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Unit 12: Oxidation, Reduction, and Electrochemistry
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Redox REDuction – OXidation Reactions (AKA Redox): rxns that involve the _____________________________; both reduction and oxidation must happen ____________________! Reduction = ___________________ by an atom or ion; __________ ___________ goes ________/____________ Oxidation = ___________________ by an atom or ion; __________ ___________ goes ________/_____________ Transfer of electrons simultaneously Gain of electrons Oxidation Number Down Reduces Loss of electrons Oxidation Number up increases
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A way to remember → L E O the lion goes G E R
A way to remember → L E O the lion goes G E R *Oxidation and reduction happen because of the __________ for electrons in a chemical reaction. Species prefer to either ________ or _________ electrons in a chemical reaction. AND **Oxidation and reduction are __________ or __________________ reactions and one cannot happen without the other. If one atom _______ electrons, there must be another atom that will _______ electrons. Lose e- oxidation Gain e- reduction desire lose gain mutual simultaneous loses gain
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Green Chunk Experiment: http://www.sciencedojo.com/?p=184
Example: ___ Al + ___ CuCl2 → ___ ___________ ___ ______ Aluminum is above Cu on Table J so it will replace it! Notice how Al is all by itself (on left of arrow) with a zero charge and then bonded (on right of arrow) where it takes on a charge 2 3 2 AlCl3 3 Cu
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IDENTIFYING REDOX REACTIONS
One way that we can begin to identify a redox reaction is to inspect the _________________________ from reactant to product side (for every element involved in the reaction). Oxidation numbers are used to track the __________________________ (electron transfer) from reactant to product side of rxn Oxidation Number (State) = ____________, ____________, OR ____________ (_______) values that can be assigned to atoms; identify how many electrons are being lost or gained by an atom/ion when they ____________________ Oxidation numbers Movement of electrons positive negative neutral zero Form bonds
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*top listed # to the upper right is the most common oxidation number for that element
Trick 1: _________________________ reactions are always REDOX! Example: Zn + HCl → ___ _________ ___ _________ *___/___ are by themselves on one side and bonded on the opposite side Trick 2: DOUBLE REPLACEMENT REACTIONS are NOT REDOX! NaOH + HCl → ___ _________ ___ _________ *charges stay the same for all elements in the rxn Single replacement reactions ZnCl2 H2 Zn H NaCl HOH
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Rules for assigning OXIDATION STATES (numbers):
1) _______________________________ (elements not bonded to another element) have an oxidation number of ________. This includes any formula that has only one chemical symbol in it (single elements & diatomic elements). Examples: _______ _______ _______ _______ 2) In ___________, the sum of the _________ for all elements must ____________________. Ex: NaCl Ex: Mg3N Ex: HNO3 Na: 1(+1) = Mg: H: Cl: 1(-1) = N: N: 0! O: Uncombined elements zero Al(s)0 Na(s)0 Cl2(g)0 H2(g)0 compounds charges Add up to zero 3(+2) = +6 2(-3) = -6 0!
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* The ________________________ is the number ___________ the _______________. It is the charge on _____ atom of that element! ** Trick: You can keep polyatomic ions together and use the charge from Table E to determine the oxidation numbers for those elements. *** Remember that we almost always write the __________________ and the ___________________ in a compound formula. EXAMPLE: EXCEPTION to this rule: Oxidation number inside parentheses one (+) element first (-) element last HCl NH3
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3) ___________________ always have a ___ oxidation number when in a compound (bonded to another species). ___________________ always have a ___ oxidation number when located within a compound. (Assign) Ox #: Ex: LiCl MgCl2 4) __________ is always a ___ in compounds. The other __________ (ex: Cl, Br, I) are also ___ as long as they are the most electronegative element in the compound. Ex: HF CaCl2 NaBr Group 1 Metals +1 Group 2 Metals +2 Fluorine -1 Halogens -1
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5) _____________ is a ___ in compounds unless it is combined with __________ or __________________, in which case it is ___. (Assign) Ox #: Ex: HCl LiH 6) _____________ is ________________ in compounds. Ex: H2O Hydrogen +1 Group 1 Group 2 Metal -1 Oxygen Usually -2
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When combined with ___________ (__), which is more electronegative, ___________________. (Assign) Ox #: Ex: OF2 When in a __________________________. A peroxide is a compound that has a formula of ________. Ex: Na2O2 H2O2 Fluorine F Oxygen is +2 Peroxide, oxygen is -1 X2O2
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7) The sum of the oxidation numbers in polyatomic ions must equal the ___________________________ (_________________). Ex: Cr2O72- a Cr: 2( ) = O: 7( ) = = (charge on ion) Charge on the ion See table E +6 +12 -2 -14 -2
