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Review of Terminology Review of Terminology for Redox Reactionsfor Redox Reactions
Review of Terminology Review of Terminology for Redox Reactionsfor Redox Reactions
• OXIDATIONOXIDATION—loss of electron(s) by a —loss of electron(s) by a species; increase in oxidation number.species; increase in oxidation number.
• REDUCTIONREDUCTION—gain of electron(s); —gain of electron(s); decrease in oxidation number.decrease in oxidation number.
• OXIDIZING AGENTOXIDIZING AGENT—electron acceptor; —electron acceptor; species is reduced.species is reduced.
• REDUCING AGENTREDUCING AGENT—electron donor; —electron donor; species is oxidized.species is oxidized.
• OXIDATIONOXIDATION—loss of electron(s) by a —loss of electron(s) by a species; increase in oxidation number.species; increase in oxidation number.
• REDUCTIONREDUCTION—gain of electron(s); —gain of electron(s); decrease in oxidation number.decrease in oxidation number.
• OXIDIZING AGENTOXIDIZING AGENT—electron acceptor; —electron acceptor; species is reduced.species is reduced.
• REDUCING AGENTREDUCING AGENT—electron donor; —electron donor; species is oxidized.species is oxidized.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Some fruits, including apples, turn brown when you cut them. What is happening on the surface of the fruit?
CHEMISTRY & YOUCHEMISTRY & YOU
Oxygen in air reacts with chemicals on the surface of
the cut fruit. The oxygen oxidizes the chemicals in the fruit, causing a redox reaction and therefore the
color change.
Why does cut fruit turn brown?
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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• The element whose oxidation number increases is oxidized
• The element whose oxidation number decreases is reduced
If changes in oxidation number occur, the reaction is a redox reaction.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Identifying Redox Reactions
Use the change in oxidation number to identify whether each reaction is a redox
reaction
Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Sample Problem 20.5Sample Problem 20.5
a. Assign oxidation numbers.
0 +1 –1 +1 –1 0
Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
The chlorine is reduced;
The bromide ion is oxidized;
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Solve ASSIGN OXIDATION NUMBERS. 2
Sample Problem 20.5Sample Problem 20.5
• NO change in oxidation number.
• This is not a redox reaction.
+1 –2 +1 +1 +6 –2 +1 +6 –1 +1 –2
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
This is an acid-base (neutralization) reaction.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Which of the following are redox reactions?
A. NH3 + HCl → NH4Cl
B. SO3 + H2O → H2SO4
C. NaOH + HCl → NaCl + H2O
D. H2S + NHO3 → H2SO4 + NO2 + H2O
B. SO3 + H2O → H2SO4
D. H2S + NHO3 → H2SO4 + NO2 + H2O
99
TRANSFER REACTIONSTRANSFER REACTIONS
Atom/Group transfer (not)
HCl + H2O ---> Cl- + H3O+
Redox: Electron transfer
Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)
1010
OXIDATION-REDUCTION REACTIONS
OXIDATION-REDUCTION REACTIONS
Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)
Why 2?
1111Balancing Balancing EquationsEquations
Need to Balance BOTH mass and CHARGE
Cu + Ag+ --give--> Cu2+ + Ag
Step 1:: Divide into half-reactions:
one for oxidation and the other for reduction.
Ox Cu ---> Cu2+
Red Ag+ ---> Ag
1212Balancing Balancing EquationsEquations
Step 2: Balance each for mass.
Already done in this case.
Step 3:Balance each half-reaction for
charge by adding electrons.
Ox Cu ---> Cu2+ + 2e-
Red Ag+ + e- ---> Ag
1313Balancing Balancing EquationsEquations
Step 4: Multiply each half-reaction by a factor to have the electrons lost equal to number gained
Cu ---> Cu2+ + 2e- 2 Ag+ + 2 e- ---> 2 Ag
Step 5:Add to give the overall equation.
Cu + 2 Ag+ ---> Cu2+ + 2Ag
The equation is now balanced for BOTH charge and mass.
1414
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Use the half-reaction method to balance the following redox equation.
FeCl3 + H2S → FeCl2 + HCl + S
Oxidation: H2S → 2H+ + S + 2e–
Reduction: 2Fe3+ + 2e– → 2Fe2+
2FeCl3 + H2S → 2FeCl2 + 2HCl + S
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Balancing Redox Equations
What are two different methods for balancing a redox equation?
