What Are the Rules for Assigning Oxidation Numbers?
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Electrochemical reactions involve the transfer of electrons . Mass and charge are conserved when balancing these reactions, but you need to know which atoms are oxidized and which atoms are reduced during the reaction. Oxidation numbers are used to keep track of how many electrons are lost or gained by each atom. These oxidation numbers are assigned using the following rules.
Rules for Assigning Oxidation Numbers
- The convention is that the cation is written first in a formula, followed by the anion . For example, in NaH, the H is H-; in HCl, the H is H+.
- The oxidation number of a free element is always 0. The atoms in He and N 2 , for example, have oxidation numbers of 0.
- The oxidation number of a monatomic ion equals the charge of the ion. For example, the oxidation number of Na + is +1; the oxidation number of N 3- is -3.
- The usual oxidation number of hydrogen is +1. The oxidation number of hydrogen is -1 in compounds containing elements that are less electronegative than hydrogen, as in CaH 2 .
- The oxidation number of oxygen in compounds is usually -2. Exceptions include OF 2 because F is more electronegative than O, and BaO 2 , due to the structure of the peroxide ion, which is [O-O] 2- .
- The oxidation number of a Group IA element in a compound is +1.
- The oxidation number of a Group IIA element in a compound is +2.
- The oxidation number of a Group VIIA element in a compound is -1, except when that element is combined with one having a higher electronegativity. The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl.
- The sum of the oxidation numbers of all of the atoms in a neutral compound is 0.
- The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion. For example, the sum of the oxidation numbers for SO 4 2- is -2.
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22.6: Assigning Oxidation Numbers
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Moving from studying the element iron to iron compounds, we need to be able to clearly designate the form of the iron ion. An example of this is iron that has been oxidized to form iron oxide during the process of rusting. Although Antoine Lavoisier first began the idea of oxidation as a concept, it was Wendell Latimer (1893-1955) who gave us the modern concept of oxidation numbers. His 1938 book The Oxidation States of the Elements and Their Potentials in Aqueous Solution laid out the concept in detail. Latimer was a well-known chemist who later became a member of the National Academy of Sciences. Not bad for a gentleman who started college planning on being a lawyer.
Assigning Oxidation Numbers
The oxidation number is a positive or negative number that is assigned to an atom to indicate its degree of oxidation or reduction. In oxidation-reduction processes, the driving force for chemical change is in the exchange of electrons between chemical species. A series of rules have been developed to determine oxidation numbers:
- For free elements (uncombined state), each atom has an oxidation number of zero. \(\ce{H_2}\), \(\ce{Br_2}\), \(\ce{Na}\), \(\ce{Be}\), \(\ce{K}\), \(\ce{O_2}\), \(\ce{P_4}\), all have an oxidation number of 0.
- Monatomic ions have oxidation numbers equal to their charge. \(\ce{Li^+} = +1\), \(\ce{Ba^{2+}} = +2\), \(\ce{Fe^{3+}} = +3\), \(\ce{I^-} = -1\), \(\ce{O^{2-}} = -2\), etc. Alkali metal oxidation numbers \(= +1\). Alkaline earth oxidation numbers \(= +2\). Aluminum \(= +3\) in all of its compounds. Oxygen's oxidation number \(= -2\) except when in hydrogen peroxide \(\left( \ce{H_2O_2} \right)\), or a peroxide ion \(\left( \ce{O_2^{2-}} \right)\) where it is \(-1\).
- Hydrogen's oxidation number is \(+1\), except for when bonded to metals as the hydride ion forming binary compounds. In \(\ce{LiH}\), \(\ce{NaH}\), and \(\ce{CaH_2}\), the oxidation number is \(-1\).
- Fluorine has an oxidation number of \(-1\) in all of its compounds.
