper manufacture, in extraction of gold, manufacture of dyes and organic solvents etc.
• Bromine is used in photography, in flame retardants, and synthesis of brominated polymers.
• Iodine is used as a laboratory reagent for testing ammonium ions (Nessler’s reagent), as disinfectant, for cloud seeding etc.

Group-18 elements

• Group-18 elements constitute the last group of p-block elements and include Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon (Rn).
• Helium is an exception as it does not have p-orbitals. All the elements in the group are gases.
• They contain eight electrons in the outer most shell i.e. a filled s-orbital and a completely filled p-orbital, with the exception of Helium (ns²np⁶).
• Exhibit very high ionization enthalpies due to high stability and it decreases down the group.
• Atomic radii, melting and boiling points increase down the group.
• Only xenon reacts with fluorine and oxygen to form fluorides and oxides respectively.
• Helium is used in filling hot air balloons and in superconducting magnets.
• Neon is used in bulbs and lights for advertisement and display purposes.
• Argon is used to provide an inert atmosphere for metallurgical reactions.

D-block elements show typical metallic behaviour. These are characterised by high tensile strength, malleability, ductility, electrical and thermal conductivity as well as metallic lustre. Except for Zinc (Zn), Cadmium (Cd), Mercury (Hg) and Manganese (Mn), they have one or more typical metallic lattices. Due to the larger extent of metallic bonding by virtue of d-electrons, they are hard and have high melting and boiling points. The group-12 elements (Zn, Cd and Hg) shows the exception in this regard also.

Melting and Boiling point

The high melting and boiling points of d-block elements can be attributed to the involvement of d-orbital electrons in addition to the s-electrons in metallic bonding. Melting and boiling points increases as the d-orbital get filled. This trend goes till d⁵ configuration and then decreases regularly as the orbital gets further filled and the electrons get paired in the orbital. A point to be noted here is that Manganese (Mn) and Technetium (Tc) have abnormally low melting and boiling points.

• The exceptional case of Mercury – the liquid metal: Mercury is one among the two elements (other being Bromine(Br) ) and the only metal that exists in its liquid state at room temperature. This is unexpected for the general outlook of a metal and can be explained by the fact that 6s valence electrons of Mercury are more closely pulled by the nucleus, rendering those outer s-electrons less involved in metallic bonding. Still, Mercury shows good electrical conductivity.

Atomic and Ionic Sizes

The ionic sizes show gradual decrease as we move right across a series. This is because as we number of electrons increase, the nuclear charge to increase. Due to poor shielding of nuclear charge by d-electrons, the net increase in nuclear charge outweighs the effect of the added electron, thereby reducing the size.

Similar trend and reasons can be observed for the case of atomic radii as well, though the decrease is much gradual. In case of going down the series, the atomic radii show increase. There is a huge jump from 3d to 4d in terms of atomic size, but the 4d and 5d series have a small difference only. This is explained on the basis of Lanthanoid contraction. In case of 5d series, the inner 4f orbitals are filled before 5d orbitals. The poor shielding ability of 4f electrons renders the outer electrons greater nuclear pull, causing lanthanoid contraction. Due to this, the expected increase in atomic size is compensated by increased nuclear pull, keeping the size nearly same.

Ionisation enthalpy

The same reasons that govern the trend in atomic size can explain the gradual increase in ionisation enthalphy as we move from left to right in a series. However, there exist large anomalies in this general trend. The anomalies arise due to the fact that removal of an electron alters the combined energy considerations of s and d orbital system. The concepts like hybridisation, pairing and exchange energy plays its role here, but let keep away such complexities for the time being. In simple terms, we can assume that the ionisation enthalpy will be high if removal of electron leads to deviation from a stable configuration – half or completely filled d-orbitals or hybridised orbitals.

Multiple oxidation states and stability

One of the most notable features of d-block elements are the variety of oxidation states exhibited by them. This leads to a large number of compounds and pool of reactions involved. The number of oxidation states shown increases with the number of electrons available to lose/share. It decreases as the number of paired electrons increase making the number of orbitals available for sharing less. Hence, the first and last elements show fewer oxidation states and the elements with a maximum number of oxidation states are towards the centre of the table. For example, Manganese (Z=25) shows states from Mn²⁺ to Mn⁷⁺, whereas Sc³⁺and Zn²⁺ are the only other states shown by Scandium (Z=23) and Zinc (Z=30) respectively.

In p-block, the heavier elements prefer lower oxidation states due to what is called inert pair effect. But in case of d-block elements, the higher oxidation states are more stable for heavier members in a group. For example, in group-6 ( Chromium (Cr), Molybdenum(Mo) and Tungsten(W) ), the +6 states for W and Mo are stable whereas Cr⁶⁺, as in potassium dichromate, easily reduces thus being a common oxidising agent.

Lower oxidation states in these metals are stabilised by ligands like CO, which are pi-electron donors, whereas the higher oxidation states are stabilised by electronegative elements like Fluorine(F) and Oxygen(O).Hence the high oxidation compounds of these metals are mainly fluorides and oxides.

Standard electrode potential

There are mainly the M²⁺/M and M³⁺/M²⁺ reduction potentials that are generally considered, for the states M²⁺ and M³⁺ respectively.  The reduction potential values are governed by the ionisation enthalpies, size, electronic configuration of the ions and the energy levels of hybridised t_2g and e_g orbitals, which are beyond the scope of this article.

Magnetic properties

Paramagnetism, Diamagnetism and ferromagnetism are the general properties exhibited by substances. Here ferromagnetism is an extreme level of paramagnetism.

