# Classification of Elements and Periodicity in Properties

Why Do We Need to Classify Elements?

Chemists discovered that all the substances are made up of atoms or elements. Slowly they discovered many elements. In 1789, Antoine Lavoisier published a list of 33 elements. With the discovery of many elements chemists felt the need of classification of elements for their easy understanding and comparison. At present 118 elements are known. Efforts to synthesize new elements are continuing. It is very difficult to study the properties of such a huge number of chemical elements individually. Scientists were trying to classify elements in a periodic manner on the basis of their various properties.

Genesis of Periodic Classification

Classification by Johann Dobereiner – German chemist, Johann Dobereiner classified certain elements on the basis of their similar properties in the groups of three elements each. He called these groups triads. In each triad, the atomic weight of the middle element was equal to the average of the atomic weights of the first and third element.

 Triad Lithium Sodium Potassium Atomic Weight 7 23 39 Na = $\frac{39+7}{2}$ = 23

Newlands Law of Octaves – English chemist John Alexander Newlands profounded the Law of Octaves in 1865. He arranged the elements in increasing order of their atomic weights and found that every 8th element shows similarity with the 1st element.

 7Li 9Be 11B 12C 14N 16O 19F 23Na 24Mg 27Al 28Si 31P 32S 35.5Cl 39K

Mendeleev’s Periodic Table – In the year 1869, Dmitri Mendeleev arranged all 63 elements in rows or columns in order of their atomic weight. He left the space for corresponding elements in his periodic table which were not even discovered then. Although he was able to predict the properties of those elements through his periodic classification of elements.

Periodic law given by Mendeleev – The properties of the elements are periodic function of their atomic weights.

• In Mendeleev’s periodic table, vertical rows were called groups while horizontal rows were called periods.

• There were nine groups (I, II, III, IV, V, VI, VII, VIII and zero group).

• Group VIII had nine elements which were arranged in triads.

• Zero group had noble gases with 0 valency.

• There were seven periods.

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Modern Periodic Law and the Present Form of the Periodic Table

English physicist Henry Moseley showed through his experiments that the atomic number of the element is its more fundamental property than its atomic mass. So, accordingly periodic law was also changed.

Modern Periodic Law - The properties of the elements are periodic functions of their atomic numbers.

Elements were rearranged in the periodic table according to the modern periodic law. Thus, the modern periodic table was formed. Presently, Modern periodic table or Long form of the periodic table is widely used by chemists. It helps in study of physical and chemical properties of elements.

• In modern periodic table elements have been arranged according to their increasing atomic numbers.

• It has 18 groups and 7 periods.

• Elements of the same group have similar outer electronic configuration.

• The period number corresponds to the highest principal quantum number (n) of the elements in the period.

• In the modern periodic table, 14 elements of both 6th and 7th period are placed in the separate panels at the bottom. These are known as lanthanides and actinides respectively.

Nomenclature of Elements with Atomic Number > 100

To avoid the confusion and conflicts between scientists IUPAC decided that until new elements discovery is proved, and its name is officially recognized, a systematic nomenclature will be followed. Systematic nomenclature is derived directly from the atomic number of the element. Numerical roots for 0 – 9 are used and the suffix ‘ium’ is used in the end.

 Atomic number of the element Notation for digit 1 Notation for digit 0 Notation for digit 1 Suffix Name according to IUPAC nomenclature 101 Un Nil Un Ium Unnilunium

 Notation for Digits 0-9 According to IUPAC Nomenclature of Elements Digit Name Abbreviation 0 nil n 1 un u 2 bi b 3 tri t 4 quad q 5 pent p 6 hex h 7 sept s 8 oct o 9 enn e

