Chapter 20 Oxidation-Reduction Reactions

Chapter 20 Oxidation-Reduction Reactions

Some Common Reactions The combustion of gasoline in an automobile engine requires oxygen Burning of wood in a fireplace requires oxygen The reactions that break down food in your body and release energy use oxygen The oxide of hydrogen is water Charcoal oxidizes when it burn forming CO2 Bleaching stains in fabric is still oxidation even though it does not burn.

CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
Oxidation When methane burns in air, it oxidizes and forms oxides of carbon and hydrogen. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) When elemental iron turns to rust, it slowly oxidizes to compounds such as iron (III) oxide. 4Fe(s) + 3O2(g) Fe2O3(s)

2Fe2O3(s) + 3C(s) 4Fe(s) + 3CO2(g)
Reduction Originally reduction meant a loss of oxygen from a compound 2Fe2O3(s) + 3C(s) Fe(s) + 3CO2(g) iron oxide carbon iron carbon dioxide Reduction of iron ore to metallic iron involves the removal of oxygen from iron (III) oxide. Involves heating the ore with carbon.

Question What happens to magnesium and oxygen when they react to form magnesium oxide? 2Mg O MgO magnesium oxygen magnesium oxide Magnesium loses electrons to form Mg2+ Oxygen gains electrons to form O2-

Electron Shift in Redox Reactions
The modern concept of oxidation and reduction have been extended to include many reactions that do not even involve oxygen. Oxygen is the most electronegative element (besides fluorine) When oxygen bonds with an atom of a different element (except fluorine), electrons from that atom shift toward oxygen.

Redox Reactions Redox reactions are currently understood to involve any shift of electron between reactants. Oxidation – a process that involves a complete or partial loss of electrons or a gain of oxygen. Results in an increase in the oxidation number of an atom Reduction – a process that involves a complete or partial gain of electrons or the loss of oxygen. Results in a decrease in the oxidation number of an atom

Redox Reactions Oxidation and reduction always occur simultaneously. The substance gaining oxygen or losing electrons is oxidized The substance losing oxygen or gaining electrons is reduced.

Redox Reactions During a reaction between a metal and a nonmetal, electrons are transferred from atoms of the metal to atoms of the nonmetal. Mg S Mg2+S2- magnesium sulfur magnesium sulfide 2 electrons are transferred from a magnesium atom to a sulfur atom. Magnesium atoms become more stable by the loss of electrons. Sulfur atoms become more stable by the gain of electrons

Redox Reactions Mg + S Mg2+S2-
magnesium sulfur magnesium sulfide Oxidation: Mg Mg e- (loss of electrons) Reduction: S e S2- (gain of electrons) Magnesium atom is said to be oxidized to a magnesium ion Sulfur atom is said to be reduced to a sulfide ion.

Redox Reactions When a metal combines with oxygen, it loses electrons
When oxygen is removed from the oxide of a metal, the metal gains electrons. This knowledge is what led to the broader definition of oxidation and reduction as an exchange of electrons.

Redox Reactions Reducing agent – the substance that loses the electrons Mg S MgS reducing agent oxidized Oxidizing agent – the substance that accepts electrons Mg S MgS oxidizing agent reduced

Sample Problem 2AgNO3 + Cu Cu(NO3)2 + 2Ag
Silver nitrate reacts with copper to form copper nitrate and silver. From the equation below, determine what is oxidized and what is reduced. Identify the oxidizing agent and the reducing agent. 2AgNO Cu Cu(NO3) Ag Rewrite the equation in ionic form 2Ag NO3- + Cu Cu NO Ag

2Ag+ + 2NO3- + Cu Cu2+ + 2NO3- + 2Ag
Sample Problem 2Ag NO3- + Cu Cu NO Ag Oxidation: 2 e- are lost from copper when it becomes a Cu2+ Reduction: 2e- are gained by two silver ions which become neutral silver atoms. oxidizing reducing Agent agent reduced oxidized

