Chemistry

Key unit competency:  Explain the concept of reduction and oxidation and balance equations for redox reactions

Learning Objectives 

By the end of this unit, a learner should be able to

 • Explain the redox reactions in terms of electron transfer and changes in oxidation state (number).

 • Explain the concept of disproportionation

 • Differentiate the reducing agent from the oxidizing agent in a redox reaction.

 • Work out the oxidation numbers of elements in the compounds.

 • Perform simple displacement reactions to order elements in terms of oxidizing or reducing ability.

 • Apply half-reaction method to balance redox reactions. 

 • Deduce balanced equations for redox reactions from relevant half equations.

Introductory activity 17 

The following figures illustrate the natural electrochemical process that is occurring on the surface of the metals. The figure A shows what we observe before and after the said process and the figure B shows what really happened chemically. Observe each of them carefully and answer the questions. 

         

                                                                    Figure A

      

                                                                      Figure B

    1. What does figure A represent?

    2.Based on some features  shown on the  figures above, suggest the name of the phenomenon  that is happening to the figures

    3. Write the overall chemical equation of the reaction that took place during the process

   4. What are the conditions for the process shown to occur?

   5. Is this process reversible? Explain.

   6. Discuss the economic impact of this process and suggest how it should be prevented.

17.1. Definition of electrochemistry and its relationship with redox reactions. 

Activity 17.1. 

1. Use examples to differentiate redox reactions from other chemical reactions 

2. Explain this statement: “Electrochemistry is a chapter of chemistry that studies the chemical reactions that produce electricity”

Electrochemistry is defined as the study of the interchange of chemical and Electrical energy. It is primarily concerned with two processes that involve oxidation–reduction reactions: the generation of an electric current from a spontaneous chemical reaction and, the opposite process, the use of a current to produce chemical change.

Electrochemistry is important in other less obvious ways. For example, the corrosion of iron, which has tremendous economic implications, is an electrochemical process. In addition, many important industrial materials such as aluminum, chlorine, and sodium hydroxide are prepared by electrolytic processes. 

       

17.2. Definitions of reduction and oxidation reactions. 

Activity 17.2 

Compare the following definitions of oxidation and reduction and answer the questions related 

 

a) Suggest at least five chemical equations where oxidation and reduction reactions occurred 

b) For chemical equations in (a), highlight  the chemical species that have lose electrons and those that have gained electrons 

c) Hence what can you conclude about oxidation and reduction reactions

17.2.1. Introduction 

The name “oxidation” was initially applied to reactions where substances combined with the element oxygen. Thus any substance burning in air was said to be oxidized, the product being some type of oxide. For example, burning carbon to produce carbon dioxide is an oxidation, as shown by the equation

                                      

Subsequently it was realized that reactions of substances with elements other than oxygen were of essentially the same type. For example, hydrogen can react with oxygen to form the compound water, but equally it can react with chlorine to form the compound hydrogen chloride. In both reactions the free element hydrogen is converted into a compound of hydrogen and another non-metal, and in both reactions, hydrogen atom tends to lose an electrons whereas the other atoms tend to gain electron; so both were classified as oxidations even though no oxygen was involved in the second case. 

                                       

The reverse reaction, conversion of compounds such as oxides of metals to the elemental metal were called “reduction” reactions, for example, the reduction of copper (II) oxide to copper by heating with charcoal (carbon). The reduction reaction was then defined as the loss of oxygen by a compound.

                                           

The gain or loss of oxygen is still a useful way of recognizing some oxidation or reduction reactions, but with knowledge of the structure of atoms, a rather wider definition of oxidation and reduction reaction has been adopted. 

17.2.2. Oxidation reactions as a loss of electrons and reduction as a gain of electrons. 

   

Hence a redox reaction is a combination of two half-reactions: an oxidation halfreaction and a reduction half-reaction. Nevertheless, one half-reaction cannot exist without the other, because electrons lost in the oxidation process must be captured in the reduction process, this explains why we talk of oxidation-reduction or redox reaction. 

The characteristic of a redox reaction is that there is exchange or transfer of electrons between chemical species participating in the reaction. 

We can compare this to the emigration-immigration movement: when a person leaves a country, emigration for that country, he/she must enter another country, immigration for that country and this constitutes an emigration-immigration movement. 

We notice that any chemical species tha sees its oxidation state increase is oxidized: 

    

17.3. Explanation of oxidizing and reducing agents 

Activity 17.3 

1. Based on the positions in the periodic table, which of the following reactions would you expect to occur?

                        

                                               

1. Aqueous copper (II) ion reacts with aqueous iodide ion to yield solid copper (I) iodide and aqueous iodine.

 a. Write the net ionic equation, 

 b. Assign oxidation numbers to all species present, and

 c. Identify the oxidizing and reducing agents.

