Plus One Chemistry Notes Chapter 3 Classification of Elements and Periodicity in Properties is part of Plus One Chemistry Notes. Here we have given Kerala Plus One Chemistry Notes Chapter 3 Classification of Elements and Periodicity in Properties.
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Kerala Plus One Chemistry Notes Chapter 3 Classification of Elements and Periodicity in Properties
The systematic classification of elements made the study of elements easy. In this unit we will study the historical development of the periodic table and also learn how elements are classified.
Genesis of Periodic Classification Law of Triads (Dobereiner)
When the elements in a triad were arranged in the order of increasing atomic masses, the atomic mass of the middle element was found to be approximately equal to the arithmetic mean of the other two elements.
Only a limited number of elements could be grouped into triads.
Newland’s Law of Octaves
Newland arranged the elements in increasing order of their atomic masses and noted that every eighth element had properties similar to the first element, like the eighth note in an octave of music.
eg., Li(7), Be(9), B(11), C(12), N(14), 0(16), F(19), Na(23)
Drawbacks of Newiand’s Law of Octaves:
- It failed to explain the heavier elements beyond calcium.
- When the noble gases were discovered, the idea of octaves could not be held.
Lother Meyer’s Curve of Atomic Volume Vs Atomic Mass
Lother meyer noticed that the elements with similar properties occupied similar positions on the curve. He observed a periodicity in the properties of the elements with atomic mass,
eg., Li, Na, K, etc. occupied the peak positions.
Mendeleev’s Periodic Law
(മെൻഡലീവിന്റെ — ആവർത്തന നിയമം)
This law states that “the properties of the elements are periodic function of their atomic mass”.
Merits of Mendeleev’s Periodic Table:
- Systematic study of the elements.
- Prediction of new elements and their proper-ties.
- Correction of certain atomic masses.
Mendeleev corrected the atomic masses of Beryllium (Be), gold (Au) and platinum (Pt).
Demerits in the Mendeleev’s Periodic Table:
- Position of hydrogen. No proper position for hydrogen in the Mendeleev’s periodic table.
- Separation of chemically similar elements in different groups,
eg., Cu and Hg.
- Grouping of chemically dissimilar elements,
eg., Alkali metals and coinage metals are grouped together.
- Inversion in the periodic table. Certain pairs of elements had to be placed in the reverse order of atomic masses in order to conform the periodic law.
- Position of lanthanides and actinides.
Modern Periodic Law
The physical and chemical properties of the elements are periodic functions of their atomic numbers.
Long form periodic table
Bohr constructed a periodic table known as the long form periodic table. It has been modified in the light of IUPAC (1984) recommendations.
Nomenclature of Elements with Atomic Number > 100
Notation. IUPAC nomenclature of elements above 100 for 0 digit (nil), 1(un), 2(bi), 3(tri),
4(quad), 5(pent), 6(hex), 7(sept), 8(oct), 9(enn).
Atomic number = Digit’s name + ium.
Symbol = First letter of digit’s name,
eg., Atomic number = 101, Unnilunium
Symbol = Unu.
IUPAC = Mendelevium (Md)
Atomic number = 109, Unnilennium
Symbol = Une
IUPAC = Meitnerium (Mt)
Electronic Configurations and Types of Elements
(ഇലക്ട്രോൺ വിന്യാസവും വിവിധതരം മുലക – ങ്ങളും )
The s-Block elements
The elements of group 1 (alkali metals) and group 2 (alkaline earth metals) which have ns1 and ns2 outermost electronic configuration belong to the s-block Elements. They are all reactive metals with low ionization enthalpies. The compounds of the s-block elements, with the exception of those of lithium and beryllium are predominantly ionic.
The p-Block Elements
The p-block elements comprise those belonging to group 13 to 18 and these together with the s-block elements are called the representative elements or main group elements. The outermost electronic configuration varies from to in each period. All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. Preceding the noble gas family are two chemically important groups of non-metals. They are the halogens (group 17) and the chalcogens (group 16). These two groups of elements have high negative electron gain enthalpies and readily add one or two electrons respectively to attain the stable noble gas configuration.