Sample Problem 2Na + S 2Na+ + S2- 4Al + 3O2 4Al3+ + 3O2- 2I- I2 + 2e-
oxidized reduced reducing agent oxidizing agent 4Al O Al O2- Oxidized reduced Reducing agent oxidizing agent 2I I e- oxidation Zn e Zn reduction

Redox with Covalent Compounds
Some reactions involve covalent compounds. In these compounds complete electron transfer does not occur. 2H2 (g) + O2 (g) H2O (l) In each reactant hydrogen molecule, the bonding electrons are shared equally between the hydrogen atoms. In water, however, the bonding electrons are pulled toward oxygen because it is much more electronegative than hydrogen.

2H2 (g) + O2 (g) H2O (l) There is a shift of bonding electrons away from hydrogen, even though there is not a complete transfer. Hydrogen is oxidized because it undergoes a partial loss of electrons.

2H2 (g) + O2 (g) H2O (l) In oxygen, the bonding electrons are share equally between oxygen atoms in the reactant oxygen molecule. When oxygen bonds to hydrogen in the water molecule, there is a shift of electrons toward oxygen. Oxygen is thus reduced because it undergoes a partial gain of electrons.

H H O O H O shift of bonding e- shared e- shared e- away from H equally equally H and toward O H is reducing O is oxidizing agent agent

Some reactions involving covalent reactants or products, the partial electron shifts are less obvious. General guideline for covalent reactants or products: for carbon compounds, the addition of oxygen or the removal of hydrogen is always oxidation

Processes Leading to Oxidation & Reduction
Complete loss of electrons (ionic reactions) Complete gain of electrons (ionic reactions) Shift of electrons away from an atom in a covalent bond Shift of electrons toward an atom in a covalent bond Gain of oxygen Loss of oxygen Loss of hydrogen by a covalent compound Gain of hydrogen by a covalent compound Increase in oxidation number Decrease in oxidation number

Corrosion Iron, often used in the form of the alloy steel,
corrodes by being oxidized to ions of iron by oxygen. 2Fe(s) + O2 (g) + 2H2O (l) Fe(OH)2 (s) 4Fe(OH)2(s) + O2 (g) + 2H2O (l) Fe(OH)3 (s) Equations describe the corrosion of iron to iron hydroxides in moist conditions.

Corrosion Water in the environment accelerates the rate of corrosion.
Corrosion occurs more rapidly in the presence of salts and acids. Salts and acids produce electrically conducting solutions that make electron transfer easier.

Resistance to Corrosion
Not all metals corrode easily. Gold and platinum are called noble metals because they are very resistant to losing their e- by corrosion. Other metals lose electrons easily but are protected from extensive corrosion by the oxide coating formed on their surface. Aluminum oxidizes quickly in air to form a coating of very tightly packed aluminum oxide particles.

Resistance to Corrosion
Iron also forms a coating when it corrodes But the coating of iron oxide that forms is not tightly packed. Water and air can penetrate the coating and attack the iron metal below it. Corrosion continues until the iron object becomes only a pile of rust.

Controlling Corrosion
To prevent corrosion, the metal surface can be coated with oil, paint, plastic or another metal. Coatings exclude air and water from the surface, preventing corrosion. If coating is scratched or worn away, however, the exposed metal will begin to corrode.

Another method of corrosion control One metal is “sacrificed” or allowed to corrode, in order to save a second metal. To protect an iron object, a piece of magnesium (or other active metal) may be placed in electrical contact with the iron. When oxygen and water attack the iron object, the iron atoms lose electrons as the iron being to be oxidized.

Because magnesium is a better reducing agent than iron and is more easily oxidized the magnesium immediately transfers electrons to the iron, preventing their oxidation to iron ions. Magnesium is “sacrificed” by oxidation and protects the iron in the process.

