An Introduction to Modern Periodic Table
In the modern periodic table, all the elements are arranged systematically. They are assigned positions in the groups and periods according to the atomic number and the similarities in the physical and Chemical properties. The modern periodic table was formed after the introduction of the modern periodic law. This law was given by Henry Moseley. According to this law, the atomic number decides the properties of the elements. In simple words, the properties of the elements repeat themselves after regular intervals if the elements are arranged in the increasing order of their atomic number.
All the elements are arranged in groups and periods in the periodic table. Groups are vertical columns and these are 18 in number. The horizontal rows are the periods and these are 7 in number.
Elements of the Modern Periodic Table
There are different types of elements present in the modern periodic table. These elements are listed below:
Elements of Main Group - In the modern periodic table, the Group 1 and 2 elements on the left-hand side and the elements of the Group 13 to 17 on the right-hand side are known as the main group elements. These are also called representative elements.
Noble Gases - The elements of the Group 18 are called noble gases. These are not reactive as the outermost shell of these elements is completely filled.
Transition Elements - The elements belonging to Groups 3 to 12 are known as transition elements. They are present in the middle block of the periodic table. The two outermost shells of these elements are not completely filled.
Inner Transition Elements - There are two rows of inner transition elements and each row comprises 14 elements. The first row is called the lanthanide series (elements from atomic number 58 to 71). The second row is known as the actinide series (elements from atomic number 90 to 103).
Metals - Metals are placed on the left-hand side of the modern periodic table. Elements of Group 1 and 2 are metals. Group 1 elements are called alkali metals and Group 2 elements are called alkali earth metals.
Non-Metals - The space on the right-hand side of the modern periodic table is occupied by the non-metals.
Metalloids - Metalloids are placed diagonally starting from boron (Group 13) to polonium (Group 16). These elements possess the properties of both metals and non-metals.
Periodic Trends in the Modern Periodic Table
Periodic table trends are systematic patterns in the properties of elements of the periodic table. Periodic trends depend upon the electronic configuration of elements. They arise from the changes in the atomic structure of elements with their respective periods(horizontal rows) and groups(vertical columns) in the periodic table. Periodic trends are based on the periodic law, which states that if the Chemical elements are listed in order of increasing atomic number(from left to right), then the physical and Chemical properties of elements will be a periodic function of their atomic numbers. We will discuss some trends of fundamental properties of elements here. ‘
Number of Shells
The number of shells increases with the increase in atomic number. The period number indicates the number of shells of the related element, so if we move from top to bottom in a group, the number of shells increases by 1 and remains the same if we move left to right across a period.
The number of electrons present in the outermost shell is called valence electrons. If we move from top to bottom in a group, the number of valence electrons remains the same. While moving across a period, the valence electrons increase.
Valency is the combining capacity of elements or the number of electrons donated, accepted or shared during compound formation by the elements. As valence electrons remain the same on moving top to bottom, so the valency also remains the same on going down the group. While we move across a period, valency first increases and then decreases. In order to become Chemically stable, the octet rule must be followed with the least energy loss. Due to energy loss, elements accept or donate a minimum number of electrons.
Valence Electrons and Valency
The valency of an element refers to the number of electrons that are lost or gained by it to have a stable configuration. The valencies of elements belonging to the s and p block are calculated as 8-number of valence electrons.
The valencies of elements belonging to the d and f block are determined by the d and f orbital electrons. The general valency of the d and f block elements is 2 and 3 respectively.
Atomic Size (Atomic Radius Periodic Table)
Atomic size refers to the distance from the atomic nucleus to the outermost electron orbit. Atomic size usually increases on going down in a group as the number of shells also increases. But, it decreases on moving left to right across the period due to the shrinking of the atom because of increasing nuclear force.
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Metallic & Non - Metallic Character
Metals have a tendency to lose electrons. Metals are the best donors while non-metals are acceptors. There are trends in metallic character as we move across the period, the metallic character decreases from left to right as the valency of elements decreases. Atoms more readily accept electrons to fill valence shells than lose them to remove the unfilled shell.
Metallic character increases on going down to the group because of the lesser nucleus force, the electrons become easier to lose as atomic radius increases. Metals are on the left-hand side of the periodic table except for hydrogen.’
Melting Point
The melting of an element is defined as the amount of energy required to change the solid phase of the element into the liquid phase. The change of the phase takes place due to the cleavage of the bonds between the atoms. If the bond will be stronger then more energy will be required to break that bond. The energy required is directly proportional to the temperature. High bond dissociation energy means high temperature. The melting point varies across the periodic table, therefore, it is not considered as the distinguishable trend. But following conclusions can be derived from the periodic table using the trend of melting point.
Generally, all the metals have high melting points.
Mostly, all the non-metals have low melting points.
There is an exception in the case of non-metals. The carbon atom has a high melting point. Also, boron, which is a semi-metal, has a high melting point.