The d-Block Elements (Transition Elements)
These are the elements of group 3 to 12 in the centre of the periodic table. These are characterised by the filling of inner d-orbitals by electrons and are therefore referred to as d- block elements. These elements have the general outer electronic configuration. They are all metals. They mostly form coloured ions, exhibit variable valence (oxidation states), paramagnetism and oftenly used as catalysts. In a way, transition metals form a bridge between the chemically active metals of s-block elements and the less active elements of groups 13 and 14 and thus take their familiar name “transition elements”.
The f-Block Elements (Inner-Transition Elements)
The two rows of elements at the bottom of the periodic table called the lanthanoids. The last electron added to each element is filled in f-orbital. These two series of elements are hence called the inner transition elements. They are all metals. Within each series, the properties of the elements are quite similar. The elements after uranium are called transuranium elements.
Those properties which vary in a periodic manner are termed as periodic properties,
eg., Ionisation enthalpy (I.E).
- Atomic radius is defined as the distance between the nucleus and the outermost electronic level in an atom.
- For a homonuclear diatomic molecule the atomic radius is half the interatomic distance.
- Covalent radius (C.R) is defined as half the distance between the centres of the nuclei of two similar atoms bonded by single covalent bond.
- Metallic radius (M.R) is taken as half the internuclear distance separating the metal cores in the metallic crystal.
- Van der Waal’s radius (V.R) is one half of the distance between the centres of the nuclei of two non bonded atoms of the adjacent molecules of the element in the solid state.
Note: V.R > M.R > C.R
Which Is the smallest among Na+, Mg2+, Al3+ and why?
A2+ is smallest because it has the highest number of protons (13) among Na+, Mg2+, Al3 ions, due to which effective nuclear charge is maximum.
Ionic radius may be defined as the distance be-tween the centre of the nucleus of an ion and its outermost orbitals.
- Anion is always larger than the neutral atom.
- Cation is much smaller than the neutral atom.
- As charge of anion increase, radius increases.
- As charge of carbon increases, radius decreases.
Note: Alkali metals have the largest size in period.
Atoms or ions having same number of electrons are called isoelectronic species.
A quantitative measure of the tendency of an element to lose electron is given by its Ionization enthalpy. It represents the energy required to remove an electron from an isolated gaseous atom in its ground state.
Electron Gain Enthalpy
It is the amount of energy released when an isolated gaseous atom accepts an electron to form a monovalent gaseous anion.
A quantitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity. Unlike ionization enthalpy and electron gain enthalpy, it is not a measurable quantity.
Write the atomic number of the element present in the third period and seventeenth group of the periodic table.
Element in 3rd period and 17th group of periodic table is Cl with atomic number 17.
Periodicity of Valency or Oxidation States
Valency or oxidation state is equal to the number of electrons in the outermost orbitals or equal to eight minus the number of outermost electrons or charge assigned to an atom.
Anomalous Properties of Second Period Elements
The first, second and third elements of second period show resemblance with second, third and fourth elements of the third period. this relationship is known as diagonal relationship.
On the basis of quantum numbers, justify that the sixth period of the periodic table should have 32 elements
Periodic trends and Chemical reactivity
- Atomic radii decrease along a period and increase down the group.
- IE increase along a period and decrease down the group.
- Electronegativity increase along a period and decrease down the group.
It is measured by the relative tendency of an element to lose or gain electrons in a chemical reaction.
- Reactivity of metals decrease along a period and increase down the group.
- Reactivity of non-metal increases along a period and decreases down the group.
- Highly reactive elements do not occur in nature in free state and occur in the combined form.
Give general electronic configuration of least reactive group. Why are they least reactive?
ns2 np6 is general electronic configuration of least reactive group because they have stable electronic configuration.
Arrange the elements F, Cl, 0 and N in the correct order of their chemical reactivity in terms of oxidising property.
F, Cl, O and N can be arranged in order of their chemical reactivity in terms of oxidising property as
F > O > Cl > N.
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