Sacrificial zinc and magnesium blocks are sometimes attached to piers and ship hulls to prevent corrosion damage in areas submerged in water. Underground pipelines and storage tanks may be connected to magnesium block for protection It is easier and cheaper to replace a block of magnesium or zinc than to replace a bride or a pipeline.

Question Can you identify the common chemical characteristic of all metal corrosion? The transfer of electrons from metals to oxidizing agents.

Questions Define oxidation and reduction in terms of the gain or loss of oxygen. Oxidation is the gain of oxygen Reduction is the loss of oxygen Define oxidation and reduction in terms of the gain or loss of electrons. LEO the lion goes GER Loss of Electrons is Oxidation Gain of Electrons is Reduction

Iron atoms are oxidized when iron corrodes
Questions What happens to the atoms in an iron nail that corrodes? Iron atoms are oxidized when iron corrodes How do you identify the oxidizing agent and the reducing agent in a redox reaction? The species reduced is the oxidizing agent. The species oxidized is the reducing agent.

Questions Use electron transfer or electron shift to identify what is oxidized and what is reduced in each reaction. (use electronegativity values for molecular compounds) 2Na(s) + Br2(l) NaBr(s) Na oxidized, Br2 reduced H2(g) + Cl2(g) HCl(g) H2 oxidized, Cl2 reduced

Questions Use electron transfer or electron shift to identify what is oxidized and what is reduced in each reaction. (use electronegativity values for molecular compounds) 2Li(s) + F2(g) LiF(s) Li oxidized, F2 reduced S(s) + Cl2(g) SCl2(g) S oxidized, Cl2 reduced

Mg(s) + Cu(NO3)2(aq) Mg(NO3)2(aq) + Cu(s)
Questions Use electron transfer or electron shift to identify what is oxidized and what is reduced in each reaction. (use electronegativity values for molecular compounds) N2(g) + 2O2(g) NO2(s) N2 oxidized, O2 reduced Mg(s) + Cu(NO3)2(aq) Mg(NO3)2(aq) + Cu(s) Mg oxidized, Cu reduced

Questions Identify the reducing agent and the oxidizing agent for each reaction. 2Na(s) + Br2(l) NaBr(s) Na reducing agent, Br2 oxidizing agent H2(g) + Cl2(g) HCl(g) H2 reducing agent, Cl2 oxidizing agent

Questions Identify the reducing agent and the oxidizing agent for each reaction. 2Li(s) + F2(g) LiF(s) Li reducing agent, F2 oxidizing agent S(s) + Cl2(g) SCl2(g) S reducing agent, Cl2 oxidizing agent

Questions Identify the reducing agent and the oxidizing agent for each reaction. N2(g) + 2O2(g) NO2(s) N2 reducing agent, O2 oxidizing agent Mg(s) + Cu(NO3)2(aq) Mg(NO3)2(aq) + Cu(s) Mg reducing agent, Cu oxidizing agent

End of Section 20.1

Oxidation Numbers Oxidation number is a + or – number assigned to an atom to indicate its degree of oxidation or reduction. General Rule A bonded atom’s oxidation # is the charge that it would have if the e- in the bond were assigned to the atom of the more electronegative element.

Rules for Assigning Oxidation Numbers
The oxidation number of a monatomic ion is = in magnitude and sign to its ionic charge. Bromide (Br1-) is -1 Iron III (Fe3+) is +3 The oxidation number of hydrogen in a compound is +1, except in metal hydrides, such as NaH, where it is -1 The oxidation number of oxygen in a compound is -2 except in peroxides, such as H2O2, where it is -1 and in compounds with the more electronegative fluorine, where it is positive.

Rules for Assigning Oxidation Numbers
The oxidation number of an atom in uncombined (elemental) form is 0. Potassium metal (K) is Nitrogen Gas (N2) is 0 For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 0 For polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.