 Atomic Number Name According to IUPAC Nomenclature Symbol IUPAC Official Name IUPAC Symbol 101 Unnilunium Unu Mendelevium Md 102 Unnilbium Unb Nobelium No 103 Unniltrium Unt Lawrencium Lr 104 Unnilquadium Unq Rutherfordium Rf 105 Unnilpentium Unp Dubnium Db 106 Unnilhexium Unh Seaborgium Sg 107 Unnilseptium Uns Bohrium Bh 108 Unniloctium Uno Hassium Hs 109 Unnilennium Une Meitnerium Mt 110 Ununnilium Uun Darmstadtium Ds 111 Unununium Uuu Rontgenium Rg 112 Ununbium Uub Copernicium Cn 113 Ununtrium Uut IUPAC Official name yet to be announced - 114 Ununquadium Uuq Flerovium Fl 115 Ununpentium Uup IUPAC Official name yet to be announced - 116 Ununhexium Uuh Livermorium Lv 117 Ununseptium Uus IUPAC Official name yet to be announced - 118 Ununoctium Uuo IUPAC Official name yet to be announced -

Electronic Configuration of Elements and the Periodic Table

Electronic Configurations in Periods – The period number indicates the value of n for the outermost or valence shell. An element placed in 2nd period will have its outermost electrons in 2s or 2p orbitals. Ne is placed in 2nd period and has electronic configuration – 1s2 2s2 2p6. This period has 8 elements (Ne - 2s2 2p6).

Electronic Configurations in Groups – Elements of the same group have similar valence shell electronic configurations. They have the same number of electrons in the outer orbitals.

 Elements of the First Group Atomic Number Symbol Electronic Configuration 3 Li [He]2s1 11 Na [Ne]3s1 19 K [Ar]4s1 37 Rb [Kr]5s1 55 Cs [Xe]6s1 87 Fr [Rn]7s1

## Types of Elements: s, p, d, f – Blocks

 Types of Elements s- block Group 1 and 2 Alkali metals and alkali earth metals p- block Group 13 – 18 Representative or main group elements d- block Group 3 – 12 Transition elements f- block Lanthanides and actinides (4f and 5f) Inner transition elements Elements after Uranium are called transuranium elements.

Periodic Trends in Properties of Elements

Following properties of elements show a very clear periodic trends in periodic table –

• Ionization energy

• Electron affinity

• Electronegativity

• Valence electrons

• Valency

• Metallic character of the elements

• Non – metallic character of the elements

• Reactivity of elements

• Melting and boiling points of elements

Periodic Trend of Atomic Radius Across a Period – As we move from left to right in a period, atomic radius gradually decreases.

Reason – As we move left to right in a period atomic number of the elements increases so nuclear charge increases while number of shells in elements remain the same.

## Example –

 Elements of 2nd Period Li Be B Atomic Number 3 4 5 Nuclear Charge or Number of Protons in the Nucleus 3 4 5 Number of Shells 2 2 2 Atomic Radius (in pm) 152 106 88

Exceptional Behavior – Noble gases show exceptional behavior. The atomic radii of inter gases suddenly increase as compared to its predecessor halogen atom. The reason for this type of exceptional behavior is that atomic radius refers to van der Waal’s radius in case of noble gases while in case of other elements it refers to covalent radius.

Across a Group – On moving top to bottom in a group, atomic radii gradually increase as nuclear charge and number of shells also increase.

Ionization Energy

Periodic trend of ionization energy across a period – As we move from left to right in a period, ionization energy gradually increases.

Reason – As we move left to right in a period atomic size or atomic radius decreases while nuclear charge increases.

## Example -

 Elements of 3rd Period Al Si P Atomic Number 13 14 15 Nuclear Charge or Number of Protons in the Nucleus 12 14 15 Number of Shells 3 3 3 First Ionization Energy 577.5 786.5 1011.8

Exceptional Behavior – Beryllium possesses more first ionization energy than Boron. Because beryllium has half - filled s – orbital and more energy is required to remove an electron from half or completely filled orbitals. That is why noble gases also show exceptionally high ionization energies.

Across a Group – On moving top to bottom in a group, ionization energy gradually decreases as atomic radius increases.

Electron Affinity

Periodic Trend of Electron Affinity Across a Period – As we move from left to right in a period, electron affinity gradually increases.