Two different methods for balancing redox equations are the
oxidation-number-change method and the half-reaction
method.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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In the oxidation-number-change method, you balance
a redox equation by comparing the increases and
decreases in oxidation numbers.
Using Oxidation-Number Changes
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Using Oxidation-Number Changes
Step 1: Assign oxidation numbers to all the atoms in the equation.
• Write the numbers above the atoms.
• The oxidation number is stated per atom.
Fe2O3(s) + CO(g) → Fe(s) + CO2(g)
+3 –2 +2 –2 0 +4 –2
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Using Oxidation-Number Changes
Step 2: Identify which atoms are oxidized and which are reduced.
• Iron is reduced. +3 to 0
• Carbon is oxidized. +2 to +4
Fe2O3(s) + CO(g) → Fe(s) + CO2(g)
+3 –2 +2 –2 0 +4 –2
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Using Oxidation-Number Changes
Step 3: Use one line to connect the atoms that undergo oxidation and another such line to connect those that undergo reduction.
–3 (reduction)
Fe2O3(s) + CO(g) → Fe(s) + CO2(g) +3 –2 +2 –2 0 +4 –2
+2 (oxidation)
• Write the oxidation-number change at the midpoint of each line.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Using Oxidation-Number Changes
Step 4: Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients.
2 × (–3) = –6
Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g) +3 –2 +2 –2 0 +4 –2
3 × (+2) = +6
• The oxidation-number increase should be multiplied by 3 and the decrease by 2.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Using Oxidation-Number Changes
Step 5: Finally, make sure the equation is balanced for both atoms and charge.
Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)
• If necessary, finish balancing the equation by inspection.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Balancing Redox Equations by Oxidation-Number Change
Balance this redox equation by using the oxidation-number-change method.
Cl2(g) + K(s) → KCl(s)0 0 +1 –1
2 × (-1) = -2
2 × (+1) = +2
2 2
2424Acid / BaseAcid / Base Redox Reactions Redox Reactions
Some redox reactions have equations that must be balanced by special
techniques.If reactions occur in the “presence of”
acid or base
MnO4- + 5 Fe2+ + 8 H+
---> Mn2+ + 5 Fe3+ + 4 H2O
2525
Step 1:: Divide into half-reactions
Permanganate and iron(II) ions are reacted in an acidic solution
MnO4- + Fe2+ ---> Mn2+ + Fe3+
Ox Fe2+ ---> Fe3+
Red MnO4- ---> Mn2+
Need to balance MASS
2626Step 2: Balance each for mass.
Already done for Iron
Fe2+ ---> Fe3+
Need to have “O” on both sided: Add water
MnO4- ---> Mn2+ + H2O
Never add O2, O atoms, or O2- to balance oxygen.
MnO4- ---> Mn2+ + 4H2O
2727
Step 2: Balance each for mass.
MnO4- ---> Mn2+ + 4H2O
Need to have “H” on both sided.
Told in an acidic solution: need H+
8 H+ + MnO4- ---> Mn2+ + 4H2O
Never add H2 or H atoms to balance hydrogen.
2828Step 3:Balance each half-reaction for
charge by adding electrons.
8 H+ + 5 e- +MnO4- ---> Mn2+ + 4H2O
5Fe2+ ---> 5Fe3+ + 5e-
Multiply each half-reaction by a factor to have the electrons lost
equal to number gained
2929Acid / BaseAcid / Base Redox Reactions Redox Reactions
Step 5: Add to obtain the overall equation
MnO4- + 5 Fe2+ + 8 H+
---> Mn2+ + 5 Fe3+ + 4 H2O
Check by adding charges on both sides and by counting atoms
The equation is now balanced for BOTH charge and mass.
3030
Balancing Balancing EquationsEquations
• Never add O2, O atoms, or O2- to balance oxygen.
• Never add H2 or H atoms to balance hydrogen.
• Be sure to write the correct charges on all the ions.
• Check your work at the end to make sure mass and charge are balanced.
• PRACTICE!
3232Balancing Balancing EquationsEquations
Balance the following in acid solution—
VO2+ + Zn ---> VO2+ + Zn2+
Step 1:Write the half-reactions
Ox Zn ---> Zn2+
Red VO2+ ---> VO2+
Step 2:Balance each for mass.