The numbers on a sports player’s jersey
Some Thought Determining oxidation numbers of elements in compounds is a way for chemists to keep track of electron transfer during redox reactions What are other examples where items are numbered to keep track of movement? The numbers on a sports player’s jersey The area codes assigned to telephone numbers in different regions

Binary Ionic Compounds
In binary ionic compounds, such as NaCl and CaCl2, the oxidation numbers of the atoms equal their ionic charges. Na1+ + Cl NaCl oxidation # neutral Ca2+ + Cl CaCl2 oxidation # Note the sign I put before the oxidation number

Molecular Compounds No ionic charges are associated with atoms of molecular compounds. However, oxygen is reduced in the formation of water for example. In water the two shared e- in the H – O bond are shifted toward the O and away from the H. Imagine the e- contributed by the two H atoms are completely transferred to the O.

Molecular Compounds The charge that would result from the transfer are the oxidation numbers of the bonded elements. The oxidation number of O is -2 and the oxidation number of each hydrogen is +1 Oxidation numbers are often written above the chemical symbols in a formula. +1 -2 H2O

Multiple Oxidation Numbers
Many elements can have several different oxidation numbers. K2CrO4 – Potassium Chromate K2CrO7 – Potassium Dichromate

Sample Problems What is the oxidation number of each kind of atom in the following ions and compounds? +4 -2 SO2 CO32- Na2SO4 (NH4)2S

Sample Problems Determine the oxidation number of each element in the following. +3 -2 S2O3 Na2O2 P2O5 NO3-

Sample Problems Determine the oxidation number of chlorine in each of the following substances. KClO3 Cl2 Ca(ClO4)2 Cl2O

Sample Problems What are the oxidation numbers of iodine in the following? HIO4 HIO3 HIO I2 HI

Oxidation Number Changes
When copper wire is placed in a solution of silver nitrate the following reaction occurs 2AgNO3(aq) + Cu(s) Cu(NO3)2(aq) + 2Ag(s) Ag is reduced from +1 to 0 Cu is oxidized from 0 to +2

Oxidation Number Changes
An increase in the oxidation number of an atom or ion indicates oxidation. A decrease in the oxidation number of an atom or ion indicates reduction. 2CuSO4(aq) + Fe2(s) Cu(s) + 2FeSO4 Cu is reduced from +2 to 0 Fe is oxidized from 0 to +2

Sample Problem Identify which atoms are oxidized and which are reduced in the following reaction. Also identify the oxidizing agent and the reducing agent. 2HBr(aq) + Cl2(g) HCl(aq) + Br2(l) Cl is reduced from 0 to -1, so Cl2 is the oxidizing agent Br is oxidized from -1 to 0, so Br1- is the reducing agent

Sample Problem Identify which atoms are oxidized and which are reduced in each reaction O2(g) + 2H2(g) H2O(l) O2 is reduced from 0 to -2 H2 is oxidized from 0 to +1

Sample Problem Identify which atoms are oxidized and which are reduced in each reaction 2KNO3(s) KNO2(s) + O2(g) N is reduced from +5 to +3 O is oxidized from -2 to 0

Sample Problem Identify the oxidizing agent and the reducing agent in each equation. O2(g) + 2H2(g) H2O(l) O2 is reduced, thus O2 is the oxidizing agent H2 is oxidized, thus H2 is the reducing agent

Sample Problem Identify the oxidizing agent and the reducing agent in each equation. 2KNO3(s) KNO2(s) + O2(g) N is reduced, thus N is the oxidizing agent O is oxidized, thus O is the reducing agent

Questions What is the general rule for assigning oxidation numbers?
The oxidation number is the charge a bonded atom would have if the electrons in the bond were assigned to the more electronegative element How is a change I oxidation number related to the process of oxidation and reduction? An increase in oxidation number indicates oxidation; a decrease in oxidation number indicates reduction.