- Halogens (\(\ce{Cl}\), \(\ce{Br}\), \(\ce{I}\)) have negative oxidation numbers when they form halide compounds. When combined with oxygen, they have positive numbers. In the chlorate ion \(\left( \ce{ClO_3^-} \right)\), the oxidation number of \(\ce{Cl}\) is \(+5\), and the oxidation number of \(\ce{O}\) is \(-2\).
- In a neutral atom or molecule, the sum of the oxidation numbers must be 0. In a polyatomic ion, the sum of the oxidation numbers of all the atoms in the ion must be equal to the charge on the ion.
Example \(\PageIndex{1}\)
What is the oxidation number for manganese in the compound potassium permanganate \(\left( \ce{KMnO_4} \right)\)?
The oxidation number for \(\ce{K}\) is \(+1\) (rule 2).
The oxidation number for \(\ce{O}\) is \(-2\) (rule 2).
Since this is a compound (there is no charge indicated on the molecule), the net charge on the molecule is zero (rule 6).
So we have:
\[\begin{align*} +1 + \ce{Mn} + 4 \left( -2 \right) &= 0 \\ \ce{Mn} - 7 &= 0 \\ \ce{Mn} &= +7 \end{align*}\nonumber \]
When dealing with oxidation numbers, we must always include the charge on the atom.
Another way to determine the oxidation number of \(\ce{Mn}\) in this compound is to recall that the permanganate anion \(\left( \ce{MnO_4^-} \right)\) has a charge of \(-1\). In this case:
\[\begin{align*} \ce{Mn} + 4 \left( -2 \right) &= -1 \\ \ce{Mn} - 8 &= -1 \\ \ce{Mn} &= +7 \end{align*}\nonumber \]
Example \(\PageIndex{2}\)
What is the oxidation number for iron in \(\ce{Fe_2O_3}\)?
\[\begin{align*} &\ce{O} \: \text{is} \: -2 \: \left( \text{rule 2} \right) \\ &2 \ce{Fe} + 3 \left( -2 \right) = 0 \\ &2 \ce{Fe} = 6 \\ &\ce{Fe} = 3 \end{align*}\nonumber \]
If we have the compound \(\ce{FeO}\), then \(\ce{Fe} + \left( -2 \right) = 0\) and \(\ce{Fe} = 2\). Iron is one of those materials that can have more than one oxidation number.
The halogens (except for fluorine) can also have more than one number. In the compound \(\ce{NaCl}\), we know that \(\ce{Na}\) is \(+1\), so \(\ce{Cl}\) must be \(-1\). But what about \(\ce{Cl}\) in \(\ce{NaClO_3}\)?
\[\begin{align*} \ce{Na} &= 1 \\ \ce{O} &= -2 \\ 1 + \ce{Cl} + 3 \left( -2 \right) &= 0 \\ 1 + \ce{Cl} - 6 &= 0 \\ \ce{Cl} - 5 &= 0 \\ \ce{Cl} &= +5 \end{align*}\nonumber \]
Not quite what we expected, but \(\ce{Cl}\), \(\ce{Br}\), and \(\ce{I}\) will exhibit multiple oxidation numbers in compounds.
- The oxidation number is a positive or negative number that is assigned to an atom to indicate its degree of oxidation or reduction.
- In oxidation-reduction processes, the driving force for chemical change is in the exchange of electrons between chemical species.
- Six rules for determining oxidation numbers are listed.
- Examples of oxidation number determinations are provided.
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Rules for Oxidation Numbers Study Guide
The oxidation number is defined as the total number of electrons an atom gains or losses in a chemical reaction to form a bond. It follows a set of oxidation number rules.
Introduction
Have you seen that there is a layer of rust forming on the surface of any of the metal rods around you? Most often metal surfaces are painted, or a coating is applied to decrease rusting. Have you ever wondered why? It is because the metals get oxidized and develop rust when left untreated. Let’s learn about oxidation in this guide.
