STATES OF MATTER

Three States Of Matter
Matter
It is defined as any thing which has mass and occupies space is called matter.
Matter is composed of small and tiny particles called Atoms or molecules. It exist in three different states which are gaseous, liquid & solid.
Properties of Gas
1. It has no definite shape.
2. It has no definite volume, so it can be compressed or expanded.
3. A gas may diffuse with the other gas.
4. The molecules of a gas are in continuous motion.
Properties of Liquids
1. A liquid has no definite shape.
2. It has a fixed volume.
3. The diffusion of a liquid into the other liquid is possible if both of the liquids are polar or non-polar.
4. It can be compressed to a negligible.
Properties of Solids
1. A solid has a definite shape.
2. It has a fixed volume.
3. The rate of diffusion of solid with each other is very slow.
4. It cannot be compressed easily.
Kinetic Theory of Gases
It was an idea of some scientist like Maxwell & Bolzmann that the properties of gases are due to their molecular motion. This motion of the molecules is related with the kinetic energy, so the postulates give by the scientist about the behaviour of gases are collectively known as kinetic molecular theory of gases.
The postulates of kinetic molecular theory are as follows.
1. All gases consists of very large number of tiny particles called molecules.
2. These molecules are widely separated from each other and are so small that they are invisible.
3. The size of the molecules is very small as compared to the distance between them.
4. There is no attractive or repulsive force between molecules so they can move freely.
5. The molecules are very hard and perfectly elastic so when they collide no loss of energy takes place.
6. The gas molecules are in continuous motion they move in a straight path until they collide. The distance between two continuous collision is called Mean Free Path.
7. During their motion these molecules are collided with one another and with the walls of the container.
8. The collision of the molecules are perfectly elastic. When molecules collide they rebound with perfect elasticity and without loss or gain of energy.
9. The pressure of the gas is the result of collision of molecules on the walls of the container.
10. The average kinetic energy of gas molecules depends upon the absolute temperature. At any given temperature the molecules of all gases have the same average kinetic energy (1/2 mv2).
Kinetic Theory of Liquids
This theory is bases on the following assumptions.
1. The particles of a liquid are very close to each other due to which a liquid has fixed volume.
2. The particles in a liquid are free to move so they have no definite shape.
3. During the motion these molecules collides with each other and with the walls of the container.
4. These molecules possess kinetic which is directly proportional to its absolute temperature.
Kinetic Theory of Solids
The assumptions of kinetic theory for solids are as follows.
1. The particles in a solid are very closely packed due to strong attractive forces between the molecules.
2. These molecules are present at a fixed position and are unable to move.
3. They have definite shape because the particles are arranged in a fixed pattern.
4. They possess only vibrational energy.

HYDROGEN BONDING

A hydrogen bond is a weak type of force that forms a special type of dipole-dipole attraction which occurs when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons. These bonds are generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds.

The evidence for hydrogen bonding

Many elements form compounds with hydrogen. If you plot the boiling points of the compounds of the group 14 elements with hydrogen, you find that the boiling points increase as you go down the group.

Figure 1: Boiling points of group 14 elemental halides.

The increase in boiling point happens because the molecules are getting larger with more electrons, and so van der Waals dispersion forces become greater. If you repeat this exercise with the compounds of the elements in groups 15, 16, and 17 with hydrogen, something odd happens.

Figure 2: Boiling points of group 15-17 elemental halides.

Although for the most part the trend is exactly the same as in group 4 (for exactly the same reasons), the boiling point of the compound of hydrogen with the first element in each group is abnormally high. In the cases of [Math Processing Error], [Math Processing Error] and[Math Processing Error] there must be some additional intermolecular forces of attraction, requiring significantly more heat energy to break. These relatively powerful intermolecular forces are described as hydrogen bonds.

Origin of Hydrogen Bonding

The molecules which have this extra bonding are:

Figure 3: The lone pairs responsible for hydrogen bonding in . The solid line represents a bond in the plane of the screen or paper. Dotted bonds are going back into the screen or paper away from you, and wedge-shaped ones are coming out towards you.

