Chemistry
Oxidation Number
The oxidation number of an atom in a compound is a measure of the atom's ability to attract or lose electrons. It indicates the charge that an atom would have if the compound was composed of ions. The oxidation number is used to track the transfer of electrons in chemical reactions.
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10 Key excerpts on "Oxidation Number"
eBook - PDF- Morris Hein, Susan Arena, Cary Willard(Authors)
- 2016(Publication Date)
- Wiley(Publisher)
Oxidation Numbers can be zero, positive, or negative. An Oxidation Number of zero means the atom has the same num- ber of electrons assigned to it as there are in the free neutral atom. A positive oxidation num- ber means the atom has fewer electrons assigned to it than in the neutral atom, and a negative Oxidation Number means the atom has more electrons assigned to it than in the neutral atom. The Oxidation Number of an atom that has lost or gained electrons to form an ion is the same as the positive or negative charge of the ion. (TABLE 17.1.) In the ionic com- pound NaCl, the Oxidation Numbers are clearly +1 for the Na + ion and −1 for the Cl − ion. The Na + ion has one less electron than the neutral Na atom, and the Cl − ion has one more electron than the neutral Cl atom. In MgCl 2 , two electrons have transferred from the Mg atom to the two Cl atoms; the Oxidation Number of Mg is +2. In covalently bonded substances, where electrons are shared between two atoms, oxi- dation numbers are assigned by an arbitrary system based on relative electronegativities. For symmetrical covalent molecules such as H 2 and Cl 2 , each atom is assigned an oxida- tion number of zero because the bonding pair of electrons is shared equally between two like atoms, neither of which is more electronegative than the other: Cl Cl H H When the covalent bond is between two unlike atoms, the bonding electrons are shared unequally because the more electronegative element has a greater attraction for them. In this case the Oxidation Numbers are determined by assigning both electrons to the more electronegative element. Thus in compounds with covalent bonds such as NH 3 and H 2 O, O H H N H H H shared pairs of electrons the pairs of electrons are unequally shared between the atoms and are attracted toward the more electronegative elements, N and O. This causes the N and O atoms to be relatively negative with respect to the H atoms.
- Morris Hein, Scott Pattison, Susan Arena, Leo R. Best(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
Oxidation Numbers can be zero, positive, or negative. An Oxidation Number of zero means the atom has the same number of electrons assigned to it as there are in the free neutral atom. A positive Oxidation Number means the atom has fewer electrons assigned to it than in the neutral atom, and a negative Oxidation Number means the atom has more electrons assigned to it than in the neutral atom. The Oxidation Number of an atom that has lost or gained electrons to form an ion is the same as the positive or negative charge of the ion. (See Table 17.1.) In the ionic compound NaCl, the Oxidation Numbers are clearly + 1 for the Na + ion and - 1 for the Cl - ion. The Na + ion has one less electron than the neutral Na atom, and the Cl - ion has one more electron than the neutral Cl atom. In MgCl 2 , two electrons have transferred from the Mg atom to the two Cl atoms; the Oxidation Number of Mg is + 2. In covalently bonded substances, where electrons are shared between two atoms, oxida- tion numbers are assigned by an arbitrary system based on relative electronegativities. For symmetrical covalent molecules such as H 2 and Cl 2 , each atom is assigned an oxidation num- ber of zero because the bonding pair of electrons is shared equally between two like atoms, neither of which is more electronegative than the other: Cl Cl H H When the covalent bond is between two unlike atoms, the bonding electrons are shared un- equally because the more electronegative element has a greater attraction for them. In this case the Oxidation Numbers are determined by assigning both electrons to the more electronegative element. Thus in compounds with covalent bonds such as NH 3 and H 2 O, O H H N H H H shared pairs of electrons the pairs of electrons are unequally shared between the atoms and are attracted toward the more electronegative elements, N and O. This causes the N and O atoms to be relatively negative with respect to the H atoms.