RULES FOR ASSIGNING OXIDATION NUMBERS
Here are some of the rules for assigning oxidation numbers. They are as follows:
The pattern is that the cation is always written first in a formula, which is then followed by the anion.
The oxidation number of any free element always remains 0.
The oxidation number in the case of a monatomic ion is always equal to the value of the charge of the ion. For instance, the oxidation number of Na 3- is 3-.
The expected oxidation number of hydrogen is plus 1. The oxidation number of hydrogen remains -1 in compounds, including elements that have less electronegativity than hydrogen, as in NaH.
The oxidation number of oxygen in the case of compounds is typically -2. Anomalies include OF2 because F turns out to be more electronegative than O, and BaO2, due to the construction of the peroxide ion, which can be expressed as (O-O) 2- .
The oxidation number of elements in Group IA in a compound is +1.
The oxidation number of elements present in Group IIA in a compound is +2.
The oxidation number of elements present in Group VIIA in a compound is -1, except for the case when that element is conjoined with one having a higher value of electronegativity. The oxidation number of Cl turns out to be -1 in HCl, but this number becomes +1 in HOCl.
The sum of the value of oxidation numbers of all the present atoms in any neutral compound is 0.
Also, the sum of the value oxidation numbers in any given polyatomic ion is always equal to the value of the charge of the ion. For instance, the sum of the value of oxidation numbers for SO 4 2- is -2.
1. What are the rules related to assigning oxidation numbers?
If you want to know how to assign oxidation numbers, you need to know some rules of oxidation numbers. Some of them are:
- The pattern is that the cation is expressed first in a formula, and then the anion is written.
- The oxidation number of any given free element always has to be 0.
- The oxidation number of a monatomic ion equals the value of the charge of the ion.
2. What is the easiest way to determine oxidation numbers?
The easiest way to determine the oxidation number is by dividing the total charge by two.
3. How do you find the oxidizing agent?
To find an oxidizing agent, you need to check the oxidation number of an atom that is before and after the chemical reaction. If the oxidation number is higher in the product, then it has lost electrons, and the substance went through oxidization. If the oxidation number is less, then it has gained electrons and was reduced.
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- Assigning Oxidation Numbers: https://www.ck12.org/c/chemistry/assigning-oxidation-numbers/lesson/Assigning-Oxidation-Numbers-CHEM/ . Accessed 7th March 2022.
- Rules for Assigning Oxidation Numbers: https://www.toppr.com/ask/question/write-the-rules-for-assigning-oxidation-number/ . Accessed 7th March 2022.
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Rules for Assigning Oxidation Numbers to Elements
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Oxidation numbers are assigned to elements using these rules:
Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). This is also true for elements found in nature as diatomic (two-atom) elements
and for sulfur, found as:
Rule 2: The oxidation number of a monatomic (one-atom) ion is the same as the charge on the ion, for example:
Rule 3: The sum of all oxidation numbers in a neutral compound is zero. The sum of all oxidation numbers in a polyatomic (many-atom) ion is equal to the charge on the ion. This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.
Rule 4: The oxidation number of an alkali metal (IA family) in a compound is +1; the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.
Rule 5: The oxidation number of oxygen in a compound is usually –2. If, however, the oxygen is in a class of compounds called peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1. If the oxygen is bonded to fluorine, the number is +1.
Rule 6: The oxidation state of hydrogen in a compound is usually +1. If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.
Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.
Notice that the zinc metal (the reactant) has an oxidation number of zero (rule 1), and the zinc cation (the product) has an oxidation number of +2 (rule 2). In general, you can say that a substance is oxidized when there’s an increase in its oxidation number.
Reduction works the same way. Consider this reaction:
The copper is going from an oxidation number of +2 to zero. A substance is reduced if there’s a decrease in its oxidation number.
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How to Assign Oxidation Numbers
The oxidation number is the positive or negative number of an atom that indicates the electrical charge the atom has if its compound consists of ions. In other words, the oxidation number gives the degree of oxidation (loss of electrons) or reduction (gain of electrons) of the atom in a compound. Because they track the number of electrons lost or gained, oxidation numbers are a sort of shorthand for balancing charge in chemical formulas.
