Mandeleev’s Periodic Law:- The elements ' properties are the periodic function of their atomic masses.
Moseley, the English physicist has shown that an element's atomic numbers are more fundamental properties than its atomic mass. The position of an element in the periodic table therefore depends on its atomic number as compared with its atomic mass.
Modern Periodic Law: Elements physical and chemical properties are the periodic features of their atomic numbers.
Elements are Classified into: s-, p-, d- and f- blocks.
Main Group Elements/ Representative Elements:
The elements of the s-and p-block are called major group elements, or representative elements.
s- block Elements: Group-1 (Alkali metals) and Group-2 (Alkaline earth metals) elements with outermost electronic configurations of ns1 and ns2, respectively.
p- Block Elements: They are part of group-13 to 18. The most electronic parameter outer is ns2 np1-6. He (1s2) is a s-block element but is positioned with the group 18 elements (ns2 np6) because it has a fully filled valence shell and thus exhibits characteristic properties of other noble gasses.
d- block Elements (Transition Elements): They are community 3 to 12 elements having external electronic configuration(n-1) d1-10 ns1-2. There are four transformation sequences 3d, 4d, 5d, and 6d. The 6d-series is incomplete. Atomic radius usually decreases over a period of time, and increases as we descend the group.
f-Block Elements (Inner- Transition Series)
Characterized by the filling of 4 f-orbitals, the lanthanum-following elements are from 58Ce to 71Lu. Actinoids distinguished by 5f-orbitals filling are the elements from Th 103Lr following actinium. The outer electronic configuration is characteristic of (n-2) f1-14 (n-1) d0-1 ns2.
Noble Gases: Group 18 gaseous elements are called noble gasses. The general electronic outermost configuration of noble gases (except for He) is ns2 np6. He 's beautifully made to have 1s2. So the outermost shell of noble gasses is filled up entirely.
Periodicity: The occurrence of similar properties after regular intervals is called periodicity.
Cause of Periodicity: The properties of the elements are that similar electronic product configuration occurs regularly as the number of atoms increases.
Atomic Properties: The physical traits of an element's atom are called atomic properties. Properties like atomic radius, ionic radius, ionization energy, electro-negativity, affinity and valence of electrons, etc., called atomic properties.
Atomic Radius- The size in the atom of any product from the middle of the nucleus to the electrons ' outermost shell is called its atomic radius.
Periodicity-
(a) In Periods- The atomic radius of elements decreases in a period from left to right.
In Group-Atomic element radius increases when moving from top to bottom in a group.
Covalent Radius- Half the inter-nuclear distance between two similar atoms is called covalent radius of any element that is covalently bonded to each other by a single covalent bond.
Van Der Waals’ Radius: Half the inter-nuclear separation of two similar adjacent atoms in solid state belonging to the two adjacent molecules of the same substance is called the van der waals ' radius of that atom.
Metallic Radius: Half the distance between the nuclei of the two adjacent metal atoms in the tightly packed metal lattice is called its metal radius.
Ionic Radius: The effective distance from the center of an ion 's nucleus to which it impacts its electron cloud is called its ionic radius.
A cation is smaller than the parent atom but the anion is larger. Among iso-electronic material, the cation with a higher positive charge has a smaller radius, but anion with a higher negative charge has the larger radius.
Ionisation Enthalpy: The enthalpy of ionisation is the change in molar enthalpy that accompanies the removal of an electron from a gaseous phase atom or ion in its ground state. So change in enthalpy for reaction.
Is the ionisation enthalpy of element M. Like the ionization energies for successive ionization, the successive ionization enthalpy can also be called the 2nd ionization enthalpy (i.e., rH2), the third ionization enthalpy (i.e., rH3), etc. The term ionization enthalpy is used for the initial ionization enthalpy (both in kg mol and in eV).
Periodicity:
The ionisation enthalpies generally follow the order (there are few exceptions):
(∆rH1) < (∆rH2)< (∆rH3)
The enthalpy of ionisation decreases when moving in a group from top to bottom.
The enthalpy of ionization increases when traveling in a time from left to right.
Periodicity:
In Period-The enthalpy contribution from the electron decreases in a time from left to right.
