Balancing equations oxidation number change

Redox reactions may be represented by balanced equations. Oxidation numbers are assigned to those substances involved in the oxidation/reduction. The total change in oxidation number is set to zero by adjusting the numbers of reactants and products. In aqueous solution, H2O, H+Caq) or OH (aq) ions are included as necessary to balance the number of hydrogen and oxygen atoms. 

Balance the following equation by the oxidation number change method ... 

Section 16.1 introduces the concept of oxidation number and how to calculate the oxidation number of an element from the formula of the compound or ion of which it is a part. Section 16.2 describes how to use the oxidation numbers to name compounds, formalizing and extending the rules given in Chapter 6. Section 16.3 shows how to predict possible oxidation numbers from the position of the element in the periodic table and how to use these oxidation numbers to write probable formulas for covalent compounds. Section 16.4 presents a systematic method for balancing equations in which oxidation numbers change. 

Balance the equation for the reduction of HNO3 to NH4NO3 by Mn by the oxidation number change method. Add other compounds as needed. 

Equations for redox reactions are sometimes difficult to balance. Use the steps in Skills Toolkit 2 below to balance redox equations for reactions in acidic aqueous solution. An important step is to identify the key ions or molecules that contain atoms whose oxidation numbers change. These atoms are the starting points of the unbalanced half-reactions. For the reaction of zinc and hydrochloric acid, the unbalanced oxidation and reduction half-reactions would be as follows ... 

To see how oxidation numbers change, start hy assigning numbers, shown in Table 20-2, to all elements in the balanced equation. Then review the changes as shown in the accompanying diagram. 

One important use of oxidation numbers is in balancing redox equations. There are essentially two methods to balance redox reactions the oxidation number change method and the ion-electron method. In the former method, the changes in oxidation number are used to balance the species in which the elements that are oxidized and reduced appear. The numbers of atoms of each of these elements is used to give equal numbers of electrons gained and lost. If necessary, first balance the number of atoms of the element oxidized and/or the number of atoms of the element reduced. Then, balance by inspection, as was done in Chapter 7. 

Since this is a reduction, the other reaction must be an oxidation. The sulfite ion will be oxidized to the sulfate ion with the sulfur atom losing two electrons (sulfur s oxidation number changes from +4 to +6). Two protons and two electrons balance the equation... 

In a balanced equation, the total change in oxidation number is zero. 

Cupric sulfide, copper(II) sulfide, reacts with hot nitric acid to produce nitric oxide gas, NO, and elemental sulfur. Only the oxidation numbers of S and N change. Write the balanced equation for the reaction. 

Note that all the added atoms in steps 4 and 5 have the same oxidation number as the atoms already in the equation. The atoms changing oxidation number have already been balanced in steps 1 and 2. 

The Law of Conservation of Mass states that the total mass remains unchanged. This means that the total mass of the atoms of each element represented in the reactants must appear as products. In order to indicate this, we must balance the reaction. When balancing chemical equations, it is important to realize that you cannot change the formulas of the reactants and products the only things you may change are the coefficients in front of the reactants and products. The coefficients indicate how many of each chemical species react or form. A balanced equation has the same number of each type of atom present on both sides of the equation and the coefficients are present in the lowest whole number ratio. For example, iron metal reacts with oxygen gas to form rust, iron(III) oxide. We may represent this reaction by the following balanced equation ... 

The usefulness of determining the oxidation number in analytical chemistry is twofold. First, it will help determine if there was a change in oxidation number of a given element in a reaction. This always signals the occurrence of an oxidation-reduction reaction. Thus, it helps tell us whether a reaction is a redox reaction or some other reaction. Second, it will lead to the determination of the number of electrons involved, which will aid in balancing the equation. These latter points will be discussed in later sections. 

Step 1 Look at the equation to be balanced and determine what is oxidized and what is reduced. This involves checking the oxidation numbers and discovering which have changed. 

In section 10.2, you learned that a redox reaction involves changes in oxidation numbers. If an element undergoes oxidation, its oxidation number increases. If an element undergoes reduction, its oxidation number decreases. When balancing equations by the half-reaction method in section 10.3, you sometimes used oxidation numbers to determine the reactant(s) and product(s) in each half-reaction. 

The key to the oxidation-number method of balancing redox equations is to realize that the net change in the total of all oxidation numbers must be zero. That is, any increase in oxidation number for the oxidized atoms must be matched by a corresponding decrease in oxidation number for the reduced atoms. Take the reaction of potassium permanganate (KMn04) with sodium bromide in aqueous acid, for example. An aqueous acidic solution of the purple permanganate anion (Mn04 ) is reduced by Br- to yield the nearly colorless Mn2+ ion, while Br- is oxidized to Br2. The unbalanced net ionic equation for the process is... 

Previously balanced equations are pretty straightforward each of the elements involved in the oxidations and reductions only occurred once on a side. This equation is different because chlorine appears in more than one place on the right. Not all of the chlorine is taking part in a change in oxidation number. Notice that Cl- is in HC1 on the left, and Cl- is also in CrCl3 on the right. Additionally, some Cl- from HC1 is being oxidized to Cl2 (neutral). This dual role of chlorine causes no problems. 

It is not necessary to balance the equation. All we need to know is that the oxidation number on the iron changes from +2 to +3, as indicated by the charges on the iron ions. The equivalent mass FeSC>4 is... 

The first step in any method of balancing oxidation-reduction equations is to identify the element that is oxidized and the one that is reduced. Because the change in oxidation number is equal to a change in the number of electrons controlled, and the electrons must be controlled by some atom, the total gain in oxidation number is equal to the total loss in oxidation number. The oxidation half of a reaction may be written in one equation, and the reduction half in another. Neither half-reaction can be carried out without the other, but they can be done in different locations if they are connected in such a way that a complete electrical circuit is made (Chapter 17). The half-reaction method is illustrated by balancing the equation for the reaction of zinc metal with dilute nitric acid to produce ammonium ion, zinc ion, and water ... 

Use changes in oxidation number to balance redox equations. 

For each of the following balanced equations, write the oxidation number above the symbol of each atom that changes oxidation state in the course of the reactions. 

Most redox equations can be balanced and (2) the change-in-oxidation-number method. Many redox equations can be balanced... 

Frequently we need more oxygen or hydrogen to complete the mass balance for a reaction or half-reaction in aqueous solution. We must be careful, however, not to introduce other changes in oxidation number or to use species that could not actually be present in the solution. We cannot add H2 or O2 to equations because these species are not present in aqueous solutions. Acidic solutions do not contain significant concentrations of OH ions. Basic solutions do not contain significant concentrations of H+ ions. 

Balance the atom which changes oxidation number in each partial ionic equation ... 

---