This is a list of rules for assigning oxidation numbers, with examples showing the numbers for elements, compounds, and ions.
Rules for Assigning Oxidation Numbers
Various texts contain different numbers of rules and may change their order. Here is a list of oxidation number rules:
- Write the cation first in a chemical formula, followed by the anion. The cation is the more electropositive atom or ion, while the anion is the more electronegative atom or ion. Some atoms may be either the cation or anion, depending on the other elements in the compound. For example, in HCl, the H is H + , but in NaH, the H is H – .
- Write the oxidation number with the sign of the charge followed by its value. For example, write +1 and -3 rather than 1+ and 3-. The latter form is used to indicate oxidation state .
- The oxidation number of a free element or neutral molecule is 0. For example, the oxidation number of C, Ne, O 3 , N 2 , and Cl 2 is 0.
- The sum of all the oxidation numbers of the atoms in a neutral compound is 0. For example, in NaCl, the oxidation number of Na is +1, while the oxidation of Cl is -1. Added together, +1 + (-1) = 0.
- The oxidation number of a monatomic ion is the charge of the ion. For example, the oxidation number of Na + is +1, the oxidation number of Cl – is -1, and the oxidation number of N 3- is -3.
- The sum of the oxidation numbers of a polyatomic ion is the charge of the ion. For example, the sum of the oxidation numbers for SO 4 2- is -2.
- The oxidation number of a group 1 (alkali metal) element in a compound is +1.
- The oxidation number of a group 2 (alkaline earth) element in a compound is +2.
- The oxidation number of a group 7 (halogen) element in a compound is -1. The exception is when the halogen combines with an element with higher electronegativity (e.g., oxidation number of Cl is +1 in HOCl).
- The oxidation number of hydrogen in a compound is usually +1. The exception is when hydrogen bonds with metals forming the hydride anion (e.g., LiH, CaH 2 ), giving hydrogen an oxidation number of -1.
- The oxidation number of oxygen in a compound is usually -2. Exceptions include OF 2 and BaO 2 .
Examples of Assigning Oxidation Numbers
Example 1: Find the oxidation number of iron in Fe 2 O 3 .
The compound has no electrical charge, so the oxidation numbers of iron and oxygen balance each other out. From the rules, you know the oxidation number of oxygen is usually -2. So, find the iron charge that balances the oxygen charge. Remember, the total charge of each atom is its subscript multiplied by its oxidation number. O is -2 There are 3 O atoms in the compound so the total charge is 3 x -2 = -6 The net charge is zero (neutral), so: 2 Fe + 3(-2) = 0 2 Fe = 6 Fe = 3
Example 2: Find the oxidation number for Cl in NaClO3.
Usually, a halogen like Cl has an oxidation number of -1. But, if you assume Na (an alkali metal) has an oxidation number of +1 and O has an oxidation number of -2, the charges don’t balance out to give a neutral compound. It turns out all of the halogens, except for fluorine, have more than one oxidation number. Na = +1 O = -2 1 + Cl + 3(-2) = 0 1 + Cl -6 = 0 Cl -5 = 0 Cl = -5
- IUPAC (1997) “Oxidation Number”. Compendium of Chemical Terminology (the “Gold Book”) (2nd ed.). Blackwell Scientific Publications. doi: 10.1351/goldbook
- Karen, P.; McArdle, P.; Takats, J. (2016). “Comprehensive definition of oxidation state (IUPAC Recommendations 2016)”. Pure Appl. Chem . 88 (8): 831–839. doi: 10.1515/pac-2015-1204
- Whitten, K. W.; Galley, K. D.; Davis, R. E. (1992). General Chemistry (4th ed.). Saunders.
