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.

SOME COMMON RULES

Pauli Exclusion Principle

So, according to these rules, if we look at the first shell, it only contains two electrons in the first subshell. The quantum number can be described as 1s^2. However, another rule is that two electrons cannot have identical quantum numbers. So if two electrons have to fit inside one s subshell, how will this work?

This matter directly relates to the Pauli Exclusion Principle, which states that two electrons cannot exist in the same location, and thus electrons in the same orbital must have opposing spins. Hydrogen only has one electron in the 1s orbital and has a quantum number of 1s^1. Helium has two electrons in this field, so they must have opposing spins, and it has a quantum number of 1s^2.

This spin factor is the fourth quantum number, m, and is described as either +1/2 for a spin 'up' or -1/2 for a spin 'down'.

Hund's Rule

So now that we've reviewed the orbital structure, let's dive in to how the electrons fill these orbitals. Because electrons are both negatively charged, there is a certain amount of repulsion that prevents them from wanting to fill the same space. So according to Hund's Rule, electrons will fall into empty orbitals of the same energy before electrons begin to pair up into the same orbital.

Aufbau Principle

Closely related to Hund's Rule is the Aufbau Principle, which states that electrons will fill the lower energy levels before moving to higher energy orbitals. Remember that each orbital has two electrons and the number of orbitals at an energy level depends on the first two quantum number.

QUANTUM NUMBERS

There are four quantum numbers; their symbols are n, l, m and s. EVERY electron in an atom has a specific, unique set of these four quantum numbers. The story behind how these numbers came to be discovered is a complex one indeed and is one best left for another day.
A warning before I proceed: this is a 100% non-mathematical discussion. The equations governing electron behavior in atoms are complex. This area of study generated LOTS of Nobel Prizes and the reasoning leading to the above mentioned equations is sophisticated and sometimes quite subtle.
Just keep this in mind: EVERY electron's behavior in an atom is governed by a set of equations and that n, l, m, and s are values in those equations. EVERY electron in an atom has a unique set of quantum numbers.
Lastly, I'm going to reserve to another discussion what these numbers mean. I will just describe their existence and the rules for how to determine them in this tutorial. The next tutorial will start with hydrogen and assign quantum numbers to its electron, then proceed to helium and do the same, then lithium, beryllium, and so on.
I. The Principal Quantum Number (signified by the letter 'n'): This quantum number was the first one discovered and it was done so by Niels Bohr in 1913. Bohr thought that each electron was in its own unique energy level, which he called a "stationary state," and that each electron would have a unique value of 'n.'
In this idea, Bohr was wrong. It very quickly was discovered that more than one electron could have a given 'n' value. For example, it was eventually discovered that when n=3, eighteen different electrons could have that value.
Keep in mind that it is the set of four quantum numbers that is important. As you will see, each of the 18 electrons just mention will have its own unique set of n, l, m, and s.
Finally, there is a rule for what values 'n' can assume. It is:

n = 1, 2, 3, and so on.

n will always be a whole number and NEVER less than one.
One point: n does not refer to any particular location in space or any particular shape. It is one component (of four) that will uniquely identify each electron in an atom.
II. The Azimuthal Quantum Number (signified by the letter 'l'): about 1914-1915, Arnold Sommerfeld realized that Bohr's 'n' was insufficient. In other words, more equations were needed to properly describe how electrons behaved. In fact, Sommerfeld realized that TWO more quantum numbers were needed.
The first of these is the quantum number signified by 'l.' When Sommerfeld started this work, he used n' (n prime), but he shifted it to 'l' after some years. I'm not sure why, but it seems easier to print l than n prime and what if the printer (of a textbook) accidently dumps a few prime symbols, leaving just the letter 'n?' Ooops!
The rule for selecting the proper values of 'l' is as follows:

l = 0, 1, 2, . . . , n-1

l will always be a whole number and will NEVER be as large as the 'n' value it is associated with.
III. The Magnetic Quantum Number (signified by the letter 'm' or m_l): this quantum number was also discovered by Sommerfeld in the same 1914-1915 time frame. I don't think he discovered one and then the other, I think that him realizing the need for two runs together somewhat. I could be wrong in this, so don't take my word for it!
The rule for selecting m is as follows:

m starts at negative 'l,' runs by whole numbers to zero and then goes to positive 'l.'

For example, when l = 2, the m values used are -2, -1, 0, +1, +2, for a total of five values.
IV. The Spin Quantum Number (signified by the letter 's' or m_s): spin is a property of electrons that is not related to a sphere spinning. It was first thought to be this way, hence the name spin, but it was soon realized that electrons cannot spin on their axis like the Earth does on its axis. If the electron did this, its surface would be moving at about ten times the speed of light (if memory serves correctly!). In any event, the electron's surface would have to move faster than the speed of light and this isn't possible.
In 1925, Wolfgang Pauli demonstrated the need for a fourth quantum number. He closed the abstract to his paper this way:

"On the basis of these results one is also led to a general classification of every electron in the atom by the principal quantum number n andtwo auxiliary quantum numbers k₁ and k₂ to which is added a further quantum number m₁ in the presence of an external field. In conjunction with a recent paper by E. C. Stoner this classification leads to a general quantum theoretical formulation of the completion of electron groups in atoms."

In late 1925, two young researchers named George Uhlenbeck and Samuel Goudsmit discovered the property of the electron responsible for the fourth quantum number being needed and named this property spin.
The rule for selecting s is as follows:

after the n, l and m to be used have been determined, assign the value +1/2 to one electron, then assign -1/2 to the next electron, while using the same n, l and m values.

For example, when n, l, m = 1, 0, 0; the first s value used is +1/2, however a second electron can also have n, l, m = 1, 0, 0; so assign -1/2 to it.