eBook - PDF- David A. Ucko(Author)
- 2013(Publication Date)
- Academic Press(Publisher)
3 The sum of the Oxidation Numbers of all atoms in a molecule or formula unit must equal 0. For example, in NaCl, the sum of the Oxidation Numbers, ( + 1) + ( -1 ) , equals 0. This rule is based on the fact that all compounds are neutral. 4 In a polyatomic ion, the sum of the Oxidation Numbers of the atoms must equal the charge of the ion. For example, in the hydroxide ion, O H , oxygen has an Oxidation Number of -2 and hydrogen has an Oxidation Number of + 1 . The sum, ( -2 ) + ( + 1 ) = -1 , is equal to the charge of the ion. 5 The most electronegative atoms in a compound have negative oxi-dation numbers because they have gained electrons. Oxygen usually has an Oxidation Number of -2 . (In peroxides, however, oxygen's Oxidation Number is -1 . ) The halogens (Group 7A)—F, CI, Br, I— generally have Oxidation Numbers of -1 . Less electronegative ele-ments, especially metals, usually have positive Oxidation Numbers. Hydrogen usually has an Oxidation Number of + 1 . However, when hydrogen is bonded to an even less electronegative element such as a metal, forming a hydride, its Oxidation Number is - 1, because hydrogen gains an electron is these compounds. In simple ionic compounds, the assignment of Oxidation Numbers is easy if we are familiar with the charges of the ions involved. For example, in potassium oxide, K 2 0 , potassium has an Oxidation Number of + 1 and oxygen, -2 . These Oxidation Numbers are the same as the charges of the ions. To keep track of Oxidation Numbers, we write them above the formula, as shown here. 482 OXIDATION-REDUCTION REACTIONS AND ELECTROCHEMISTRY / 15.1 Oxidation NumberS Oxidation Number of potassium + 1 - 2 Oxidation Number of oxygen K 2 0 We then check that the sum of the Oxidation Numbers is zero: 2(+ 1) + ( -2 ) = 2 - 2 = 0. Note that the Oxidation Numbers written above the formula are for only one atom (or ion) of each element.
eBook - PDFFoundations of Chemistry
An Introductory Course for Science Students
- Philippa B. Cranwell, Elizabeth M. Page(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
d. The zinc atoms lose two electrons and are oxidised to Zn 2+ . The Cu 2+ ions in CuSO 4 gain two electrons and are reduced to Cu(0) in the metal. 5.1.2 Oxidation Number The Oxidation Number of an element (ON) provides a useful tool to help keep track of electrons in chemical reactions and indicate which species is being oxi-dised and which reduced. The Oxidation Number of an element in a compound tells us about the degree of oxidation of the element. Sometimes the term oxida-tion state is used in place of Oxidation Number, but both have the same meaning. Atoms in elemental species such as magnesium metal (Mg) or oxygen mole-cules (O 2 ) have been neither oxidised nor reduced so their oxidation state is said to be zero. In an ionic compound, the Oxidation Number is effectively the charge on the ion. In the compound magnesium oxide, MgO, the Oxidation Number of the magnesium ion is +2, as it is present as the Mg 2+ ion, and so is in a higher oxi-dation state than in the metal. O 2– ion O atom O O 2– Figure 5.2 Electron arrangements in an O atom and O 2 -ion. 5.1 Redox reactions 133 The oxide ion has a charge of -2, so the Oxidation Number of the oxygen is therefore -2. The compound could be written as Mg 2+ O 2 -. There are a few important principles to note here: 1. The Oxidation Number of an element in an ionic compound is the same as the charge on the ion. 2. An Oxidation Number can be positive, e.g. +2, or negative, e.g. – 2. If the Oxidation Number is positive, the atom is more oxidised than in the ele-ment. If the Oxidation Number is negative, the atom is more reduced than in the element. 3. The Oxidation Number of an atom in the elemental state is zero. For example, the Oxidation Number of the magnesium atom in the metal is zero. The Oxidation Number of the oxygen atoms in oxygen gas is zero. 4. Because a compound is uncharged overall, the sum of the oxidation num-bers in a neutral compound is zero.
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
216 Chapter 5 | Oxidation–Reduction Reactions Rules for the Assignment of Oxidation Numbers We can use some basic knowledge learned earlier plus the set of rules below to determine the Oxidation Numbers of the atoms in almost any compound. These rules are a hierarchy where a rule can supersede a rule below it if necessary. This allows us to deal with conflicts between two rules. Rules for Assigning Oxidation Numbers Note: When there is a conflict between two of these rules or an ambiguity in assigning an Oxidation Number, apply the rule that comes first and ignore the conflicting rule. 1. Oxidation Numbers must add up to the charge on the molecule, formula unit, or ion. 2. All the atoms of free elements have Oxidation Numbers of zero. 3. Metals in Groups 1A, 2A, and Al have +1, +2, and +3 Oxidation Numbers, respectively. 4. H and F, in compounds, have +1 and -1 Oxidation Numbers, respectively. 5. Oxygen has a -2 Oxidation Number. 6. Group 7A elements have a -1 Oxidation Number. 7. Group 6A elements have a -2 Oxidation Number. 8. Group 5A elements have a -3 Oxidation Number. These rules can be applied to molecular and ionic compounds as well as to elements and ions. For example, oxygen, O 2 , is a free element and has a zero charge; the Oxidation Number for each oxygen atom in O 2 must be zero, following Rules 1 and 2. Or take the simple ionic compound, NaCl: the sodium is in Group 1A and has an Oxidation Number of +1, from Rule 3; and from Rule 6, chlorine has an Oxidation Number of -1. These add up to zero, which is the overall charge of the formula unit which, necessarily, also obeys Rule 1. The numbered rules and their order in the list given above come into play when an element is capable of having more than one oxidation state. For example, you learned that transition metals can form more than one ion. Iron, for example, forms Fe 2+ and Fe 3+ ions, so in an iron compound we have to use the rules to figure out which iron ion is present.