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OXIDATION NUMBERS CALCULATOR
To calculate oxidation numbers of elements in the chemical compound, enter it's formula and click 'Calculate' (for example: Ca2+ , HF2^- , Fe4[Fe(CN)6]3 , NH4NO3 , so42- , ch3cooh , cuso4*5h2o ).
The oxidation state of an atom is the charge of this atom after ionic approximation of its heteronuclear bonds. The oxidation number is synonymous with the oxidation state. Determining oxidation numbers from the Lewis structure (Figure 1a) is even easier than deducing it from the molecular formula (Figure 1b). The oxidation number of each atom can be calculated by subtracting the sum of lone pairs and electrons it gains from bonds from the number of valence electrons. Bonds between atoms of the same element (homonuclear bonds) are always divided equally.
When dealing with organic compounds and formulas with multiple atoms of the same element, it's easier to work with molecular formulas and average oxidation numbers (Figure 1d). Organic compounds can be written in such a way that anything that doesn't change before the first C-C bond is replaced with the abbreviation R (Figure 1c). Unlike radicals in organic molecules, R cannot be hydrogen. Since the electrons between two carbon atoms are evenly spread, the R group does not change the oxidation number of the carbon atom it's attached to. You can find examples of usage on the Divide the redox reaction into two half-reactions page.
Rules for assigning oxidation numbers
- The oxidation number of a free element is always 0.
- The oxidation number of a monatomic ion equals the charge of the ion.
- Fluorine in compounds is always assigned an oxidation number of -1.
- The alkali metals (group I) always have an oxidation number of +1.
- The alkaline earth metals (group II) are always assigned an oxidation number of +2.
- Oxygen almost always has an oxidation number of -2, except in peroxides (H 2 O 2 ) where it is -1 and in compounds with fluorine (OF 2 ) where it is +2.
- Hydrogen has an oxidation number of +1 when combined with non-metals, but it has an oxidation number of -1 when combined with metals.
- The algebraic sum of the oxidation numbers of elements in a compound is zero.
- The algebraic sum of the oxidation states in an ion is equal to the charge on the ion.
Assigning oxidation numbers to organic compounds
- cysteine: HO2CCH(NH2)CH2SH
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AP®︎/College Chemistry
Unit 4: lesson 5.
- Oxidation–reduction (redox) reactions
Worked example: Using oxidation numbers to identify oxidation and reduction
- Balancing redox equations
- Worked example: Balancing a simple redox equation
- Worked example: Balancing a redox equation in acidic solution
- Worked example: Balancing a redox equation in basic solution
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Video transcript
IMAGES
VIDEO
COMMENTS
This chemistry tutorial discusses how to assign oxidation numbers and includes examples of how to determine the oxidation numbers in a
The convention is that the cation is written first in a formula, followed by the anion. · The oxidation number of a free element is always 0.
In a neutral atom or molecule, the sum of the oxidation numbers must be 0. In a polyatomic ion, the sum of the oxidation numbers of all the
RULES FOR ASSIGNING OXIDATION NUMBERS · The pattern is that the cation is always written first in a formula, which is then followed by the anion. · The oxidation
Proceed to remove all atoms around carbon assigning oxidation states to H (+1), O. (-2) and halogens (-1) until oxidation state of the central carbon atom
The sum of all oxidation numbers in a polyatomic (many-atom) ion is equal to the charge on the ion. This rule often allows chemists to calculate
Rules for Assigning Oxidation Numbers · Write the cation first in a chemical formula, followed by the anion. · Write the oxidation number with the
Rules for assigning oxidation numbers · The oxidation number of a free element is always 0. · The oxidation number of a monatomic ion equals the charge of the ion
Assigning Oxidation Numbers Using Rules · The oxidation number of a free, neutral element is zero. · The oxidation number of a monoatomic ion is
By assigning oxidation numbers to the atoms of each element in a redox equation, we can determine which element is oxidized and which element is reduced