VO2+ ---> VO2+ 2 H+ +
Add HAdd H22O on O-deficient side and add HO on O-deficient side and add H++
on other side for H-balance.on other side for H-balance.
+ H2O
3333Balancing Balancing EquationsEquations
Step 3: Add electrons to half reaction.
Zn ---> Zn2+ + 2e-
e- + 2 H+ + VO2+ ---> VO2+ + H2O
Step 4:Multiply by an appropriate factor.
2e- + 4 H+ + 2 VO2+ ---> 2 VO2+ + 2 H2O
Zn ---> Zn2+ + 2e-
Balancing EquationsBalancing EquationsStep 5:Add balanced half-reactions
Zn + 4 H+ + 2 VO2+
---> Zn2+ + 2 VO2+ + 2 H2O
Check by adding charges on both sides and by counting atoms
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Balancing Redox Equations by Half-Reactions
Balance this redox equation using the half-reaction
method.
KMnO4(aq) + HCl(l) → MnCl2(aq) + Cl2(g) + KCl(aq)
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Step 1:Write the equation in ionic form.
Sample Problem 20.7Sample Problem 20.7
K+(aq) + MnO4–(aq) + H+(aq) + Cl–(aq) →
Mn2+(a) + 2Cl–(a) + Cl2(g) + H2O + K+(a) + Cl–(a)
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Step 2: Write half-reactions. Determine the oxidation and reduction process.
Oxidation half-reaction:
Cl– → Cl2
Reduction half-reaction:
MnO4– → Mn2+
–1 0
+7 +2
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Step 3: Balance the atoms in each half-reaction.
Oxidation:
2Cl–(aq) → Cl2(g) (atoms balanced)
Reduction:
MnO4–(aq) + 8H+(aq) →
Mn2+(aq) + 4H2O(l) (atoms balanced)
• The solution is acidic, so use H2O and H+ to balance the oxygen and hydrogen.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Step 4: Balance the charges by adding electrons.
Oxidation:
2Cl–(aq) → Cl2(g) + 2e– (charges balanced)
Reduction:
MnO4–(aq) + 8H+(aq) + 5e– →
Mn2+(aq) + 4H2O(l) (charges balanced)
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Step 5: Make the numbers of electrons equal.
Oxidation: 10Cl–(aq) → 5Cl2(g) + 10e–
Reduction:
2MnO4–(aq) + 16H+(aq) + 10e– →
2Mn2+(aq) + 8H2O(l)
• Multiply the oxidation half-reaction by 5 and the reduction half-reaction by 2.
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Step 6: Add the half-reactions. Then, subtract the terms that appear on both sides.
10Cl–(aq) + 2MnO4–(aq) + 16H+(aq) + 10e– →
5Cl2(g) + 10e– + 2Mn2+(aq) + 8H2O(l)
Step 7: Add the spectator ions, making sure the charges and atoms are balanced.
10Cl– + 2MnO4– + 2K+ + 16H+ + 6Cl– →
5Cl2 + 2Mn2+ + 4Cl– + 8H2O + 2K+ + 2Cl–
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Combine the spectator and nonspectator Cl– on each side.
16Cl–(a) + 2MnO4–(a) + 2K+(a) + 16H+(a) →
5Cl2(g) + 2Mn2+(a) + 6Cl–(a) + 8H2O(l) + 2K+(a)
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Show the balanced equation for the substances given in the question (rather than for ions).
2KMnO4(aq) + 16HCl(aq) →
2MnCl2(aq) + 5Cl2(g) + 8H2O(l) + 2KCl(aq)
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Use the half-reaction method to balance the following redox equation.
FeCl3 + H2S → FeCl2 + HCl + S
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Glossary TermsGlossary Terms
• oxidation-number-change method: a method of balancing a redox equation by comparing the increases and decreases in oxidation numbers
• half-reaction: an equation showing either the oxidation or the reduction that takes place in a redox reaction
• half-reaction method: a method of balancing a redox equation by balancing the oxidation and reduction half-reactions separately before combining them into a balanced redox equation
20.3 Describing Redox Equations >20.3 Describing Redox Equations >
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Redox equations can be balanced by two methods, the oxidation-number-change method and balancing the oxidation and reduction half-reactions.
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