Notice that in each of these molecules:

• The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant positive charge.
• Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one "active" lone pair.

Lone pairs at the 2-level have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge. Lone pairs at higher levels are more diffuse and not so attractive to positive things.

Types of hydrogen bonds

Hydrogen bonds can occur within one single molecule, between two like molecules, or between two unlike molecules.

Intramolecular hydrogen bonds

Intramolecular hydrogen bonds are those which occur within one single molecule. This occurs when two functional groups of a molecule can form hydrogen bonds with each other. In order for this to happen, both a hydrogen donor an acceptor must be present within one molecule, and they must be within close proximity of each other in the molecule. For example, intramolecular hydrogen bonding occurs in ethylene glycol (C2H4(OH)2) between its two hydroxyl groups due to the molecular geometry.

Intermolecular hydrogen bonds

Intermolecular hydrogen bonds occur between separate molecules in a substance. They can occur between any number of like or unlike molecules as long as hydrogen donors and acceptors are present an in positions in which they can interact.For example, intermolecular hydrogen bonds can occur between NH3 molecules alone, between H2O molecules alone, or between NH3 and H2O molecules.

Properties and effects of hydrogen bonds

On Boiling Point

When we consider the boiling points of molecules, we usually expect molecules with larger molar masses to have higher normal boiling points than molecules with smaller molar masses. This, without taking hydrogen bonds into account, is due to greater dispersion forces (see Interactions Between Nonpolar Molecules). Larger molecules have more space for electron distribution and thus more possibilities for an instantaneous dipole moment. However, when we consider the table below, we see that this is not always the case.

Compound	Molar Mass	Normal Boiling Point
A	18 g/mol	373 K
B	20 g/mol	292.5 K
C	17 g/mol	239.8 K
D	34 g/mol	212.9 K
E	36.4 g/mol	197.9 K
F	34 g/mol	185.2 K

We see that H2O, HF, and NH3 each have higher boiling points than the same compound formed between hydrogen and the next element moving down its respective group, indicating that the former have greater intermolecular forces. This is because H2O, HF, and NH3 all exhibit hydrogen bonding, whereas the others do not. Furthermore, [Math Processing Error] has a smaller molar mass than HF but partakes in more hydrogen bonds per molecule, so its boiling point is consequently higher.

On Viscosity

The same effect that is seen on boiling point as a result of hydrogen bonding can also be observed in the viscosity of certain substances. Those substances which are capable of forming hydrogen bonds tend to have a higher viscosity than those that do not. Substances which have the possibility for multiple hydrogen bonds exhibit even higher viscosities.

Factors preventing Hydrogen bonding

Electronegativity

Hydrogen bonding cannot occur without significant electronegativity differences between hydrogen and the atom it is bonded to. Thus, we see molecules such as PH3, which no not partake in hydrogen bonding. PH3 exhibits a trigonal pyramidal molecular geometry like that of ammmonia, but unlike NH3 it cannot hydrogen bond. This is due to the similarity in the electronegativities of phosphorous and hydrogen. Both atoms have an electronegativity of 2.1, and thus, no dipole moment occurs. This prevents the hydrogen bonding from acquiring the partial positive charge needed to hydrogen bond with the lone electron pair in another molecule. (see Polarizability)

Atom Size

The size of donors and acceptors can also effect the ability to hydrogen bond. This can account for the relatively low ability of Cl to form hydrogen bonds. When the radii of two atoms differ greatly or are large, their nuclei cannot achieve close proximity when they interact, resulting in a weak interaction.

Hydrogen Bonding in Nature

Hydrogen bonding plays a crucial role in many biological processes and can account for many natural phenomena such as the Unusual properties of Water. In addition to being present in water, hydrogen bonding is also important in the water transport system of plants, secondary and tertiary protein structure, and DNA base pairing.

