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What are Elements’ Properties? Examples and Definitions
Periodic table displays chemical elements according to many features like configuration , affinity etc
Periods are the rows, and the periodic table has seven of them. Metals can be found on the left rows, whereas nonmetals can be found on the right. The columns, on the other hand, are referred to as groups. Elements in groups exhibit a variety of chemical properties. There are 18 groups, with halogens being found in group 17 and noble gasses being found in group 18.
Periodic Trends
Periodic trends are defined as distinct patterns in the properties of chemical elements found in the periodic table. The following are the major trends:
The Ionization Energy
Metal Characteristics
Atomic Radius
Electronegativity
Ionic radius,
Affinity for electrons
Reactivity of chemicals
Shielding effect
Changes in the structure of atoms of the elements within their groups and periods cause these patterns. There are a few exclusions, such as the ionisation energy of groups 3 and 6.
Periodic Law
Periodic trends are founded on periodic law. “The chemical elements are listed in order of increasing atomic number, and main properties undergo cyclic changes,” according to periodic law. In regular periods, elements with comparable chemical characteristics reappear.”
Dmitri Mendeleev established this principle. He also asserted that the periodic chart was based on different physical and chemical properties of elements, not merely atomic weights
The recurrence of features was later discovered to be related to the recurrence of comparable electronic configurations in the outer shells of atoms.
- Energy of Ionization
The ionisation potential is defined as follows:
“Minimum energy required by an isolated atom in its neutral or gaseous state to remove one electron”
The ionisation energy grows as one progresses through the interval. The reason for this is that when the nuclear charge increases over time, the nucleus holds the electrons more tightly.
However, as one moves down the group, the ionisation energy diminishes. The reason for this is that as the nuclear charge lowers, the valence electrons get farther away from the nucleus
Ionisation energy is affected by a variety of factors.
The Ionization Energy Levels are affected by a number of factors.
- Nuclear Charges
The force of attraction between the nucleus and valence electrons decreases as the nuclear charge decreases, resulting in decreased ionisation energy.
- Shielding Effect
Shielding effect increases as nuclear charge grows, hence as shielding effect increases, so does ionisation energy.
- Radius of an atom
The force of attraction between the nucleus and valence electrons reduces as the atomic radius grows. When a result, as the atomic radius increases, the ionisation reduces
- Valence Shells, Half-Filled
Ionisation energy is high in pseudo-filled or half-filled valence shells.
A basic rule to remember is that if the primary quantum number is low, the ionisation number for the electron in that shell will be high.
Exceptions
The periodic trend does not apply to any of the elements in the oxygen and boron families. They take a little less energy than the standard trend.
- Metallic Characteristics
The capacity of an element to conduct electricity is known as its metallic property. As the nuclear charge lowers down the group, the metallic characteristics rise. The valence electrons are able to conduct electricity well because they are loosely bound by the nucleus.
However, when nuclear charge accumulates, the metallic property reduces over time. The valence electrons’ force of attraction with the nucleus rises as a result, preventing them from conducting electricity or heat.
3.Atomic Radius
The atomic radius is the distance between the nucleus of an atom and its outermost stable electron orbital when it is in equilibrium. As the nuclear charge grows, the atomic radius decreases with time. The reason for the drop is because as nuclear charge rises, the force of attraction between the nucleus and the valence electrons rises as well, causing the nucleus to tightly grip the electron, resulting in smaller atomic radii.
The atomic radius of a group grows as it progresses. The reason for this is that as new shells are added, the nuclear charge lowers. However, the atomic radii expand diagonally as well, generating some
Example
Along the Period – Li > Be > B > C > N > O > F
Down the Group – Li < Na < K < Rb < Cs
4)Electronegativity
The ability of an atom or molecule to attract a pair of electrons is known as electronegativity. The difference in electronegativity of the atoms determines the bond formed as a result of this.
Electronegativity increases as nuclear charge grows over time. The electronegativity falls as the nuclear charge lowers as you move down a group. Because the distance between the atom’s nucleus and the valence electrons is so great, the electrons are easily lost.
Example
Along the Period- Li < Be < B < C < N < O < F
Down the Group – Li > Na > K > Rb > Cs
Exception
The group 13 elements are an exception, and electronegativity rises from aluminum to thallium as a result. In addition, tin has a stronger electronegativity than lead in group 14.
5)Affinity for Electrons
The tendency of an atom to accept an electron or an electron pair is known as electron affinity. As nonmetals gain electrons and form anions, this is a characteristic property. The electron affinity rises with time as the nuclear charge increases.
As the nuclear charge drops, it decreases along the group. Noble gasses are not included and fluorine has the highest electronegativity. Because they have a complete valence shell, they cannot gain or lose electrons.
6.)Shielding Effect
It can be characterised as the inner electrons resisting an outer electron. It can also be used to describe how many nuclei are capable of controlling electrons in the outer shell. Because of the increasing shielding effect, the effective nuclear charge drops along the group. The effective nuclear charge rises over time as the nuclear charge rises.
7)Atomic Radius
The electrons in an ion’s many shells and the nucleus make up the ion. The ionic radius of an ion is the distance between the nucleus and the electron in the ion’s last outermost shell. There is a tendency that may be seen in the periodic table based on the ionic radius of different elements. This pattern can be summarised as follows:
The ionic radius of the elements increases in value as we proceed down the periodic table from top to bottom. This occurs because the number of layers or shells of electrons increases as we proceed down the periodic table.
The ionic radius tends to shrink as we move sideways from left to right on the periodic table. Although it appears strange that as additional protons, electrons, and neutrons are added, the ionic size decreases. However, this occurs because metals shed their outer electron layers as we move sideways on the periodic table in order to generate cations. When the number of electrons in an ion exceeds the number of protons, the ionic radius increases, resulting in a considerable decrease in nuclear charge.
This tendency is true not only for ionic radius but also for atomic radius; nonetheless, the two are distinct.
8)Chemical Reactivity
The capacity of an atom to react with any other substance is referred to as its reactivity. Ionisation energy and electronegativity are frequently used to govern chemical reactivity. The chemical reactivity trend in the periodic table is based on this mechanism of electron transfer and interchanging.
Chemical reactivity of metals reduces as we move from left to right on the periodic table. As we proceed from the top to the bottom groups of the table, however, the reactivity increases. The interchange of electrons increases easier and faster as we move downhill or to the left, increasing the chemical reactivity of the elements.
In non-metals, the situation is reversed. As we proceed from the left to the right side of the table, the chemical reactivity increases. As we proceed from the top groupings to the lowest groups, the reactivity drops. The easier it is for atoms to shed their electrons in exchange for other electrons as they move upwards or to the right, the higher the electronegativity, which increases the chemical reactivity of the elements.
CONCLUSION:–
Certain periodic tendencies in element properties are shown by the systematic arrangement of elements in a periodic table. Atomic and ionic radii, for example, decrease in a period from left to right. Understanding the patterns in fundamental properties of elements like atomic, ionisation enthalpy etc will lead you to the conclusion that property periodicity is mainly determined by an element’s electronic configuration.