Sample Problem Identify which atoms are oxidized and which are reduced in each reaction. Also identify the oxidizing agent and the reducing agent. 2Na(s) + Cl2(g) NaCl(s) Cl2 is reduced, thus Cl2 is the oxidizing agent Na is oxidized, thus Na is the reducing agent

Sample Problem Identify which atoms are oxidized and which are reduced in each reaction. Also identify the oxidizing agent and the reducing agent. 2HNO3(aq) + 6HI(aq) NO(g) + 3I2(s) + 4H2O(l) O2 is reduced, thus O2 is the oxidizing agent H2 is oxidized, thus H2 is the reducing agent

End of section 20.2

Identifying Redox Reactions
In general, all chemical reaction can be assigned to one of two classes Redox reactions in which electrons are transferred from one reacting species to another. Many single-replacement reactions, combination reactions, decomposition reactions and combustion reactions are redox reactions. All other reactions in which no electron transfer occurs. Double-replacement reactions and acid-base reactions are not redox reactions

What Kind of Reaction Is It?
2Mg(s) + O2(g) MgO(s) Combination reaction – two or more substances react to form a single new substance 2HgO (s) Hg (l) + O2 (g) Decomposition reaction – a single compound breaks down into two or more simpler products. Many are redox reactions

Zn(s) + Cu(NO3)2(aq) Cu(s) + Zn(NO3)2 (aq) Single Replacement Reaction – one element replaces a second element in a compound. 2C6H18(l) + 25O2 (g) CO2(g) H2O(l) Combustion Reactions – an element or a compound reacts with oxygen often producing energy in the form of heat and light.. Many are redox reactions

Na2S(aq) + Cd(NO3)2(aq) CdS(s) + 2NaNO3(aq) Double Replacement Reaction – involving an exchange of positive ions between two compounds. 2C6H18(l) + 25O2 (g) CO2(g) H2O(l) Combustion Reactions – an element or a compound reacts with oxygen often producing energy in the form of heat and light.. Many are not redox reactions

Identifying Redox Reactions
If the oxidation number of an element in a reacting species change, then that element has undergone either oxidation or reduction. Many reactions in which color changes occur are redox reactions.

Balancing Redox Equations
Many redox reaction are too complex to be balanced by trial and error. Two systematic methods are available to balance redox reactions The two methods are based on the fact that the total number of electrons gained in reduction must equal the total number of electrons lost in oxidation. One method used oxidation number changes, and the other used half reactions.

Using Oxidation Number Changes to Balance Redox Equations
Oxidation Number Method You balance by comparing the increases and the decreases in oxidation numbers. Fe2O3(s) + CO(g) Fe(S) + CO2(g) Step 1 – assign oxidation numbers to all the atoms in the equation.

Oxidation Number Method C oxidized (+2) Fe2O3(s) + CO(g) Fe(S) + CO2(g) Fe reduced (-3) Step 2 – Identify which atoms are oxidized and which are reduced

Oxidation Number Method (+2) oxidized Fe2O3(s) + CO(g) Fe(S) + CO2(g) (-3) reduced Step 3 – Use one bracketing line to connect the atoms that undergo oxidation and another such line to connect those that undergo reduction.

In a balanced redox equation, the total increase in oxidation number of the species oxidized must be balance by the total decrease in the oxidation number of the species reduced.

Oxidation Number Method 3 x (+2) = +6 Fe2O3(s) + 3CO(g) Fe(S) + 3CO2(g) 2 x (-3) = -6 Step 4 – Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients.

Oxidation Number Method Fe2O3(s) + 3CO(g) Fe(S) + 3CO2(g) Step 5 – Make sure that the equation is balanced for both atoms and charge.