Plants

The cohesion-adhesion theory of transport in vascular plants uses hydrogen bonding to explain many key components of water movement through the plant's xylem and other vessels. Within a vessel, water molecules hydrogen bond not only to each other, but also to the cellulose chain which comprises the wall of plant cells. This creates a sort of capillary tube which allows for capillary action to occur since the vessel is relatively small. This mechanism allows plants to pull water up into their roots. Furthermore,hydrogen bonding can create a long chain of water molecules which can overcome the force of gravity and travel up to the high altitudes of leaves.

HYBRIDISATION

Hybridisation

When bonds are formed between atoms in the formation of molecules, there is a change in the nature of the Atomic Orbitals in each atom. Effectively, new molecular orbitals are formed, and the electrons in these orbitals are those of the molecule as a whole. The alteration of the structure of the atomic orbitals is called hydridisation, as it involves combining a number of orbitals to create an equal number of new orbitals, where each of the new hybrid orbitals have properties which are an average of those of the orbitals from which they were created.
A number of different types of hybrid orbitals are known for carbon in the organic compounds.

• Hybridisation sp
• Hybridisation sp2
• Hybridisation sp3

The different hybrid orbitals, which have different orientations in space, account for the geometry of the organic compounds of carbon in which they appear.

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Hybridisation sp

The simplest alkyne, Ethyne (i.e. Acetylene), has a linear structure, where the two carbon atoms and the two hydrogens atoms that are attached to these carbon atom lie along a straight line. The carbon to carbon and carbon to hydrogen bonds are arranged as far apart in space as possible. Thus, these bonds are at 1800 to each other. The geometry of this structure cannot be explained using the shape of the Atomic Orbitals on the carbon atom.
The Electronic Configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If we leave two of the 2p sub-orbitals (e.g. the 2py(1) and 2pz(1) sub-orbitals) to form the second and third bonds of the carbon to carbon triple bond, we can rearrange the other two sub-orbitals (i.e. 2s(1) and 2px(1)) to form two equivalent hybrid orbitals. When these new orbitals are arranged as far apart in space as possible, the new orbitals are arranged in a plane and are 180 degrees apart.
The particular hybridisation of the orbitals of carbon atoms in ethene, are called sp1 hybrids orbitals, (i.e. they arises from the hybridising of one sigma orbital and one pi orbitals). The bonds between the hydrogen atoms and carbons, and the first bond (of the double bond) between the carbon atoms of ethene are formed by the end-on overlap of these sp1 hybrids orbitals to form sigma bonds, and accounts for the 180 degree bond angles observed in this compound. The second and third bonds (of the triple bond) between the carbon atoms in ethyne is formed by the side-on overlap of the original atomic unhybridised orbitals (i.e. the 2py(1) and 2pz(1) orbitals) and these are pi bonds.

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Hybridisation sp2

The simplest alkene, Ethene, has a planar structure, where the two carbon atoms and four hydrogens atoms that are attached to these carbon atom lie in a plane. The carbon to carbon and carbon to hydrogen bonds are arranged as far apart in space as possible. Thus, these bonds are at 120 degrees to each other. The geometry of this structure cannot be explained using the shape of the Atomic Orbitals on the carbon atom.
The Electronic Configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If we leave one of the 2p sub-orbitals (e.g. the 2pz(1) sub-orbital) to form the second bond of the carbon to carbon double bond, we can rearrange the other three sub-orbitals (i.e. 2s(1) 2px(1) 2py(1)) to form three equivalent hybrid orbitals. When these new orbitals are arranged as far apart in space as possible, the new orbitals are arranged in a plane and are 120 degrees apart.
The particular hybridisation of the orbitals of carbon atoms in ethene, are called sp2 hybrids orbitals, (i.e. they arises from the hybridising of one sigma orbital and two pi orbitals). The bonds between the hydrogen atoms and carbons, and the first bond (of the double bond) between the carbon atoms of ethene are formed by the end-on overlap of these sp2 hybrids orbitals to form s bonds (sigma bonds), and accounts for the 120 bond angles observed in this compound. The second bond (of the double bond) between the carbon atoms in ethene is formed by the side-on overlap of the original atomic unhybridised orbitals (i.e. the 2pz(1) orbital) and this pi bond prevents free rotation in the carbon to carbon axis.

