Explanation of Group 16 Elements Electronic Configuration
The electronic configuration of any element is defined as the arrangement of the electrons around the nucleus. The electronic configuration of any element determines its physical state and reactivity with other elements. When one looks at the Group 16 elements, the electronic configuration of all the elements in that group is categorized by the presence of six electrons in their last shell or the valence shell.
The elements present in group 16 consist of oxygen (0), sulfur (S), selenium (Se), tellurium (Te), and polonium (Po). Of all the elements, only Polonium is radioactive. All these elements can exist in a free state in nature. However, due to its electronic configuration, it can react with other elements and also exists in a combined state.
As stated earlier, electronic configuration refers to the arrangement of electrons on its orbital shells and subshells. All elements in group 16 have six electrons in its last shell; for example, the total number of electrons for oxygen is 8, which is distributed in two shells as 2 and 8. To understand the electronic distribution for other members of group 16, it is important to learn some basic principles to do so.
The first rule is to fill the lower energy shells with electrons first before moving to the higher shells. Hund’s rule, Pauli’s exclusion principle, and Aufbau’s rule are needed to be followed while distributing the electrons. According to Pauli’s exclusion principle, no two electrons in the same atom can not have the same quantum numbers (n,l,m, and s); the first three might have the same quantum number, but it will differ from the fourth value.
According to Hund’s rule, similar energy orbitals accommodate one electron, and then other electrons can pair with them in half-filled orbitals. According to Aufbau’s principle, electrons first occupy the lowest energy levels. All these three principles can be followed while determining the electronic configuration of an atom.
Atomic Orbital Diagonal Rule
Considering all the three above mentioned principles is the best approach for deciphering the electronic configuration of any element. For example, the total number of electrons in an oxygen atom is 8. The first step is to fill the lowest energy shell 1s with two electrons. The remaining six electrons are distributed in 2s and 2p orbitals. 2s orbital will have two electrons, and 2p orbital will have four electrons.
According to Neil Bohr, all the members of the same group of the Periodic table have a similar electronic configuration. Therefore, it also stands true for all the members of group 16, and the electronic configuration of oxygen follows the pattern of the general electronic configuration of group 16 elements, which is ns2 np4.
Since the noble gas is considered to have a complete electronic shell, the electronic configuration of most elements is represented in terms of its nearest noble gas.
It is important to note that the electronic configuration of any element determines its chemical properties. Group 16 electron configuration indicates that its members of Group 16 have six elements in its valence shell, and therefore require two elements to complete the octet valency. Therefore all the elements of group 16 are negatively charged since it can receive two electrons from other elements.
These anions can interact with positively charged cations that can donate electrons so that their octet is also completed along with these anions. For example, oxygen receives two electrons, one from each hydrogen atom to form water. Sulfur receives two electrons, one from each hydrogen atom to form hydrogen sulfide gas.
About Electronic Configuration
Students have been introduced to the atomic structure of elements in the previous classes. We all know that Atoms consist of electrons, protons and neutrons. In this class, we will get to learn about the arrangement of these subatomic particles inside the atoms of any particular element.
The rules and principles used for determining the position of all electrons are studied under the Physical Chemistry of Class 11 NCERT textbooks. It also mentions the story of various scientists working to understand and identify the difference between various types of materials available on Earth and their properties. It is the story of the gradual development of chemistry as a subject of study.
The electrons are the freely moving subatomic particles revolving around a nucleus centre consisting of protons and neutrons. The degree of freedom of electrons varies according to the orbit they choose for their position. Several protons present in the atoms of an element are known as the 'Atomic Number' of the element.
This Atomic number also gives us information about the number of electrons present in the atom which are equal in number to balance the net positive charge of protons. The Periodic Table arranges the elements in the increasing order of their atomic number and categorizes all elements into different groups. Neil Bohr suggested that the elements of the same group have similar electronic configurations.
For example, the atoms Hydrogen and Helium with single and double electrons respectively occupy the lowest valency positions. The rule and principle to arrange the electrons by assigning different orbits around the nucleus are known as the Electronic configuration of atoms. As the chemical properties of the elements depend on the number of outermost free electrons. We all know that the properties of the elements under the same group are the same so the number of electrons in the outermost cell is equal. The Elements of group 16 such as Oxygen, Sulfur, Selenium, Tellurium, Polonium and all have 6 electrons on the outermost cell. The 's' valence shells always contain 2 electrons, The 'p' valence shell contains 4 electrons.