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A reaction is REDOX if…___________________________ _________________________________________________ Reduction (GER) = ________________________ by an atom or ion; _________________________ goes _______/______________ Oxidation (LEO) = ________________________ by an atom or ion; _________________________ goes _______/______________ Oxidation numbers change for two elements within a reaction Gain of electrons Oxidation number down reduces Loss of electrons Oxidation number up oxidizes
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Assign oxidation numbers for all elements and complete the tables: Example 1: C + H2O → CO + H2 Example 2: MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O Charge: Increases/Decreases e-: Lost/Gained Oxidized/Reduced C0 H+1 increases lost oxidized decreases gained reduced Charge: Increases/Decreases e-: Lost/Gained Oxidized/Reduced Cl-1 Mn+4 increases lost oxidized decreases gained reduced
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HALF REACTIONS Half reactions allow us to show the _______________ in a redox rxn. For each redox reaction, we can illustrate two __________. One half-reaction shows ____________ and other shows ____________. Example of a Reduction Half Reaction: Fe3+ + 3e- → Fe *Electrons on left hand side, __________ in the rxn (____). Notice also how the charge for Fe goes down from left to right, ____________ (____). Charge goes down because Fe __________ e-. Exchange of e- Half rxns. oxidation reduction gained GER reduction GER gained
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Example of an Oxidation Half Reaction:
F e → Fe e- *Electrons on left hand side, _________ in the rxn (____). Notice also how the charge for Fe goes up from left to right, _____________ (____). Charge goes up because Fe _______ e-. NOTICE: Always _________________ to the side of rxn that has a _________________ (remember: electrons are ___________!) FOLLOWING THE LAW OF CONSERVATION: Half reactions follow the _____________________________. This means that there must be the _______________________ on both sides of the reaction arrow. There must also be a _____________________________. In half reactions, the ____________________________________ of the equation, although it doesn’t necessarily need to equal zero. lost LEO oxidation LEO lost Add electrons Higher total charge negative Law of conservation of mass Same number of atoms Conservation of charge Net charge must be the same on both sides
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RULES FOR SETTING UP HALF REACTIONS
1) Assign oxidation numbers to all elements in reaction 2) Draw brackets and identify oxidation & reduction 3) Begin to set up half reactions. Pull out brackets bringing element symbol and assigned charge with you. Set up as a reaction with arrow connecting two sides that have different oxidation numbers assigned. Only trick: diatomics must be pulled out as a pair. This is the only time you ever “bring subscripts with you” in creating half reactions! 4) FOR REACTIONS INVOLVING DIATOMIC ELEMENTS ONLY: Balance mass 1st (make sure there are the same number of elements on each side of each half reaction) 5) Lastly, balance charge in each half reaction by inserting appropriate amount of electrons into each half reaction to attain conservation of charge. Always add electrons to side that has a more positive charge. REMEMBER, electrons are negative in nature! Net charges on each side of rxn should be equal after adding electrons.
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Assign oxidation numbers to all elements or polyatomic ions
Assign oxidation numbers to all elements or polyatomic ions. Label the brackets for reduction (red) or oxidation (ox). Ex. 1: Mg + ZnCl2 → MgCl2 + Zn OXIDATION Half Reaction: Mg0 → ______ + ______ REDUCTION Half Reaction: Zn+2 + ______ → ______ Mg+2 2e- 2e- Zn0
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Ex. 2 (balance masses): Hg + I2 → HgI OXIDATION Half Reaction (make sure to balance the masses ): ______ → ______ + ______ REDUCTION Half Reaction: ______ + ______ → ______ 2 Hg0 2 Hg+1 2e- I20 2 e- 2 I-
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Ex. 3 (balance charges): Cu + AgNO3 → Cu(NO3)2 + Ag OXIDATION Half Reaction: ______ → ______ + ______ REDUCTION Half Reaction: ______ + ______ → ______ Now, we need to balance the charges: ___ x (Ag+1 + 1e- → Ag0) = ___Ag+1 + ___e- → ___Ag0 Cu0 Cu+2 2 e- Ag+1 1 e- Ag0 2 2 2 2
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Table J and Spontaneous Reactions
General Rule: elements ___________ on Table J are ________ reactive than the elements below them Spontaneous rxn = rxn occurs w/out adding energy to system If the “single” element is more active than the “combined” element, the reaction will be spontaneous. Non-spontaneous rxn = rxn will not occur unless energy is added to system If the “single” element is less active than the “combined” element, the reaction will NOT be spontaneous. higher more
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Complete the following equations by writing in the products formed or “no rxn” Ex 1: Zn + PbCl2 → Ex 2: Zn + BaO → Ex 3: Ca + CrF2 → Ex 4: Mn + NiS → Ex 5: Fe + MgI2 → Ex 6: Co + PbCl2 → ZnCl2 + Pb0 No rxn CaF2 + Cr0 MnS + Ni0 No rxn CoCl2 + Pb0
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TWO TYPES of ELECTROCHEMICAL CELLS