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Hybridisation sp3

The simplest alkane, Methane, CH4, has a tetrahedral structure, where the four hydrogens that are attached to the central carbon atom are arranged symmetrically about the carbon atom, and arranged as far apart in space as possible. The geometry of this structure cannot be explained using the shape of the Atomic Orbitals on the carbon atom.
The Electronic Configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). In this state, the carbon atom would be expected to produce three bonds of one sort (i.e. from the three Pi orbitals) and one bond of another sort (i.e. from the single Sigma orbital). However, as the four bonds in methane are known to be equivalent, it is necessary to postulate the existence of mixed or hybrid orbitals.
More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If the four sub-orbitals in the second shell are re-arranged to form four equivalent hybrid orbitals, the result is a new second shell with four new orbitals. These new orbitals are arranged as far apart in space as possible. Thus, the new orbitals point to the corners of a regular tetrahedron, and the angle between any two hybrid orbitals 109.5 degrees.
The particular hybridisation of the orbitals of carbon atom, which results in the formation of the hybrid orbitals on the carbon atom in the methane molecule, is called sp(3) (i.e. they arises from the hybridising of one Sigma orbital and three Pi orbitals).

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1s(2) 2s(1) 2p(3). In this state, the carbon atom would be expected to produce three bonds of one sort (i.e. from the three p orbitals) and one bond of another sort (i.e. from the single s orbital). However, as the four bonds in methane are known to be equivalent, it is necessary to postulate the existence of mixed or hybrid orbitals. More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If the four sub-orbitals in the second shell are re-arranged to form four equivalent hybrid orbitals, the result is a new second shell with four new orbitals. These new orbitals are arranged as far apart in space as possible. Thus, the new orbitals point to the corners of a regular tetrahedron, and the angle between any two hybrid orbitals 109.5 degrees. The particular hybridisation of the orbitals of carbon atom, which results in the formation of the hybrid orbitals on the carbon atom in the methane molecule, is called sp(3) (i.e. they arises from the hybridising of one s orbital and three p orbitals). #$K+Hybridisation sp2 The simplest alkene, Ethene, C2H4, has a planar structure, where the two carbon atoms and four hydrogens atoms that are attached to these carbon atom lie in a plane. The carbon to carbon and carbon to hydrogen bonds are arranged as far apart in space as possible. Thus, these bonds are at 120 to each other. The geometry of this structure cannot be explained using the shape of the atomic orbitals on the carbon atom. The electronic configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If we leave one of the 2p sub-orbitals (e.g. the 2pz(1) sub-orbital) to form the second bond of the carbon to carbon double bond, we can rearrange the other three sub-orbitals (i.e. 2s(1) 2px(1) 2py(1)) to form three equivalent hybrid orbitals. When these new orbitals are arranged as far apart in space as possible, the new orbitals are arranged in a plane and are 120 degrees apart. The particular hybridisation of the orbitals of carbon atoms in ethene, are called sp2 hybrids orbitals, (i.e. they arises from the hybridising of one s orbital and two p orbitals). The bonds between the hydrogen atoms and carbons, and the first bond (of the double bond) between the carbon atoms of ethene are formed by the end-on overlap of these sp2 hybrids orbitals to form s bonds (sigma bonds), and accounts for the 120 bond angles observed in this compound. The second bond (of the double bond) between the carbon atoms in ethene is formed by the side-on overlap of the original atomic unhybridised orbitals (i.e. the 2pz(1) orbital) and this p bond (pi bond) prevents free rotation in the carbon to carbon axis. #$K+Hybridisation sp The simplest alkyne, Ethyne, C2H2, (i.e. Acetylene), has a linear structure, where the two carbon atoms and the two hydrogens atoms that are attached to these carbon atom lie along a straight line. The carbon to carbon and carbon to hydrogen bonds are arranged as far apart in space as possible. Thus, these bonds are at 1800 to each other. The geometry of this structure cannot be explained using the shape of the atomic orbitals on the carbon atom. The electronic configuration of carbon in the ground state (i.e. the lowest energy state) is 1s(2) 2s(2) 2p(2). If energy is supplied to raise one of the 2s electrons to a higher energy level to fill the vacant 2p orbital, the electronic configuration of carbon in the excited state, is 1s(2) 2s(1) 2p(3). More specifically, the electronic configuration of the excited carbon atom is 1s(2) 2s(1) 2px(1) 2py(1) 2pz(1). If we leave two of the 2p sub-orbitals (e.g. the 2py(1) and 2pz(1) sub-orbitals) to form the second and third bonds of the carbon to carbon triple bond, we can rearrange the other two sub-orbitals (i.e. 2s(1) and 2px(1)) to form two equivalent hybrid orbitals. When these new orbitals are arranged as far apart in space as possible, the new orbitals are arranged in a plane and are 180 degrees apart. The particular hybridisation of the orbitals of carbon atoms in ethene, are called sp1 hybrids orbitals, (i.e. they arises from the hybridising of one s orbital and one p orbitals). The bonds between the hydrogen atoms and carbons, and the first bond (of the double bond) between the carbon atoms of ethene are formed by the end-on overlap of these sp1 hybrids orbitals to form s bonds (sigma bonds), and accounts for the 1800 bond angles observed in this compound. The second and third bonds (of the triple bond) between the carbon atoms in ethyne is formed by the side-on overlap of the original atomic unhybridised orbitals (i.e. the 2py(1) and 2pz(1) orbitals) and these are p-bonds (pi bonds).