In Group- The enthalpy of the electron production decreases from top to bottom in a group.
Electronegativity: "In a molecule the electro-negtivity is called the relative tendency of an atom to draw the mutual pair of electrons towards itself.
Periodicity:
Over Time- The electro-negtivity increases in a period from left to right.
In Group-The electro-negative diminishes in a group from top to bottom.
Valence Electrons: The electrons in outermost shell are referred to as valence electron. Because in the outermost shell the electrons decide the valence of an element.
Valency Of An Element: The number of hydrogen or halogen atom, or double the amount of oxygen atom, which is taken as its valence when paired with one atom of the element. According to the electronic valence principle, "the number of electrons that an atom loses or receives or shares with other atoms in order to achieve the noble gas configuration is called its valence."
Periodicity:
In Period- The valence increases first, and then decreases in a time from left to right.
In Group- The valence remains constant within a group from top to bottom.
Electropositive or Metallic Character: The propensity of an element to lose electrons is called electropositive or metallic character and forms positive ions (cations). Elements with lower ionization energies have a greater tendency to lose electrons, so in their behavior they are electropositive or metallic.
The most strongly electropositive elements are the alkali metals.
Periodicity: In Period-Electropositive or metallic characters diminish in a time from left to right.
In Group- The electropositive or metallic characters in a group are rising from top to bottom.
Electro-Negative or Non- Metallic Characters:
An element's propensity to admit electrons into forming anion is called its non-metallic or electronegative character. The high electro-negativity elements have a higher tendency to gain anion forms and electrons. In the upper right portion of the periodic table, therefore, the elements are electro-negative or non-metallic in nature.
Periodicity:
In period-The electro-negative or non - metallic characters increase in a certain time from left to right.
In Group- The electro-negative or non - metallic characters in a group diminish from top to bottom.
Reactivity of Metals:
Periodicity:
In period- The tendency of an element to lose electrons decreases in period. So, the reactivity of metals decreases from left to right in a period.
In group-In a period, an element 's tendency to lose electrons increases. Thus the reactivity of metals in a group increases from top to bottom.
Reactivity of Non- Metals:
In Period- An element 's tendency to gain electrons increases over time. So the non - metals reactivity increases in a time from left to right.
In Group- The tendency of an element to gain electrons decreases in a group. So, the reactivity of non-metals increases from top to bottom in a group.
Solubility of Alkali Metals Carbonates And Bicarbonates:
Periodicity in Group: Alkali metal carbonates and bicarbonates solubility in water increases down the group (from Lithium to Caesium).
Solubility of Alkaline Earth Metal Hydroxides And Sulphates:
Periodicity in Group: Alkaline earth metal hydroxide and sulphates solubility in water is increasing down the group (from Beryllium to Barium).
Basic Strength of Alkaline Earth Metal Hydroxides:
Periodicity in Group: The basic strength in water of alkaline earth metal hydroxide increases the group (from Beryllium to Barium), i.e.,
Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2
Basic strength increases…………………………………………………………………………………….>>>>>>>
Thermal Stability Of Carbonates Of Alkali And Alkaline Earth Metals:
With the exception of lithium carbonate (LiCO3), all other alkali metals ' carbonates are heat stable, i.e. alkali metal carbonates (except LiCO3) do not decompose when heated. LiCO3 decomposes to give lithium oxide (LiCO3) when heated.
The alkaline earth metal carbonates are relatively lower in stability. They decompose on heating to give corresponding oxide and carbon dioxide. The temperature of decomposition for alkaline earth metal carbonates increases as we descend the group.
Anomalous Properties of Second Period Elements
A anomalous behavior is due in a valence shell to their small size, large charge / radius ratio, high electro negativity, non - availability of d-orbitals. The first member of each group of p - Block elements shows greater ability to form several pp-pp bonds to itself (e.g. C= C, CBC O= O, NB) and other period elements (e.g. C= O, CBC, N= O) compared to subsequent group members.