Substances that cause changes in the oxidation state are called Oxidizing agents or Reducing agents. To increase oxidation state of a chemical species, the oxidizing agent must take one or more electrons from the element.

As the chemical species being oxidized loses electron(s), its oxidation number increases. However, the electrons don’t disappear. The oxidizing agent takes those electrons, and therefore the oxidation number of the oxidizing agent decreases. 

Let us consider the reaction between hydrogen chloride solution and zinc metal: 

                             

In the above reaction, hydrogen in hydrogen chloride takes an electron from the zinc metal. The oxidation state of zinc metal increases from 0 to +2; it is oxidized. By taking the electrons from zinc metal, the oxidation state of hydrogen decreases from  +1 to 0; it is reduced. 

Hydrochloric acid  is the oxidizing agent that causes oxidation of zinc to occur. Similarly, the zinc metal donates electrons to the hydrogen ion in hydrochloric acid, causing the oxidation state of hydrogen to decrease from +1 to 0. By providing the  electrons necessary to reduce hydrogen ion, the oxidation number of zinc increases from 0 to +2. 

Oxidizing agents are substances or chemical species containing the element or species that accepts electrons allowing another element or species to be oxidized By accepting electrons, the element or species in the oxidizing agent are reduced. 

Reducing agents are substances or chemical species containing the element or species that donate electrons, allowing another element or species to be reduced. By giving up electrons, the element or species in the reducing agent are oxidised. 

Table 17.1: Comparison between a reducing agent and an oxidizing agent 

In general, metals give up electrons and act as reducing agents, while reactive nonmetals such as O₂ and the halogens accept electrons and act as oxidizing agents.

      

       

17.4. Rules used to determine oxidation numbers of elements in chemical compounds and species 

              

The oxidation number of an atom is the apparent or real charge that the atom has when all bonds between atoms of different elements are assumed to be ionic. By comparing the oxidation number of an element or chemical species before and after reaction, we can tell whether the atom has gained or lost electrons. Note that oxidation numbers don’t necessarily imply ionic charges; they are just a convenient device to help keep track of electrons during redox reactions. 

The rules for assigning oxidation numbers are as follows: 

 1. The oxidation number of an element in its elemental form is 0 (zero).

 2. In compounds, the oxidation number of oxygen is almost always –2. The most common exception is in peroxides, when the oxidation number is –1. Peroxides are compounds having two oxygen atoms bonded together. For example, hydrogen 

     

          

17.5. Balancing of redox equations 

Activity 17.5 

To study the redox reactions between iron(II) sulphate and hydrogen peroxide 

Apparatus/materials 

Iron powder, 2M sulphuric acid, Bunsen burner, hydrogen peroxide (20 cm3), filter paper, test-tube and test-tube racks, 2M sodium hydroxide, filter funnel, 2M ammonia solution. 

Procedure

 1. Put about 0.5g of iron powder in a clean test tube and add 2cm³ of dilute sulphuric acid(2M).Warm the mixture gently and test the gas  produced with a burning splint. Name the gas formed in the reaction. Allow the solution to cool then filter and divide the filtrate into three portions

 2. To the first portion add a few drops of 2M sodium hydroxide and shake the contents of the test-tube then allow it to settle. Identify the product formed then add dilute sulphuric acid to the product. Record your observation. Empty the contents immediately and then rinse the test-tube with plenty of water brushing it thoroughly as you do so.

 3. To the second portion add a few drops of hydrogen peroxide followed by one or two drops of dilute surphuric acid and warm gently. Allow the solution to cool (or cool it under running tap water). To the cold solution add drop wise 2M NaOH until there is no further change. Record your observations. Add dilute sulphuric acid to the resultant product and note down your observations. Rinse the test tube thoroughly.

 4. To the third portion, add about 1 cm³ of dilute hydrogen peroxide solution followed by one or two drops of dilute sulphuric acid. Warm gently and test the gas produced with a glowing splint. Allow the solution to cool (or cool it using running tap water).To the cold solution add ammonia solution drop wise until no further change. Compare the product formed when ammonia solution to that obtained when sodium hydroxide was used.