Chemical bonding

Chemical bonding, any of the interactions that account for the association of atoms into molecules, ions, crystals, and other stable species that make up the familiar substances of the everyday world. When atoms approach one another, their nuclei and electrons interact and tend to distribute themselves in space in such a way that the total energy is lower than it would be in any alternative arrangement. If the total energy of a group of atoms is lower than the sum of the energies of the component atoms, they then bond together and the energy lowering is the bonding energy.

The ideas that helped to establish the nature of chemical bonding came to fruition during the early 20th century after the electron had been discovered and quantum mechanics had provided a language for the description of the behaviour of electrons in atoms. However, even though chemists need quantum mechanics to attain a detailed quantitative understanding of bond formation, much of their pragmatic understanding of bonds is expressed in simple, intuitive models. These models treat bonds as primarily of two kinds—namely, ionic and covalent. The type of bond that is most likely to occur between two atoms can be predicted on the basis of the location of the elements in the periodic table, and to some extent the properties of the substances so formed can be related to the type of bonding.

A key concept in a discussion of chemical bonding is that of the molecule. Molecules are the smallest units of compounds that can exist. One feature of molecules that can be predicted with reasonable success is their shape. Molecular shapes are of considerable importance for understanding the reactions that compounds can undergo, and so the link between chemical bonding and chemical reactivity is discussed briefly in this article.

Although simple models of bonding are useful as rules of thumb for rationalizing the existence of compounds and the physical and chemical properties and structures of molecules, they need to be justified by appealing to more sophisticated descriptions of bonding. Moreover, there are some aspects of molecular structure that are beyond the scope of the simple theories. To achieve this insight, it is necessary to resort to a fully quantum mechanical description. In practice, these descriptions entail heavy reliance on computers. Such numerical approaches to the chemical bond provide important information about bonding.

This article begins by describing the historical evolution of the current understanding of chemical bonding and then discusses how modern theories of the formation of chemical bonds have emerged and developed into a powerful description of the structure of matter. After the historical introduction, qualitative models of bonding are discussed, with particular attention given to the formation of ionic and covalent bonds and the correlation of the latter with molecular shapes. The more sophisticated quantum mechanical approaches to bond formation are then described, followed by a survey of a number of special cases that raise interesting problems or lead to important insights.