Sample Problem Oxidation Number Method
S oxidized (+4) x 3= +12 2K2Cr2O7(aq) + H2O(l) + 3S(s) KOH(aq) + 2Cr2O3(s) + 3SO2 (g) Cr reduced (-3) x 4 = -12 4 Cr atoms must be reduced for each 3 S atom that are oxidized

2K2Cr2O7) + 2H2O + 3S 4KOH + 2Cr2O3 + 3SO2
Sample Problem Oxidation Number Method S oxidized (+4) x 3= +12 2K2Cr2O7) + 2H2O + 3S KOH + 2Cr2O3 + 3SO2 Cr reduced (-3) x 4 = -12 Check the equation and balance by inspection 4 in front of KOH balances potassium 2 in front of H2O balances hydrogen and oxygen.

Sample Problems Balance each redox equation using the oxidation number change method KClO3(s) KCl (s) + O2 (g) 2KClO3(s) KCl (s) + 3O2 (g) HNO2 (aq) + HI (aq) NO (g) + I2 (s) + H2O (l) 2HNO2 (aq) + 2HI (aq) NO (g) + I2 (s) + 2H2O (l)

Sample Problems Balance each redox equation using the oxidation number change method Bi2S3 + HNO Bi(NO3)3 + NO + S + H2O Bi2S3 + 8 HNO Bi(NO3) NO + 3S + 4H2O SbCl5 + KI SbCl3 + KCl + I2 SbCl KI SbCl KCl + I2

Using Half Reactions to Balance Redox Equations
Half Reaction Method Half reaction is an equation showing just the oxidation or just the reduction that takes place You write and balance the oxidation and reduction half reaction separately before combining them into a balanced redox equation Then you balance the electrons gained in the reduction with the electrons lost in the oxidation.

Half Reaction Method S + HNO SO2 + NO + H2O (unbalanced) S + H+ + NO SO2 + NO + H2O In this case only HNO3 is ionized. The products are covalent compounds Step 1 – write the unbalanced equation in ionic form

Half Reaction Method Oxidation Half Reaction S SO2 Reduction Half Reaction NO NO Step 2 – Write separate half reactions for the oxidation and reduction processes.

Half Reaction Method S SO2 Sulfur is already balanced, but oxygen is not. The reaction takes place in acid solution, so H2O and H+ are present and can be used to balance oxygen and hydrogen as needed. 2H2O + S SO H+ Step 3 – Balance the atoms in the half reactions.

Half Reaction Method NO NO Nitrogen is already balanced, but oxygen is not. 4H+ + NO NO + 2H2O

Half Reaction Method 2H2O + S SO H+ + 4e - S is oxidized going from 0 to +4, a loss of 4 e- 3e H+ + NO NO + 2H2O N is reduced going from +5 to +2, a gain of 3 e- Step 4 – Add enough electrons to one side of each half reaction to balance the charges.

Half Reaction Method (x 3) 6H2O + 3S SO H+ + 12e - S is oxidized going from 0 to +4, a loss of 4 e- (x 4) 12e H NO NO + 8H2O N is reduced going from +5 to +2, a gain of 3 e- Step 5 – Multiply each half reaction by an appropriate number to make the numbers of electrons equal in both.

Half Reaction Method 6H2O + 3S SO H+ + 12e - 12e H NO NO + 8H2O 6H2O + 3S + 16H+ + 4NO e SO H+ + 4NO + 8H2O + 12e – 3S + 4H+ + 4NO SO2 + 4NO + 2H2O Step 6 – Add the balanced half reactions to show an overall equation and then subtract the terms that appear in both sides of the equation

Half Reaction Method 3S + 4HNO SO2 + 4NO + 2H2O Spectator ion – are present during a reaction, but do not participate in or change during a reaction. Because none of the ions in the reactants appear in the products, there are no spectator ions in this particular example. Step 7 – Add the spectator ions and balance the equation.