Ionization Enthalpy/ Ionization Energy
It is the minimum amount of energy required to remove the outermost electron or the most loosely bound electron. On moving left to right in a period, atomic size decreases so the nucleus force increases. Due to this, across a period in the periodic table, ionization energy usually increases.
When moving down the group, the number of shells increases. Due to this, the effective nuclear charge reduces. Atoms lose electrons easily, so less energy is required to donate electrons.
Thus, ionization energy decreases on moving top to bottom.
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Electron Affinity
Electron affinity is the energy released by an atom when an electron is added to it. Electron affinity generally decreases on going down a group of elements because each atom is larger than the atom above it. With the larger distance between the nucleus and the outermost electron, the nucleus force decreases. On moving left to right across a period, the nuclear force increases, so electron affinity also increases, caused by the decrease in atomic radius.
Electronegativity
Electronegativity is a property that describes an atom's ability to attract and bind with electrons. If the valence shell of an atom is less than half full, It requires less energy to lose an electron than to gain one. While if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate it. Sorry, electronegativity increases on moving left to right across the period.
Atomic number increases down the group so atomic size also increases. Thus, electronegativity decreases on moving top to bottom in a group.
FAQs on Modern Periodic Table Trend
1. What are the key periodic trends in the properties of elements in the modern periodic table?
The properties of elements show gradual and predictable changes, known as periodic trends, across periods and down groups. The main trends studied are:
- Atomic Radius: The size of an atom.
- Ionisation Enthalpy: The energy required to remove an electron from an isolated gaseous atom.
- Electron Gain Enthalpy: The energy change when an electron is added to an isolated gaseous atom.
- Electronegativity: The ability of an atom in a chemical compound to attract a shared pair of electrons.
- Metallic and Non-metallic Character: The tendency of an element to lose or gain electrons, respectively.
2. How does the atomic radius of elements change as we move across a period and down a group?
The atomic radius shows two distinct trends in the modern periodic table. When moving from left to right across a period, the atomic radius generally decreases. This is because electrons are added to the same valence shell, while the nuclear charge increases, pulling the electron cloud closer to the nucleus. Conversely, when moving from top to bottom down a group, the atomic radius increases as a new electron shell is added for each successive element.
3. Why does ionisation enthalpy generally increase across a period but decrease down a group?
The trend in ionisation enthalpy is a direct consequence of the changes in atomic size and nuclear charge.
- Across a period (left to right): The atomic size decreases and the effective nuclear charge increases. This holds the valence electrons more tightly, making it harder to remove them. Therefore, ionisation enthalpy increases.
- Down a group (top to bottom): The atomic size increases due to the addition of new shells. The outermost electrons are farther from the nucleus and experience increased shielding from inner electrons, which reduces the nucleus's pull. This makes it easier to remove an electron, so ionisation enthalpy decreases.
4. What is the difference between electron gain enthalpy and electronegativity?
While both properties relate to an atom's attraction for electrons, they are fundamentally different. Electron gain enthalpy is a measurable, absolute quantity representing the energy change (often released, hence negative) when a single electron is added to an isolated gaseous atom. In contrast, electronegativity is a relative, unitless value that describes the tendency of an atom to attract a shared pair of electrons within a covalent bond. It is not measured directly but calculated based on other atomic properties.
5. How can we predict the metallic or non-metallic character of an element based on its position in the periodic table?
An element's position provides clear clues about its character. Metallic character refers to the tendency to lose electrons, which is highest for elements on the left side and at the bottom of the periodic table. Therefore, metallic character increases down a group and decreases across a period. Non-metallic character is the tendency to gain electrons, which is strongest for elements on the top-right side of the table (excluding noble gases). Thus, non-metallic character increases across a period and decreases down a group.
6. Why do noble gases have exceptionally high ionisation enthalpies and near-zero electronegativity?
Noble gases, located in Group 18, possess a completely filled valence shell, resulting in a highly stable electronic configuration. This stability means they have very little tendency to alter their electronic state. Consequently, a very large amount of energy is required to remove an electron, leading to an exceptionally high ionisation enthalpy. For the same reason, they have no tendency to attract a shared pair of electrons in a bond, giving them an electronegativity value of or close to zero.
7. Explain the anomaly in the first ionisation enthalpy of Beryllium (Be) and Boron (B).
According to the general trend, Boron (Group 13) should have a higher first ionisation enthalpy than Beryllium (Group 2) because it is further to the right in the same period. However, the opposite is true. This is an important exception explained by their electronic configurations. Beryllium (1s²2s²) has a fully-filled 2s orbital, which is a stable arrangement. Boron (1s²2s²2p¹) has a single electron in the 2p orbital. It is easier to remove this single 2p electron than to remove an electron from the stable, fully-filled 2s orbital of Beryllium. Therefore, Beryllium has a higher first ionisation enthalpy than Boron.