For a detailed discussion of the structure and properties of atoms, see the article atom. Chemical compounds are surveyed in the article chemical compound, and the elements are described in the article chemical element.

The Modern Periodic Table

Later, Henry Gywn-Jeffreys Moseley showed that the atomic number of an element is numerically equal to the number of electrons round the nucleus. The number of electrons in turn is equal to the number of protons in the nucleus. He suggested that atomic number is a more fundamental property of an element than its atomic mass. When the elements are arranged in the increasing order of their atomic number, most of the defects of Mendeleev's classification get rectified.

The periodic law given earlier is now modified and followed today. It states that "the properties of the elements are periodic functions of their atomic numbers."

The modern periodic table is also known as the long form of the periodic table or the extended form of the periodic table.

In this table, the elements are arranged in the order of increasing atomic numbers in such away that elements with similar properties fall in the same vertical column. These vertical columns are referred to as 'Groups'. There are eight groups, which are further sub-divided into eighteen sub-groups. There are 7 horizontal rows in the periodic table. These are called the periods.

Noble (inert) gases are placed in the 18 or VIII A, which is at the extreme right of the table. Halogens are placed in group 17 or VII A, just ahead of the noble gas elements. The alkali metals (Na, K, etc.) are placed in group 1 or I A and the alkaline earth metals (Ca, Ba, etc.) are placed in 2 or II A which is at the extreme left of the table. Transition elements such as copper, iron, etc., are placed in the 3 to 12 groups, which occur at the middle of the table. The inner transition elements, lanthanides and actinides are placed in two separate series at the bottom of the main body of the periodic table.

The recurrence of similar properties of the elements when they are arranged in the order of increasing atomic number, after certain regular intervals, is called periodicity.

Remember

Relation between Mass Number and Atomic Number Atomic Number (Z) is the number of protons in the nucleus of an atom. It is also equal to the number of electrons since the atom is electrically neutral. Mass Number (A) is the total number of neutrons and protons present in the nucleus of an atom. = +

MENDELEEV PERIODIC TABLE

Mendeleev's Periodic Table

In 1869, just five years after John Newlands put forward his Law of Octaves, a Russian chemist called Dmitri Mendeleev published a periodic table. Mendeleev also arranged the elements known at the time in order of relative atomic mass, but he did some other things that made his table much more successful.

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• Indestructibility of Matter
• Mendeleev's Opinion on Global warming
• Mendeleev eight periodic principles
• Mendeleev on the law of periodicity
• Primary Object of Chemistry
• Principles of Chemistry by Dmitrii Mendeleev
• Rusting of Metals

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Mendeleev realized that the physical and chemical properties of elements were related to their atomic mass in a 'periodic' way, and arranged them so that groups of elements with similar properties fell into vertical columns in his table.

Gaps and predictions Sometimes this method of arranging elements meant there were gaps in his horizontal rows or 'periods'. But instead of seeing this as a problem, Mendeleev thought it simply meant that the elements which belonged in the gaps had not yet been discovered. He was also able to work out the atomic mass of the missing elements, and so predict their properties. And when they were discovered, Mendeleev turned out to be right. For example, he predicted the properties of an undiscovered element that should fit below aluminum in his table. When this element, called gallium, was discovered in 1875 its properties were found to be close to Mendeleev's predictions. Two other predicted elements were later discovered, lending further credit to Mendeleev's table.

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Modern day periodic tables are expanded beyond Mendeleev's initial 63 elements. Most of the current periodic tables include 108 or 109 elements. It is also important to notice how the modern periodic table is arranged. Although we have retained the format of rows and columns, which reflects a natural order, the rows of today's tables show elements in the order of Mendeleev's columns. In other words the elements of what we now call a 'period' were listed vertically by Mendeleev. Chemical 'groups' are now shown vertically in contrast to their horizontal format in Mendeleev's table. (reference)

It is also worthy to note that Mendeleev's 1871 arrangement was related to the atomic ratios in which elements formed oxides, binary compounds with oxygen whereas today's periodic tables are arranged by increasing atomic numbers, that is, the number of protons a particular element contains. Although we can imply the formulas for oxides from today's periodic table, it is not explicitly stated as it was in Mendeleev's 1871 table. The oxides ratio column was not shown in earlier Mendeleev versions.