Study Questions

 1. Name the products formed when dilute sulphuric acid reacts with iron powder. Write a balanced formula equation for the reaction

 2. When dilute sulphuric acid reacts with iron powder, iron atoms are oxidized and hydrogen ions are reduced. Write a balanced

 a) oxidation half-equation

 b)  reduction  half-equation and

 c) overall redox equations for the reaction between iron and sulphuric acid

 3. What is the effect of adding a hydrogen peroxide in step 4?

 4. What will be the effect of adding concentrated nitric acid to any iron salt? Explain why concentrated nitric acid does not react with pure iron metal

17.5.1. Rules for balancing redox reactions 

The Half-Reaction Method for Balancing Oxidation–Reduction Reactions in Aqueous Solutions

For oxidation–reduction reactions that occur in aqueous solution, it is useful to separate the reaction into two half-reactions: one involving oxidation reaction  and the other involving reduction reaction. Then after balancing those half reactions, find the overall oxidation-reduction (redox) reaction by summing up the two halfreactions. 

For example, consider the unbalanced equation for the oxidation– reduction reaction between cerium(IV) ion and tin(II) ion:

      

             

a) The half-reaction method for balancing equations for oxidation–reduction reactions occurring in acidic solution

 1. Write separate equations for the oxidation and reduction half-reactions.

 2. For each half-reaction, a. balance all the elements except hydrogen and oxygen. b. then balance oxygen using H2O. c. then balance hydrogen using H+. d. then balance the charge using electrons.

 3. If necessary, multiply one or both balanced half-reactions by an integer to equalize the number of electrons transferred in the two half-reactions. 

 4. Add the half-reactions, and cancel identical species. 

 5. Checking that the elements and charges are balanced.

17.5.2. Concept of disproportionation reactions 

Disproportionation is a chemical reaction, typically a redox reaction, where an element in a molecule or chemical species is simultaneously oxidized and reduced. An example is the disproportionation of copper in the following reaction: 

17.6. Reactivity series of metals 

Activity 17.6 

Carry out the following experiment regarding the displacement  reaction and the reaction of sodium, magnesium, zinc, and copper with cold/hot water and dilute hydrochloric acid

Equipments/materials

Metal samples (of consistent mass and surface area): Na, Mg, Zn, Cu 1.0 M solution of copper (II) sulphate, Cu(SO4)2, Distilled water Test tubes/ test tube rack 3 M HCl 400 mL Beaker 

Procedure 

1. Place a small sample of each metal in test tubes containing 5 mL of cold water.  For metals like magnesium, it may be necessary to clean the outside of the metal with steel wool.  Watch for evidence of reaction and note any changes you observe

2. Place the test tubes in a 400 mL beaker that is about 1/3 full of boiling water.  After a few minutes, look for evidence of reaction.  Note any changes.  Did some metals that didn’t react with cold water, react with hot water? 

3. Place a small sample of each metal in test tubes containing 5 mL of 1.0 mol/L hydrochloric acid, HCl.  Watch for evidence of reaction.  Note any changes 

4. Place a small sample of magnesium ribbon in test tube containing 5 mL of 1M copper (II) sulphate.  Watch for evidence of reaction and note any changes 

Study questions

 1)  Considering sodium, magnesium, zinc, and copper: Arrange the metals in order of increasing reactivity (from least reactive to most reactive)

 2) Which of the four metals are reacting with cold water?  For those metals that did react, write a balanced symbolic equation.

 3) Which of the four metals are reacting with hot water?  For those metals that did react, write a balanced symbolic equation. 

 4) Which of the four metals are reacting with the hydrochloric acid?  For those metals that did react, write a balanced symbolic equation. 

 5) Which metal did not react with either water or hydrochloric acid?

 6) Which of the four metals would be suitable for making saucepans?  Explain why the others are not. 

7. Describe what you would see if you dropped a piece of magnesium ribbon into some copper (II) sulphate solution in a test tube.  Write a chemical equation for the reaction.

The reactivity series is a series of metals, in order of reactivity, as reducing agents, from highest to lowest reducing agent. It is used to determine the products of single displacement reactions, whereby metal A will displace another metal B in a solution if A is higher in the series. Although hydrogen is not a metal, it is included in the reactivity series for comparison (Table 17.2). 

When a metal is placed in a solution of another metal salt, and if the metal is more active than the metal in the salt, the more active metal displaces the other metal from its salt:

         

In this example, since Zn is above Cu in the reactivity series, it displaces copper from its salts.

Examples of displacement reactions

Only a metal higher in the reactivity series will displace another. A metal can displace metal ions listed below it in the activity series, but not above. For example, zinc is more active than copper and is able to displace copper ions from solution.

Conclusion: a more reactive metal will displace a less reactive metal from its salt solution (irrespective of which salt)

17.7. End unit assessment