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Balancing Redox Reactions Chapter 20: Day 2 and 3
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Review of Terminology for Redox Reactions
OXIDATION—loss of electron(s) by a species; increase in oxidation number. REDUCTION—gain of electron(s); decrease in oxidation number. OXIDIZING AGENT—electron acceptor; species is reduced. REDUCING AGENT—electron donor; species is oxidized.
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Redox: Electron transfer
TRANSFER REACTIONS Atom/Group transfer (not) HCl + H2O ---> Cl H3O+ Redox: Electron transfer Cu(s) Ag+(aq) ---> Cu2+(aq) Ag(s)
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Copper + Silver Ion
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OXIDATION-REDUCTION REACTIONS
Direct Redox Reaction Oxidizing and reducing agents in direct contact. Cu(s) Ag+(aq) ---> Cu2+(aq) Ag(s) Why 2?
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Cu + Ag+ --give--> Cu2+ + Ag Need to Balance BOTH mass and CHARGE
Balancing Equations Cu + Ag+ --give--> Cu Ag Need to Balance BOTH mass and CHARGE Step 1: Divide into half-reactions: one for oxidation and the other for reduction. Ox Cu ---> Cu2+ Red Ag+ ---> Ag
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Already done in this case. charge by adding electrons.
Balancing Equations Step 2: Balance each for mass. Already done in this case. Step 3: Balance each half-reaction for charge by adding electrons. Ox Cu ---> Cu e- Red Ag+ + e- ---> Ag
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Balancing Equations Step 4: Multiply each half-reaction by a factor to have the electrons lost equal to number gained Cu ---> Cu e- 2 Ag e- ---> 2 Ag Step 5: Add to give the overall equation. Cu Ag > Cu Ag The equation is now balanced for BOTH charge and mass.
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5+ 2- 4+ 3- 2+ 2- 2+ 2- 6+ 4- 4+ 4+ 2+ 3+ 3+ 5+ 3- 3- 2-
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Ready for more complex reactions?
YES!
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Acid / Base Redox Reactions
Some redox reactions have equations that must be balanced by special techniques. If reactions occur in the “presence of” acid or base MnO Fe H+ ---> Mn Fe H2O
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Step 1: Divide into half-reactions
Permanganate and iron(II) ions are reacted in an acidic solution MnO4- + Fe > Mn2+ + Fe3+ Step 1: Divide into half-reactions Ox Fe > Fe3+ Red MnO > Mn2+ Need to balance MASS
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Step 2: Balance each for mass. Already done for Iron Fe2+ ---> Fe3+
Need to have “O” on both sided: Add water MnO > Mn2+ + H2O Never add O2, O atoms, or O2- to balance oxygen. MnO > Mn2+ + 4H2O
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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+ + MnO > Mn2+ + 4H2O Never add H2 or H atoms to balance hydrogen.
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charge by adding electrons. 8 H+ + 5 e- +MnO4- ---> Mn2+ + 4H2O
Step 3: Balance each half-reaction for charge by adding electrons. 8 H+ + 5 e- +MnO > Mn2+ + 4H2O 5Fe > 5Fe3+ + 5e- Multiply each half-reaction by a factor to have the electrons lost equal to number gained
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Acid / Base Redox Reactions
Step 5: Add to obtain the overall equation MnO Fe H+ ---> Mn Fe H2O Check by adding charges on both sides and by counting atoms The equation is now balanced for BOTH charge and mass.
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Balancing Equations 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!
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Reduction of VO2+ with Zn
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Balancing Equations Balance the following in acid solution—
VO Zn ---> VO Zn2+ Step 1: Write the half-reactions Ox Zn ---> Zn2+ Red VO > VO2+ Step 2: Balance each for mass. VO > VO2+ + H2O 2 H+ + Add H2O on O-deficient side and add H+ on other side for H-balance.
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Balancing Equations e- + 2 H+ + VO2+ ---> VO2+ + H2O
Step 3: Add electrons to half reaction. Zn ---> Zn e- e H+ + VO > VO H2O Step 4: Multiply by an appropriate factor. 2e H VO > 2 VO H2O
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Balancing Equations Step 5: Add balanced half-reactions Zn H VO > Zn VO H2O Check by adding charges on both sides and by counting atoms
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