Mendeleev rewrote each edition of Principles of Chemistry, including all new scientific data-particularly confirmations of the periodic law-and reanalyzing difficulties that had arisen to hinder its confirmation (inert gases, radioactivity, radioactive and rare-earth elements)"

EARLY CLASSIFICATION OF ELEMENTS

Dobereiner's Triads

In the famous atomic theory of John Dalton (1805), it was suggested that the atoms of an element have a characteristic mass. So, attempts were made to classify elements on the basis of their atomic masses.

In the year 1829, Johann Wolfgang Dobereiner, a German scientist, was the first to classify elements into groups based on John Dalton's assertions. He grouped the elements with similar chemical properties into clusters of three called 'Triads'. The distinctive feature of a triad was the atomic mass of the middle element. When elements were arranged in order of their increasing atomic mass, the atomic mass of the middle element was approximately the arithmetic mean of the other two elements of the triad.

Examples of Dobereiner's Triads 

Element	Lithium	Beryllium	Potassium	Arithmetic mean
Atomic mass	7.0	9.0	11.0	9.0
Element	Lithium	Beryllium	Potassium	Arithmetic mean
Atomic mass	7.0	9.0	11.0	9.0
Element	Lithium	Beryllium	Potassium	Arithmetic mean
Atomic mass	7.0	23.0	39.0	23.0
Element	Carbon	Nitrogen	Oxygen	Arithmetic mean
Atomic mass	12.0	14.0	16.0	14.0
Element	Calcium	Strontium	Barium	Arithmetic mean
Atomic mass	40.0	87.5	137	88.1
Element	Chlorine	Bromine	Iodine	Arithmetic mean
Atomic mass	35.0	80.0	127.0	80.6

Defects of Triad Classification

• A large number of similar elements could not be grouped into triads e.g., iron, manganese, nickel, cobalt, zinc and copper are similar elements but could not be placed in the triads.

• It was possible that quite dissimilar elements could be grouped into triads.

• Dobereiner could only classify 3 triads successfully (highlighted in the table).

Since he failed to arrange the then known elements in the form of triads his attempt at classification was not very successful.

In the famous atomic theory of John Dalton (1805), it was suggested that the atoms of an element have a characteristic mass. So, attempts were made to classify elements on the basis of their atomic masses.

Law of Octaves

At an earlier stage, when elements were discovered, there were very less number of elements, and hence there was no need to classify them. But as time passed, the number of elements increased and it became difficult to explain the physical and chemical properties and relations between them. 

Hence, scientists developed various classifications at different times and arranged all the known elements in certain forms. First Prout gave the hypothesis that all elements are made up of hydrogen because the atomic mass of all elements is a simple integral of the atomic mass of hydrogen. But this hypothesis was not valid for a long time as some elements were discovered with fractional atomic mass. 

After Prout’s hypothesis, Dobereiner gave the concept of triad in 1829. According to Dobereiner, elements can be arranged in increasing order of their atomic mass. He arranged all the known elements in groups of three each. These groups are calledtriads. In each triad, elements are arranged in increasing order of their atomic mass.The atomic mass of the middle element is equal to the average atomic mass of both the terminal elements. These triads are also termed as Dobereiner triads. 

Some examples of the Dobereiner triad are as follows,

Elements with atomic weight		Average atomic weight of elements
Li7	Na23	K39	7+392 =23
Cl35.5	Br90	I127	35.5+1272 = 81.25
Ca40	Sr88	Ba137	40+1372 = 88.5
S32	Se79	Te127	32+1272 = 79.5
P31	As75	Sb120	31+1202 = 75.5

Newlands Law of Octaves

After Dobereiner triad, John Newland purposed another way of classification in 1864. 