The parametric behaviour of d-block elements is due to the presence of unpaired electrons. Such electrons contribute to ‘orbital magnetic moment’ and ‘spin magnetic moment’. However, for 3d series, the orbital angular moment is negligible and the approximate spin-only magnetic moment is given by the formula:

μ =  √n(n+2)

where n is the number of unpaired electrons. Its unit is Bohr Magneton (BM). For higher d-series, the actual magnetic moment includes components from the orbital moment in addition to the spin moment.

‘Periodic Law

The periodic law was developed independently by Dmitri Mendeleev and Lothar Meyer in 1869. Mendeleev created the first periodic table and was shortly followed by Meyer. They both arranged the elements by their mass and proposed that certain properties periodically reoccur. Meyer formed his periodic law based on the atomic volume or molar volume, which is the atomic mass divided by the density in solid form. Mendeleev’s table is noteworthy because it exhibits mostly accurate values for atomic mass and it also contains blank spaces for unknown elements.

Introduction

In 1804 physicist John Dalton advanced the atomic theory of matter, helping scientists determine the mass of the known elements. Around the same time, two chemists Sir Humphry Davy and Michael Faraday developed electrochemistry which aided in the discovery of new elements. By 1829, chemist Johann Wolfgang Doberiner observed that certain elements with similar properties occur in group of three such as; chlorine, bromine, iodine; calcium, strontium, and barium; sulfur, selenium, tellurium; iron, cobalt, manganese. However, at the time of this discovery too few elements had been discovered and there was confusion between molecular weight and atomic weights; therefore, chemists never really understood the significance of Doberiner’s triad.

In 1859 two physicists Robert Willhem Bunsen and Gustav Robert Kirchoff discovered spectroscopy which allowed for discovery of many new elements. This gave scientists the tools to reveal the relationships between elements. Thus in 1864, chemist John A. R Newland arranged the elements in increasing of atomic weights. Explaining that a given set of properties reoccurs every eight place, he named it the law of Octaves.

The Periodic Law

In 1869, Dmitri Mendeleev and Lothar Meyer individually came up with their own periodic law “when the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically.” Meyer based his laws on the atomic volume (the atomic mass of an element divided by the density of its solid form), this property is called Molar volume.

Atomic (molar) volume (cm3/mol)= molar mass (g/ mol)ρ (cm3/g)(1)(1)Atomic (molar) volume (cm3/mol)= molar mass (g/ mol)ρ (cm3/g)

Mendeleev’s Periodic Table

Mendeleev’s periodic table is an arrangement of the elements that group similar elements together. He left blank spaces for the undiscovered elements (atomic masses, element: 44, scandium; 68, gallium; 72, germanium; & 100, technetium) so that certain elements can be grouped together. However, Mendeleev had not predicted the noble gases, so no spots were left for them.

Figure 1: Mendeleev’s original periodic table4

In Mendeleev’s table, elements with similar characteristics fall in vertical columns, called groups. Molar volume increases from top to bottom of a group 3

Atomic Number as the Basis for the Periodic Law

Assuming there were errors in atomic masses, Mendeleev placed certain elements not in order of increasing atomic mass so that they could fit into the proper groups (similar elements have similar properties) of his periodic table. An example of this was with argon (atomic mass 39.9), which was put in front of potassium (atomic mass 39.1). Elements were placed into groups that expressed similar chemical behavior.

In 1913 Henry G.J. Moseley did researched the X-Ray spectra of the elements and suggested that the energies of electron orbitals depend on the nuclear charge and the nuclear charges of atoms in the target, which is also known as anode, dictate the frequencies of emitted X-Rays. Moseley was able to tie the X-Ray frequencies to numbers equal to the nuclear charges, therefore showing the placement of the elements in Mendeleev’s periodic table. The equation he used:

ν=A(Z−b)^2

with

• νν: X-Ray frequency
• ZZ: Atomic Number
• AA and bb: constants

With Moseley’s contribution the Periodic Law can be restated:

Similar properties recur periodically when elements are arranged according to increasing atomic number.”

Atomic numbers, not weights, determine the factor of chemical properties. As mentioned before, argon weights more than potassium (39.9 vs. 39.1, respectively), yet argon is in front of potassium. Thus, we can see that elements are arranged based on their atomic number. The periodic law is found to help determine many patterns of many different properties of elements; melting and boiling points, densities, electrical conductivity, reactivity, acidic, basic, valance, polarity, and solubility.

The table below shows that elements increase from left to right accordingly to their atomic number. The vertical columns have similar properties within their group for example Lithium is similar to sodium, beryllium is similar to magnesium, and so on.

Elements in Group 1 (periodic table) have similar chemical properties and are called alkali metals. Elements in Group 2 have similar chemical properties, they are called the alkaline earth metals.

Short form periodic table

The short form periodic table is a table where elements are arranged in 7 rows, periods, with increasing atomic numbers from left to right. There are 18 vertical columns known as groups. This table is based on Mendeleev’s periodic table and the periodic law.

Long form Periodic Table

In the long form, each period correlates to the building up of electronic shell; the first two groups (1-2) (s-block) and the last 6 groups (13-18) (p-block) make up the main-group elements and the groups (3-12) in between the s and p blocks are called the transition metals. Group 18 elements are called noble gases, and group 17 are called halogens. The f-block elements, called inner transition metals, which are at the bottom of the periodic table (periods 8 and 9); the 15 elements after barium (atomic number 56) are called lanthanides and the 14 elements after radium (atomic number 88) are called actinides.

Hope this article provided a fair idea of the periodic classification of elements. You can read about the s-block elements here. All the Best!