To know the significance of the electronic configuration of the Group 16 elements, log into Vedantu and seek deeper insights from the expert mentors. Study and understand the concepts of this topic to answer critical questions easily.
FAQs on Electronic Configuration of Group 16 Elements
1. What is the general valence shell electronic configuration that defines the Group 16 elements?
The general electronic configuration for the valence shell of Group 16 elements is ns²np⁴. This means that every element in this group has six electrons in its outermost shell, which determines their chemical properties and their position in the periodic table.
2. Which elements are included in Group 16 of the periodic table?
Group 16, also known as the chalcogens or the oxygen family, consists of the following elements:
- Oxygen (O)
- Sulfur (S)
- Selenium (Se)
- Tellurium (Te)
- Polonium (Po)
3. As an example, what is the complete electronic configuration for Sulfur (S)?
Sulfur has an atomic number of 16. Its complete electronic configuration is written as 1s² 2s² 2p⁶ 3s² 3p⁴. The condensed, or noble gas, configuration is [Ne] 3s² 3p⁴, which clearly shows its 6 valence electrons in the third energy level.
4. How do fundamental principles like the Aufbau principle and Hund's rule apply to the electronic configuration of a Group 16 element like Oxygen?
For Oxygen (atomic number 8), the principles are applied as follows:
- Aufbau Principle: Electrons fill the lowest energy orbitals first. They go into 1s, then 2s, and finally 2p.
- Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons with opposite spins. Thus, the 1s and 2s orbitals are filled (1s², 2s²).
- Hund's Rule: When filling degenerate orbitals like the three p-orbitals, electrons are placed in separate orbitals first before any are paired up. For Oxygen's four 2p electrons, one electron goes into each of the 2pₓ, 2pᵧ, and 2p₂ orbitals, and the fourth electron then pairs up in the 2pₓ orbital.
5. Why do Group 16 elements have a strong tendency to form anions with a -2 charge?
Group 16 elements have six valence electrons (ns²np⁴). By gaining two more electrons, they can achieve a highly stable, completely filled outer shell, identical to the electronic configuration of a noble gas (ns²np⁶). This stable octet is an energetically favourable state, which is why these elements are highly electronegative and readily form anions with a -2 charge, such as the oxide (O²⁻) or sulfide (S²⁻) ions.
6. How does the electronic configuration of Group 16 explain its variety of oxidation states, such as -2, +4, and +6?
The various oxidation states are a direct result of their ns²np⁴ configuration:
- The -2 oxidation state is the most common, achieved by gaining two electrons to complete the octet.
- The +4 oxidation state occurs when the element loses or shares its four p-electrons (np⁴).
- The +6 oxidation state is shown when the element loses or shares all six of its valence electrons (both ns² and np⁴).
Oxygen, due to its high electronegativity and lack of d-orbitals, primarily shows the -2 state. For heavier elements like Sulfur and Selenium, the +4 and +6 states become more common in compounds with highly electronegative elements like oxygen and fluorine.
7. Which element in Group 16 is radioactive, and what is its atomic number?
The radioactive element in Group 16 is Polonium (Po). Its atomic number is 84. All isotopes of Polonium are radioactive.
8. How does the electronic configuration of Group 16 differ from its neighbours, Group 15 and Group 17?
The primary difference lies in the number of electrons in their valence p-orbitals:
- Group 15 (Pnicogens): Have a configuration of ns²np³. They are three electrons short of a stable octet.
- Group 16 (Chalcogens): Have a configuration of ns²np⁴. They are two electrons short of a stable octet.
- Group 17 (Halogens): Have a configuration of ns²np⁵. They are just one electron short of a stable octet.
This difference explains why halogens are generally more reactive and electronegative than chalcogens, as they only need to gain one electron to achieve stability.
9. Based on its electronic structure, why does Oxygen show properties that are different from other Group 16 elements?
Oxygen exhibits anomalous behaviour primarily due to three factors related to its electronic configuration and structure:
- Small Atomic Size: It is significantly smaller than Sulfur and other elements in the group.
- High Electronegativity: It is the second most electronegative element after Fluorine.
- Absence of d-orbitals: Unlike S, Se, and Te, Oxygen has no available d-orbitals in its valence shell. This limits its covalency and prevents it from forming compounds in higher oxidation states like +4 and +6.
This also allows Oxygen to form strong pπ-pπ multiple bonds, resulting in the stable diatomic molecule O₂, whereas other elements in the group are solids.