Voltaic (similar to a battery) Electrolytic (similar to alternator in cars) SIMILARITIES BETWEEN THE TWO: Both involve ___________ reactions; chemical reactions which involve the flow of ______________ Both involve the flow of __________________, or __________, measured in _________ Both have ______________ (conductive surfaces where oxidation or reduction occurs); called the _________ and the _________ _______________ or _______________ occurs in each half cell Redox electrons Electrical energy current volts 2 electrodes anode cathode oxidation reduction
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RED CAT AN OX ______________________ ______________________
______________________ ______________________ (__________________) (__________________) Electrons flow through the _______ from the _________ to the _________. Reduction ALWAYS occurs at the cathode Oxidation ALWAYS occurs at the anode Ions gain e- Metal loses e- wire anode cathode
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Voltaic Cells Cells that _____________________ convert _____________ energy into ________________ energy or electric _____________. ____________________ CATHODE The ______________________ of the 2 metals (Table J) _______________________________________ to it the _____________ electrode in a ______________________ electrode where _______________ occurs (_____________) ANODE _______________________________________ to cathode the ______________ electrode in a ______________________ spontaneously chemical electrical current batteries Less active Spontaneously attracts electrons positive Voltaic cell reduction RED CAT More active Spontaneously loses electrons negative Voltaic cell oxidation AN OX
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Example 1: Wet Cell ______________, are a form of ______________ battery Consists of LEAD ANODE and LEAD OXIDE CATHODE Both electrodes immersed in a ________________ solution Advantage: process is readily ____________ (by alternator) Disadvantage: very ________, somewhat ______________ Example 2: Dry Cell _______________________ are the type of batteries in a portable radio, remote control, etc. CARBON (GRAPHITE) CATHODE surrounded by moist electrolyte paste Usually ZINC ANODE *SALT BRIDGE provides a path for the _________________ between the half-cells prevents the __________________________ Car batteries Lead storage Sulfuric acid reversible heavy dangerous Dry cell batteries Flow of ions Build-up of charge
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Voltaic Cells (a.k.a Galvanic Cells)
→ 1. Use Table J to predict the direction that electrons will spontaneously flow. Draw arrows to indicate the direction on the wire. 2. Based on your answer above, which would be the negative electrode and which would be the positive electrode? ____________________________________________________________ 3. Explain your answer to #2. ____________________________________________________ ____________________________________________________________________________ ____________________________________________________________________________ Pb is negative electrode and Ag is positive electrode Ag is gaining electrons and therefore reduced which is the cathode and +, Pb is losing electrons and therefore oxidized which is the anode and -.
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4. At which electrode or in which half-cell does reduction occur? _____
5. At which electrode or in which half-cell does oxidation occur? _____ 6. Which electrode is the cathode? _____ 7. Which electrode is the anode? _____ *Electrons don’t flow to the cathode, they flow through it to the ions in solution. That’s why the cathode never becomes negative 2 1 2 1
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Electrolytic Cells nonspontaneous
Cells that use _________________ to force a ___________________ ____________________________ to occur. This process is for __________________ and ___________________ Example: _________________________ (keeps the car battery replenished with energy) Electrical energy nonspontaneous Chemical reaction electrolysis electroplating Alternator in car
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Electrolysis Experiment Animation (w/ tutorial):
Standard Hydrogen Electrode (Zinc) Standard Hydrogen Electrode (Copper)
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Electrolytic Cells CATHODE
electrode where _____________ are __________ the ________________ electrode (opposite of voltaic cell) electrode where _______________ occurs (____________) ANODE electrode where _____________ are _______________________ NOTICE: There is ___________________. This is a forced chemical reaction. You will always see a _______________ hooked up to an electrolytic cell which drives the ____________________ electrons sent negative reduction RED CAT electrons Drawn away from positive oxidation AN OX No salt bridge Power source Forced rxn
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Flow of e- (spontaneous or forced)
Compare and contrast the two types of electrochemical cells: GALVANIC/VOLTAIC ELECTROLYTIC Flow of e- (spontaneous or forced) (+) electrode (-) electrode *Direction of e- flow Reduction ½ cell Oxidation ½ cell spontaneous forced cathode anode anode cathode Anode → cathode Anode → cathode Cathode Cathode anode anode *Direction of e- flow is either “Anode → Cathode” or “Cathode → Anode”
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Nonspontaneous Reactions on Electrolytic Cells
Electroplating (see electroplating video)
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