eBook - PDF- Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
PRACTICE EXERCISE 12.1 Identify the substances oxidised and reduced and the oxidising and reducing agents in the reaction of iron and chlorine to form FeCl 3 . 570 Chemistry Oxidation Numbers The reaction of magnesium with oxygen in worked example 12.1 is clearly a redox reaction. However, not all reactions with oxygen produce ionic products. For example, the reaction of sulfur with oxygen is a redox reaction, but the product sulfur dioxide, SO 2 , is covalent. To help us follow electron transfers in such reactions, we use the convenient concept of the Oxidation Number. The Oxidation Number (or oxidation state) is the hypothetical charge that an individual atom or ion in a molecule would possess if the shared electron pairs in each covalent bond were assigned to the more electronegative element in the bond. Thus, the Oxidation Number is the charge that each atom would have if the compound were divided into monatomic ions. Oxidation Numbers are always given for an individual atom or ion and not for groups of atoms or ions. The Oxidation Number of an element in a particular compound is assigned according to the following basic rules. 1. The Oxidation Number of any free element (an element not combined chemically with a different element) is 0. For example, Ar, Fe, O in O 2 , P in P 4 and S in S 8 all have Oxidation Numbers of 0. 2. The Oxidation Number for any simple, monatomic ion (e.g. Na + and Cl - ) is equal to the charge on the ion. 3. The sum of the Oxidation Numbers of all atoms in a neutral molecule must equal zero. The sum of all Oxidation Numbers in a polyatomic ion must equal the charge on the ion. 4. In all of its compounds, fluorine has an Oxidation Number of -1. 5. In most of its compounds, hydrogen has an Oxidation Number of +1. 6. In most of its compounds, oxygen has an Oxidation Number of -2. In addition to these basic rules, some other chemical knowledge is required.
eBook - PDF- Morris Hein, Susan Arena, Cary Willard(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
Therefore, the Oxidation Number of the N atom is −3, and the Oxidation Number of each H atom is +1 (see illustration). Assigning correct Oxidation Numbers to elements is essential for balancing oxidation– reduction equations. The oxidation number or oxidation state of an element is an integer value assigned to each element in a compound or ion that allows us to keep track of electrons associated with each atom. Oxidation Numbers have a variety of uses in chemistry—from writing formulas to predicting properties of compounds and assisting in the balancing of oxidation–reduction reactions in which electrons are transferred (see photo). As a starting point, the Oxidation Number of an uncombined element, regardless of whether it is monatomic or diatomic, is zero (see photo). Rules for assigning oxidation num- bers are summarized below. 17.1 Oxidation Number 421 Many elements have multiple Oxidation Numbers; for example, nitrogen: N 2 N 2 O NO N 2 O 3 NO 2 N 2 O 5 NO 3 − Oxidation Number 0 +1 +2 +3 +4 +5 +5 CHECK YOUR UNDERSTANDING 17.1 Oxidation Numbers 17.2 Oxidation Numbers of Elements 3. O is −2, except in peroxides, where it is −1 and in OF 2 , where it is +2. 4. The metallic element in an ionic compound has a positive Oxidation Number. 5. In covalent compounds the negative Oxidation Number is assigned to the most electronegative atom. 6. The algebraic sum of the Oxidation Numbers of the elements in a compound is zero. 7. The algebraic sum of the Oxidation Numbers of the elements in a polyatomic ion is equal to the charge of the ion. Problem-Solving Strategy Finding the Oxidation Number of an Element in a Compound 1. Write the Oxidation Number of each known atom below the atom in the formula. 2. Multiply each Oxidation Number by the number of atoms of that element in the compound. 3. Write an expression indicating the sum of all the Oxidation Numbers in the compound. Note: The sum of the Oxidation Numbers in a compound must equal zero.
