Oxidation Number or Oxidation State

OXIDATION NUMBER OR OXIDATION STATE This fundamental concept pertains to the numeric value that signifies the quantity of electrons an atom has either acquired or lost as it undergoes a transition within a compound from its isolated, unbound state. This value, often referred to as the "oxidation number" or "oxidation state," plays a central role in understanding chemical reactions and the behavior of atoms within compounds. The oxidation number serves as a quantitative measure of the electron distribution around an atom in a compound. It indicates whether the atom has gained or lost electrons in comparison to its neutral, free state. Oxidation numbers are crucial for tracking the flow of electrons in redox (reduction-oxidation) reactions. When an atom's oxidation number increases, it signifies electron loss, while a decrease in the oxidation number indicates electron gain. This concept is vital in determining how atoms interact in chemical reactions and how electrons are transferred, a fundamental aspect of the study of chemistry. The oxidation number provides a quantitative understanding of the electron changes that occur during chemical reactions and enables chemists to predict and interpret the behavior of atoms within compounds. It is a foundational concept that underpins many aspects of chemical science and facilitates the comprehension of electron transfer in various chemical processes. (a) When an atom loses electrons during compound formation, its oxidation number is indicated with a positive (+) sign. Conversely, if an atom gains electrons in the formation of a compound, its oxidation number is denoted with a negative (-) sign. (b) Oxidation numbers represent the actual charge in the context of ionic compounds and an (c) imagined charge in the case of covalent compounds. The maximum oxidation number of an element equals its group number on the periodic table. (d) (e) Minimum oxidation no. of an element is equal to group no. 8. IA group elements always show +1 oxidation no. group elements always shows +2 oxidation no. group elements show +3 oxidation no. but +1 becomes more stable going down the group (due to inert pair effect) group shows-4 to +4 oxidation no. IV A VA group shows-3 to +5 oxidation no. VI A group shows-2 to +6 oxidation no. VII A group shows -1 to +7 oxidation no. Inert gases show zero oxidation no. Oxidation Number or Oxidation State Determination Oxidation No. for Coordinate Bond (a) In compounds where a coordinate bond is established, a specific pattern emerges in terms of the oxidation numbers of the elements involved. When such a bond forms between a low electronegative element (the electron donor) and a high electronegative element (the electron acceptor), there is a distinct shift in the oxidation numbers: The element that acts as the electron donor exhibits an oxidation number of +2. This positive oxidation number signifies the loss of two electrons during the formation of the coordinate bond. Essentially, this element is undergoing oxidation, which involves losing electrons. Conversely, the element that functions as the electron acceptor displays an oxidation number of-2. This negative oxidation number reflects the gain of two electrons during the coordinate bond formation. In this case, the element is experiencing reduction, which entails gaining electrons. This phenomenon is a crucial aspect of understanding the electron transfer and oxidation-reduction (redox) processes that transpire in the creation of coordinate bonds. It exemplifies how electrons are shifted between elements, resulting in specific changes in their oxidation numbers, and it plays a significant role in elucidating the behavior of bonded compounds where such coordinate bonds are prevalent. For example, in H2SO4 H-O-S-O-H Here 'S' is low electronegative element than 0. therefore, number of S=+2 and O. N. of O=-2 (b) In compounds where a coordinate bond is established between two identical, equally electronegative elements, a specific pattern emerges in terms of the oxidation numbers of these elements: The element that acts as the electron donor exhibits an oxidation number of +2. This positive oxidation. number signifies that during the formation of the coordinate bond, this element essentially loses two electrons. This process is akin to oxidation, where electrons are lost. On the other hand, the element that functions as the electron acceptor displays an oxidation number of-2. This negative oxidation number reflects the gain of two electrons during the formation of the coordinate bond. This element undergoes a reduction process, signifying that it gains electrons. This phenomenon showcases a key aspect of electron transfer and oxidation-reduction (redox) processes within compounds that feature coordinate bonds. It highlights how electrons are transferred between the identical, electronegative elements, leading to specific changes in their oxidation numbers. Understanding this behavior is pivotal for comprehending the intricacies of bonded compounds where coordinate bonds are prevalent. For example: In Na2S2O3 NaO-S-ONa Here O.N. of 'S' is +2, Because it is e donor and the other 'S' is -2, Because it is e acceptor. (C) The oxidation state or oxidation number of an atom is contingent upon several factors, particularly its electronic configuration and its relationship to the periodic properties of elements. Oxidation State Determination The oxidation state of an atom is essentially its oxidation number per atom and is commonly referred to as such in practice. It is frequently represented as the oxidation number. The rules to derive oxidation number or oxidation state (a) The O.S. of an element in its free state is zero. Example: O.S. s of Na, Cu, 12, Cl2, 02 etc. are zero (b) Sum of O.S's of all the atoms in neutral molecule is zero. (c) Sum of O.S' s of all the atoms in a complex ion is equal number of charges present on it. (d) In complex compounds, O.S. of some neutral molecules (ligands) is zero. Example: CO, NO, NH3, H2O. (e) Generally, O.S. of oxygen is-2 but in H202 it is -1 and in OF2 it is +2. (f) Generally, O.S. of Hydrogen is +1 but in metallic hydrides it is -1. (g) Generally, O.S. of halogen atoms is-1 but in interhalogen compounds it changes. NOTE: Sometimes same atom in a compound has different O.S. For example, structure of Na2S2O3 is Na-O-S1-O-Na 0 Here S₁ and S2 both are sulphur atoms but they have different O.S. O.S. of O.S. of S1=6 S2-2 (it is accepting two electrons from S₁) Average O.S. of S=6-2-2 2 (h) Generally, O.S. of alkali metals is +1 and that of alkaline earth metals is +2. O.S. of transition elements vary from compound to compound. (1) Mn has O.S. from +1 to +7. Mn2O+1,MnO +2,Mn304-8 3.MnO2+4,Mn205+5,Mn 4+6.MnO4+7 (j) O.S. of an atom may be fractional, negative, zero as well as Positive. Oxidation State as A Periodic Property The oxidation state or oxidation number of an atom is contingent upon several factors, particularly its electronic configuration and its relationship to the periodic properties of elements. (a) 1 A group or alkali metals shows +1 oxidation state. (b) II A group or alkaline earth metals show +2 O.S. (c) The maximum normal oxidation state, show by III A group elements is +3. These elements also show +2 to +1 oxidation states also. (d) Elements of IVA group show their max & min. oxidation states +4 &-4 respectively. (e) Non-metals shows number of oxidation states, the relation between max & min, oxidation states for non-metals is equal to maximum O.S. -minimum O.S. = 8 For example, sulphur has maximum oxidation number +6 as being in VI A group element. Electronic Configuration The electronic configuration of an atom is determined by the distribution of its electrons within different energy levels or orbitals. This configuration plays a fundamental role in defining an element's chemical behavior and its ability to form compounds. The arrangement of electrons in an atom's valence shell, where chemical interactions primarily occur, is of particular significance. Periodic Properties Elements in the periodic table are organized based on their atomic number and electron configuration. As one moves across a period (horizontal row) from left to right, there is a general trend of increasing electronegativity. Electronegativity is the atom's tendency to attract electrons when participating in chemical reactions. It is a key factor in determining the oxidation state of an atom. Oxidation State: The oxidation state of an atom reflects its ability to lose, gain, or share electrons when forming compounds. Elements with a lower electronegativity, often found on the left side of the periodic table, tend to exhibit positive oxidation states by losing electrons. In contrast, elements with higher electronegativity, typically on the right side of the table, are more prone to taking on negative oxidation states by gaining electrons. Periodic Trends The periodic table displays trends in electronegativity, with elements becoming more electronegative as one progresses from left to right across a period. This trend influences the oxidation states elements are likely to assume in chemical reactions. For instance, alkali metals like sodium (Na) tend to exhibit an oxidation state of +1, as they readily lose one electron from their valence shell, while halogens like chlorine (CI) frequently display an oxidation state of-1, having a strong tendency to gain an additional electron. Understanding an atom's electronic configuration, its place within the periodic table, and the corresponding trends in electronegativity is crucial in determining its oxidation state. This knowledge is indispensable for predicting and comprehending chemical reactions and the formation of compounds involving various elements. Oxidising And Reducing Agent Oxidant or Oxidising Agent In the realm of chemical reactions, there are particular species that possess a distinct role. These species are capable of oxidizing other entities that coexist within a given reaction while simultaneously undergoing self-reduction. Such species are officially designated as oxidants or oxidizing agents. An alternative perspective on these species defines them as substances that partake in a reaction by accepting electrons from another species. As a consequence of this electron exchange, the species that functions as an oxidant or oxidizing agent experiences a decrease in its oxidation number throughout the course of the reaction. The role of oxidants in chemical reactions is pivotal, as they facilitate the transfer of electrons and are fundamental in redox (reduction-oxidation) processes. By accepting electrons from other species, oxidants lead to the oxidation of those entities while, in turn, undergoing reduction themselves. This electron exchange principle lies at the heart of understanding how chemical reactions progress and how electrons are transferred between species during these transformations. The concept of oxidants plays a central role in various chemical reactions and serves as a foundation for exploring the intricate dynamics of electron transfer in chemistry. Some Important Oxidising Agent or Oxidant 1. All elements with high electronegative character like N, O, F, Cl, etc. 2. All metallic oxides like Li2O, Na2O, Na2O2, CaO, MgO, BaO2 etc. 3. Some nonmetallic oxides like CO2, SO2, H202, 03. 4. All neutral compound or ion in which element shows their higher oxidation no. or state are act as oxidant or oxidising agent like KMnO4, H2SO4, SnCl4, H3PO4, K2Cr2O7, HCIO4, CuCl2, HNO3, H2505, FeCl3, HgCl2, etc. Reductant or Reducing Agent An element or compound that, in a reaction, donates electrons, causing its own oxidation number to increase while reducing another element, is termed a reductant or reducing agent. Some Important Reducing Agent or Reductant 1. All metals like, K, Mg, Ca, etc. 2. All metallic hydrides like NaH, CaH2, LiAlH4, NaBH4. AlH3, etc. 3. All hydracids like HF, HCl, HBr, H2S etc. 4. Some organic compounds like Aldehyde, formic acid, oxalic acid, tartaric acid. 5. All neutral compounds or ions, which show their lower oxidation state. MnO, HCIO, HCIO2, H3PO2. HNO2, H2SO3, FeCl2, SnCl2, Hg2Cl2, CH2Cl2 etc. Some Important compound which can act as oxidant and reductant both HNO2, SO2, H2O2, 03, Al2O3, CrO2, MnO2, ZnO, CuO, NOTE: Al2O3, CrO2, MnO2, ZnO, CuO are called as amphoteric oxide.