"According to the law of octave, when elements are placed in increasing order of their atomic mass, every eighth element shows similarity with the preceding eighth element in their physical and chemical properties."

Define Octave

Octave can be defined as a group or series of eight or a series of eight notes occupying the interval between two notes, one having twice or half the frequency of vibration of the other. For example, in the series of eight musical notes, after a certain interval the note will repeat itself.

Newland's law of octaves was also based on musical notes. He arranged all the known elements in increasing order of their atomic mass. He found that in this arrangement, the first and eighth elements are similar in their chemical and physical properties. Hence there is a periodicity in elements when they are arranged in increasing order of their atomic masses.

For example, if we start from Lithium (atomic mass = 6) and arrange elements in increasing order of their atomic masses, the eighth element will be sodium (atomic mass = 23). Both elements show the same physical and chemical properties. Today, we known that both are alkali metals, located in group-1, and hence show same properties. 

John Newland

John Alexander Reina Newland discovered the Periodic Table and arranged all known elements in a tabular form on the basis of their atomic weight. John Newland was a British citizen, born in London. His father was a Scottish Presbyterian minister and his mother was Italian.

His initial schooling was done at home by his father. He went to the Royal College of Chemistry for further studies but he was more interested in social reform. Hence he served as a volunteer with Giuseppe Garibaldi in 1860 to unify Italy. After that he started a laboratory in London to practice as an analytical chemist in 1864 and later he became chief chemist in James Duncan's London sugar refinery. In this refinery, he introduced a number of improvements in processing.

Newlands Periodic Table

Newland arranged all the known elements in the form of a table. There are some vertical rows, later termed as periods, and vertical columns which were later termed as groups. Newland's periodic table is shown below:

H	F	Cl	Co/Ni	Br	Pd	I	Pt/Ir
Li	Na	K	Cu	Rb	Ag	Cs	Tl
GI	Mg	Ca	Zn	Sr	Cd	Ba/V	Pb
Bo	Al	Cr	Y	Ce/La	U	Ta	Th
C	Si	Ti	In	Zr	Sn	W	Hg
N	P	Mn	As	Di/Mo	Sb	Nb	Bi
O	S	Fe	Se	Ro/Ru	Te	Au	Os

In Newland's periodic table, each row of elements consists of a total of seven elements and the eighth element falls under the first element. In the first column Lithium is the first element and sodium is the eighth element. Similarly, the eighth element after sodium is potassium. Hence from Newland's law of octaves, the elements Li, Na and K must have similar chemical and physical properties, and they do.
In the second column, Beryllium (Be) is the first element and Magnesium (Mg) is the eighth one. After magnesium the next eighth elements is calcium (Ca). Hence according to Newland's law, these elements should show similar chemical and physical properties.

In another vertical column with carbon (C) as the first element, the eighth element is silicon (Si). Both C and Si also show similar properties. Both are non-metals and show tetravalency. Thus Newland's law of octaves holds good for all these given columns. Similarly, the last group of halogens, from fluorine (F) to chlorine (Cl) also show similarity in their properties.

Limitations of Newlands Law of Octaves

1. This classification was good with a few elements only, and was not applicable to all elements. It was not valid for elements having atomic masses higher than Ca.
2. After the discovery of Nobel gases, it became difficult to fit them in Newland’s periodic table.
3. Newland’s octave law gave a very important term of periodicity which was repeated in further classifications. 
4. His classification was based on the concept that the atomic mass of an element is the periodic function of it's chemical and physical properties.
5. The same concept of periodicity of elements was repeated in Moseley’s classification and the latest long form of periodic table is also based on the same concept. The only change is the use of atomic number instead of atomic mass. 
6. Newland’s octave law was good till 1913. But in 1913 Henry Moseley established a new concept that the properties of elements varied periodically according to atomic number, not atomic weight. 

But this concept was not applicable to elements. Only a few elements could be arranged in a given triad and some elements were not following the given criteria of atomic mass. In some triads, elements have almost same atomic weight like Fe, Co, Ni and Ru , Rh , Pd.