- Peter V. Hobbs(Author)
- 2000(Publication Date)
- Cambridge University Press(Publisher)
[Charge on sulfur atom = molecular charge - charge on one hydrogen atom - charge on three oxygen atoms = -1 -1 - 3 (-2) = +4.] This fictitious charge is called the Oxidation Number of sulfur in HSOi(aq). Following the same procedure, the Oxidation Number of sulfur in HSCX^aq) is: -1 -1 - 4(-2) = +6. Thus, according to this method of book-keeping, the two electrons on the right side of Reaction (6.5) originate from the sulfur atom, which changes its Oxidation Number from +4 to +6 or, stated another way, from changing its oxidation state from S(IV) to S(VI) - read sulfur four to sulfur six. 2 106 Oxidation-reduction reactions The (arbitrary) rules used in assigning Oxidation Numbers are as follows: 1. The Oxidation Number of a monoatomic substance is the charge on the atom [e.g., Cu + (aq) and S 2 ~(aq) have Oxidation Numbers of+1 and -2, respectively]. 2. In ionic binary compounds, the Oxidation Numbers are the charges per ion. [For example, CdCl 2 is an ionic compound, as indicated more clearly by Cd 2+ (C1~) 2 . Thus, the Oxidation Number of the cadmium ion is +2, and the Oxidation Number of each of the two chloride ions is -1.] The algebraic sum of the Oxidation Numbers of the atoms in an ion is equal to the charge on the ion (e.g., zero charge for CdCl 2 ). 3. In nonionic (covalent) compounds, the electrons involved in bond formation are shared, more or less equally, by the bonding atoms. However, to assign Oxidation Numbers, it is assumed that each bonding electron is attached to a particular atom. If these atoms are identical, the bonding electrons are shared equally between the two atoms. If the atoms are different, all of the electrons in the bond are assigned to the atom that has the greater attraction for electrons (as indicated by its electronegativity). The most electronegative elements, in order of decreasing electronegativity, are F, O, N, and Cl.
- Morris Hein, Susan Arena, Cary Willard(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
282 CHAPTER 17 OXIDATION-REDUCTION SOLUTIONS TO REVIEW QUESTIONS 1. (a) Iodine is oxidized. Its Oxidation Number increases from 0 to +5. (b) Chlorine is reduced. Its Oxidation Number decreases from 0 to −1. 2. The Oxidation Number for an atom in an ionic compound is the same as the charge of the ion that resulted when that atom lost or gained electrons to form an ionic bond. In a covalently bonded compound electrons are shared between the two atoms making up the bond. Those shared electrons are assigned to the atom in the bond with a higher electronegativity giving it a negative Oxidation Number. 3. Oxidation and reduction are complementary processes because one does not occur without the other. The loss of e − in oxidation is accompanied by a gain of e − in reduction. 4. Redox reactions are usually very difficult to balance by inspection so other more efficient methods are utilized to balance them. 5. + + → 2 6 Mn Mn 4 electrons are needed on the right-hand side of the equation. 6. − → 0 2 O O 2 electrons are needed on the left-hand side of the equation. 7. Break the reaction into half reactions; one for the oxidation reaction and one for the reduction reaction. 8. − + − + + → + 3 2 2 2 IO 12 H 10 e I 6 H O To balance this equation in base you would add 12 OH − ions to both sides of the reaction. On the left side, these OH − s would combine with the H + to form water. Cancel 6 H 2 O from each side of the equation. − − − + + → + 3 2 2 2 IO 6 H O 10 e I 12 OH 9. Oxidation of a metal occurs when the metal loses electrons. The easier it is for a metal to lose electrons, the more active the metal is. 10. The higher metal on the activity series list is more reactive. (a) Ca (b) Fe (c) Zn 11. If the free element is higher on the activity series than the ion with which it is paired, the reaction occurs. The higher metal on the activity series is more reactive.
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
When a substance is oxidized, it loses electrons, so the electron will move from A to B. Substance oxidized Substance reduced e – A B When the electron moves, the Oxidation Numbers will change. A has lost an electron, so its Oxidation Number will increase, and B has gained an electron, so its Oxidation Number will decrease. A A + + e – B + e – B – Substance oxidized Oxidation Number increased 0 +1 Substance reduced Oxidation Number decreased 0 –1 e – A B Then we need to label the oxidizing agent and the reducing agent. The oxidizing agent oxidizes the other substance, so it is reduced. B is reduced, it is the oxidizing agent. Therefore, A is the reducing agent. A A + + e – B + e – B – Substance oxidized Oxidation Number increased Reducing agent 0 +1 Substance reduced Oxidation Number decreased Oxidizing agent 0 –1 e – A B Summary 249 Summary Organized by Learning Objective Define oxidation, reduction, oxidizing agents, reducing agents, and Oxidation Numbers. Oxidation is the loss of electrons or an algebraic increase in Oxidation Number; reduction is the gain of electrons or an algebraic decrease in Oxidation Number. Both always occur together in redox reactions. The substance oxidized is the reducing agent; the substance reduced is the oxidizing agent. Oxidation Numbers are a bookkeeping device that we use to follow changes in redox reactions. They are assigned according to the rules in Section 5.1. The term oxidation state is equivalent to Oxidation Number. These terms are summarized in Table 5.1. Balance equations for oxidation–reduction reactions in acidic or basic solutions. For equations that cannot be balanced by inspection, the ion– electron method provides a systematic method for balancing a redox reaction in aqueous solution. According to this method, the skeleton Finally, the question asks for some examples of redox reactions. There are some given in the chapter, and we can make some up using the activity series.
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