Reason – As we move left to right in a period atomic size or atomic radius decreases while nuclear charge increases.

 Elements of 4th Period Ti V Cr Atomic Number 22 23 24 Nuclear Charge or Number of Protons in the Nucleus 22 23 24 Electron Affinity (eV) 0.075 0.527 0.675

Exceptional Behavior – Beryllium does not form a stable anion, so it releases less energy than boron on adding an electron. While nitrogen neither releases nor requires a significant amount of energy on adding an electron so it has electron affinity almost equal to zero.

Across a Group – On moving top to bottom in a group, electron affinity gradually decreases.

Electronegativity

Across a Period – As we move left to right across a period, electronegativity increases in the periodic table. Fluorine is the most electronegative element.

Reason – As the nuclear charge increases of an atom, its electron loving character also increases.

## Example –

 Elements of 3rd Period Na Mg Al Atomic Number 11 12 13 Nuclear Charge or Number of Protons in the Nucleus 11 12 13 Electronegativity (Pauling scale) 0.93 1.31 1.61

Across a Group – As we move top to bottom in a group, electronegativity decreases.

Valence Electrons

Across a Period – As we move left to right across a period in the periodic table, the number of valence electrons increases.

## Example –

 Elements of 3rd Period Na Mg Al Atomic Number 11 12 13 Electronic Configuration 2,8,1 2,8,2 2,8,3 Valence Electrons 1 2 3

Across a Group – Across a group, valence electrons remain constant. It means elements present in the same group have the same number of valence electrons. For example, hydrogen, lithium, and sodium elements are present in the 1st group and have the same number of valence electrons which is one.

Valency

Valency is the combining capacity of an atom.

Across a Period – On moving left to right across a period in the periodic table, first valency increases then decreases.

## Example –

 Elements of 2nd Period Li Be B C N O F Ne Atomic Number 3 4 5 6 7 8 9 10 Electronic Configuration 2,1 2,2 2,3 2,4 2,5 2,6 2,7 2,8 Valency 1 2 3 4 3 2 1 0

Across a Group – There is no change in valency across a group. Elements of the same groups show the same valency.

Metallic Character of the Elements

Across a Period – As we move left to right across a period in the periodic table, metallic character of elements decreases.

## Example –

 Elements of 2nd Period Li Be B C N O F Ne Metallic Character Metal Metal Metalloid Nonmetal Nonmetal Nonmetal Nonmetal Nonmetal

Across a Group – As we move top to bottom in a group of the periodic table the metallic character increases.

Non - Metallic Character of the Elements

Across a Period – As we move left to right across a period in the periodic table, nonmetallic character of elements increases.

## Example –

 Elements of 2nd Period Li Be B C N O F Ne Nonmetallic Character Metal Metal Metalloid Nonmetal Nonmetal Nonmetal Nonmetal Nonmetal

Across a Group – As we move top to bottom in a group of periodic table nonmetallic character decreases.

## Example -

 Group 15 Nonmetallic Character N Nonmetal P Nonmetal As Metalloid Sb Metalloid Bi Metal

Reactivity of Elements

Reactivity of metals depends on its electropositive character. So, more is the metallic character, more is the electropositive nature of the element and more is its reactivity. As metallic character decreases across a period left to right, so reactivity also decreases. Although reactivity of nonmetals increases on moving left to right across a period. Thus, we can conclude, as we move left to right in a period, the reactivity of elements gradually decreases up to the group thirteen and then starts increasing.

 Elements of 3rd Period Na Mg Al Si P S Cl Ar Group 1 2 13 14 15 16 17 18 Reactivity Very reactive Reactive Reactive Least reactive Reactive Reactive Very reactive Inert Reactivity decreases 🡪 Reactivity increases🡪

Melting and Boiling Points of Elements

Melting and boiling points of metals decrease gradually from top to bottom in a group. While melting and boiling points of nonmetals increase on moving from top to bottom in a group of the periodic table.

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