Sample Problem Half Reaction Method
KMnO4 + HCl MnCl2 + Cl2 + H2O + KCl K+ + MnO4- + H+ + Cl Mn2+ + 2Cl- + Cl2 + H2O + K+ + Cl- Step 1 – write the unbalanced equation in ionic form

Sample Problem Half Reaction Method Reduction Half Reaction MnO4- Mn2+
MnO Mn2+ Oxidation Half Reaction 2Cl Cl2 Step 2 – Write separate half reactions for the oxidation and reduction processes.

Sample Problem Reduction Half Reaction 8H+ + MnO4- Mn2+ + 4H2O
8H+ + MnO Mn H2O Oxidation Half Reaction 2Cl Cl2 Solution is acidic, so H2O and H+ ions are used to balance equation. (If solution is basic, H2O and OH- are used) Step 3 – Balance the atoms in the half reactions.

Sample Problem Half Reaction Method 5e- + 8H+ + MnO4- Mn2+ + 4H2O
5e H+ + MnO Mn H2O Mn is reduced going from +7 to +2, a gain of 5 e- 2Cl Cl2 + 2e- Cl is oxidized going from -1 to 0, a loss of 2 e- Step 4 – Add enough electrons to one side of each half reaction to balance the charges.

Sample Problem Half Reaction Method
(x 2) 10e H+ + 2MnO Mn H2O Mn is reduced going from +7 to +2, a gain of 5 e- (x5) 10Cl Cl2 + 10e- Cl is oxidized going from -1 to 0, a loss of 2 e- Step 5 – Multiply each half reaction by an appropriate number to make the numbers of electrons equal in both.

16H+ + 2MnO4- + 10Cl- 2Mn2+ + 8H2O + 5Cl2
Sample Problem Half Reaction Method 10e H+ + 2MnO Mn H2O 10Cl Cl2 + 10e- 10e H+ + 2MnO Cl Mn H2O + 5Cl2 + 10e- 16H+ + 2MnO Cl Mn H2O + 5Cl2 Step 6 – Add the balanced half reactions to show an overall equation and then subtract the terms that appear in both sides of the equation

Sample Problem Half Reaction Method 16H+ + 6Cl- + 2MnO4- + 2K+ + 10Cl-
5Cl2 + 2Mn Cl- + 8H2O + 2K+ + 2Cl- 16H+ = 6Cl- + 10Cl- (from the HCl in original equation) 2MnO4- = 2K+ (from the KMnO4 in original equation) 2Mn2+ = 4Cl- (from the MnCl2 in original equation) 2K+ + 2Cl- (from the KCl in original equation) Step 7 – Add the spectator ions and balance the equation.

Sample Problem Half Reaction Method 16H+ + 16Cl- + 2MnO4- + 2K+
5Cl2 + 2Mn Cl- + 8H2O + 2K+ Adding spectator and non-spectator Cl- on each side gives 16HCl + 2KMnO Cl2 + 2MnCl2 + 8H2O + 2KCl Step 7 – Add the spectator ions and balance the equation.

4Zn + NO3- + 6H2O + 7OH- 4Zn(OH4)2- + NH3
Sample Problem 2 The following takes place in basic solution. Balance the equation using the half reaction method Zn + NO NH3 + Zn(OH4)2- 4Zn + NO3- + 6H2O + 7OH Zn(OH4)2- + NH3

6Zn + As2O3 + 9H2O 6Zn2+ + 2AsH3 + 12OH-
Sample Problem 3 The following takes place in basic solution. Balance the equation using the half reaction method Zn + As2O AsH3 + Zn2+ 6Zn + As2O3 + 9H2O Zn AsH OH-

Choosing a Balancing Method
In some redox reactions, the same element is both oxidized and reduced. (called a disproportional reaction) 3Br2 + 6KOH KBr + KBrO3 + 3H2O Br is reduced from 0 to -1 Br is oxidized from 0 to + 5 Equations like above and equation for reactions that take place in acidic or alkaline solution are best balanced by the half reaction method.

End of Chapter 20

Chapter 20 Oxidation-Reduction Reactions

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