FAQs on Classification of Elements and Periodicity in Properties of Elements
1. What is the fundamental basis for classifying elements in the Modern Periodic Table?
The fundamental basis for the classification of elements in the Modern Periodic Table is their atomic number (Z). Unlike earlier classifications that used atomic mass, the modern periodic law states that the physical and chemical properties of the elements are a periodic function of their atomic numbers. This arrangement resolves anomalies found in Mendeleev's table and provides a more consistent framework based on electronic configuration.
2. Why is the classification of elements important for studying chemistry?
The classification of elements is crucial because it systematises the study of the 118+ known elements. Instead of memorising the properties of each element individually, classification allows us to:
- Study elements systematically by grouping them with similar properties.
- Predict the properties of newly discovered elements based on their position in the periodic table.
- Understand relationships and trends in properties like atomic size, electronegativity, and metallic character across periods and groups.
- Simplify the study of compounds formed by these elements.
3. What is meant by the periodicity of properties, and what is its underlying cause?
Periodicity of properties refers to the recurrence of similar physical and chemical properties of elements at regular intervals when they are arranged in increasing order of their atomic numbers. The underlying cause of periodicity is the repetition of a similar outer shell electronic configuration after certain regular intervals. For example, all alkali metals (Group 1) have one electron in their outermost shell (ns¹), which is why they exhibit similar chemical properties.
4. How are elements in the periodic table categorised into s, p, d, and f blocks?
Elements are categorised into blocks based on the type of orbital into which the last electron, or differentiating electron, enters.
- s-Block Elements: The last electron enters the outermost s-orbital. These include Groups 1 and 2.
- p-Block Elements: The last electron enters any of the three p-orbitals of their outermost shell. These include Groups 13 to 18.
- d-Block Elements: The last electron enters any of the five d-orbitals of the penultimate (second to last) shell. These are the transition metals in Groups 3 to 12.
- f-Block Elements: The last electron enters any of the seven f-orbitals of the anti-penultimate (third to last) shell. These are the lanthanoids and actinoids placed separately at the bottom.
5. Explain the general trends for atomic radius and ionization enthalpy across a period and down a group.
Atomic Radius:
- Across a Period (left to right): The atomic radius generally decreases. This is because the nuclear charge increases, pulling the electrons in the same shell closer to the nucleus.
- Down a Group: The atomic radius increases. This is due to the addition of a new electron shell with each successive element, which outweighs the increase in nuclear charge.
- Across a Period (left to right): The ionization enthalpy generally increases. With decreasing atomic size and increasing nuclear charge, it becomes more difficult to remove the outermost electron.
- Down a Group: The ionization enthalpy decreases. The outermost electron is farther from the nucleus and shielded by inner electrons, making it easier to remove.
6. Why is the first ionization enthalpy of nitrogen significantly higher than that of oxygen?
The first ionization enthalpy of nitrogen is higher than that of oxygen, despite oxygen being further to the right in the period. This is an important exception due to electronic configuration. Nitrogen has a half-filled p-orbital (1s² 2s² 2p³), which is a particularly stable configuration. Removing an electron from this stable state requires more energy. In contrast, oxygen (1s² 2s² 2p⁴) can achieve a stable half-filled configuration by losing one electron, a process which requires comparatively less energy.
7. How does the screening effect (or shielding effect) influence periodic trends?
The screening effect is the phenomenon where the inner-shell electrons shield or block the outer-shell electrons from the full attractive force of the nucleus. This effect is crucial in determining periodic trends:
- Atomic Radius: Down a group, the addition of new shells increases the screening effect, which pushes the valence electrons farther away and increases the atomic radius.
- Ionization Enthalpy: A stronger screening effect makes it easier to remove a valence electron because it feels a weaker pull from the nucleus. Therefore, as screening increases down a group, the ionization enthalpy decreases.
8. What are isoelectronic species, and how do their radii compare?
Isoelectronic species are atoms or ions that have the same number of electrons. For example, N³⁻, O²⁻, F⁻, Ne, Na⁺, and Mg²⁺ are all isoelectronic because they each have 10 electrons. When comparing their radii, the species with the highest nuclear charge (more protons) will be the smallest, as its nucleus pulls the same number of electrons more strongly. Therefore, for this series, the radius decreases as follows: N³⁻ > O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺.
