Unit 12 Oxidation Reduction and Electrochemistry Redox REDuction

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Unit 12: Oxidation, Reduction, and Electrochemistry

Unit 12: Oxidation, Reduction, and Electrochemistry

Redox • REDuction – OXidation Reactions (AKA Redox): rxns that involve the _______________; both

Redox • REDuction – OXidation Reactions (AKA Redox): rxns that involve the _______________; both Transfer of electrons reduction and oxidation must happen simultaneously __________! • Reduction = __________ Gain of electrons by an Oxidation ______ atom or ion; _____ goes Number ____/______ Down Reduces • Oxidation = __________ by an Loss of electrons atom or ion; _____ goes Number Oxidation ______/_______ up increases

A way to remember → L E O the lion goes Lose e- oxidation

A way to remember → L E O the lion goes Lose e- oxidation G E R Gain e- reduction *Oxidation and reduction happen because of the _____ for electrons in a chemical reaction. desire Species prefer to either ____ or _____ gain lose electrons in a chemical reaction. mutual AND **Oxidation and reduction are _____ simultaneous or _________ reactions and one cannot happen without the other. If one atom loses electrons, there must be another atom _______ that will _______ gain electrons.

Green Chunk Experiment: http: //www. sciencedojo. com/? p=184 Example: 2 Al + ___ 3

Green Chunk Experiment: http: //www. sciencedojo. com/? p=184 Example: 2 Al + ___ 3 Cu. Cl 2 _______ Al. Cl 3 3 ______ Cu → ___ + ___ 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

IDENTIFYING REDOX REACTIONS One way that we can begin to identify a redox reaction

IDENTIFYING REDOX REACTIONS One way that we can begin to identify a redox reaction is to inspect the Oxidation numbers _____________ from reactant to product side (for every element involved in the reaction). Oxidation numbers are used to track the _____________ (electron Movement of electrons transfer) from reactant to product side of rxn positive Oxidation Number (State) = ______, negative neutral ______, OR ______ (_______) zero values that can be assigned to atoms; identify how many electrons are being lost or gained by an atom/ion when they __________ Form bonds

*top listed # to the upper right is the most common oxidation number for

*top listed # to the upper right is the most common oxidation number for that element Single replacement reactions Trick 1: _____________ are always REDOX! Example: Zn + HCl → _________ + _________ H 2 Zn. Cl 2 Zn H are by themselves on one side and *___/___ bonded on the opposite side Trick 2: DOUBLE REPLACEMENT REACTIONS are NOT REDOX! Example: Na. OH + HCl → _________ + _________ Na. Cl HOH *charges stay the same for all elements in the rxn

Rules for assigning OXIDATION STATES (numbers): 1) ________________ (elements not Uncombined elements bonded to

Rules for assigning OXIDATION STATES (numbers): 1) ________________ (elements not Uncombined elements bonded to another element) have an oxidation number of ____. This includes any formula that has only one zero chemical symbol in it (single elements & diatomic elements). Examples: Al(s)0 _______ Na(s)0 _______ Cl 2(g)0 _______ H 2(g)0 _______ charges for all elements compounds the sum of the _____ 2) In ______, Add up to zero must __________. Ex: Na. Cl Ex: Mg 3 N 2 Ex: HNO 3 Na: 1(+1) = +1 Mg: 3(+2) = +6 H: Cl: 1(-1) = -1 N: 2(-3) = -6 N: 0! O: 0!

Oxidation number * The ____________ is the number ______ the ________. It is the

Oxidation number * The ____________ is the number ______ the ________. It is the charge on inside parentheses _____ one 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 (+) element first (-) element last _________ and the __________ in a compound formula. EXAMPLE: HCl EXCEPTION to this rule: NH 3

+1 oxidation 3) __________ always have a ___ Group 1 Metals number when in

+1 oxidation 3) __________ always have a ___ Group 1 Metals number when in a compound (bonded to another species). __________ always have a ___ oxidation +2 Group 2 Metals number when located within a compound. (Assign) Ox #: Ex: Li. Cl Mg. Cl 2 Fluorine is always a ___ -1 in compounds. The other 4) __________ -1 as long as they Halogens (ex: Cl, Br, I) are also ___ are the most electronegative element in the compound. (Assign) Ox #: Ex: HF Ca. Cl 2 Na. Br

Hydrogen +1 in compounds unless it is 5) _______ is a ___ combined with

Hydrogen +1 in compounds unless it is 5) _______ is a ___ combined with _____ in Group 1 or _________, Group 2 Metal which case it is ___. -1 (Assign) Ox #: Ex: HCl Li. H Oxygen Usually -2 6) _______ is ________ in compounds. (Assign) Ox #: Ex: H 2 O

Fluorine F which is more When combined with ______ (__), electronegative, __________. Oxygen is

Fluorine F which is more When combined with ______ (__), electronegative, __________. Oxygen is +2 (Assign) Ox #: Ex: OF 2 Peroxide, oxygen is -1 When in a _____________. A peroxide is X 2 O 2 a compound that has a formula of ____. (Assign) Ox #: Ex: Na 2 O 2 H 2 O 2

7) The sum of the oxidation numbers in polyatomic ions must equal the ______________

7) The sum of the oxidation numbers in polyatomic ions must equal the ______________ Charge on the ion (_________). See table E Ex: Cr 2 O 72 - a Cr: 2( +6 ) = +12 O: 7( -2 ) = -14 = -2 (charge on ion)

Oxidation numbers change for A reaction is REDOX if…______________ two elements within a reaction

Oxidation numbers change for A reaction is REDOX if…______________ two elements within a reaction _________________________ Gain of electrons Reduction (GER) = ____________ by an atom Oxidation number or ion; _____________ goes _______/_______ reduces down Loss of electrons Oxidation (LEO) = ____________ by an atom or ion; _____________ goes Oxidation number oxidizes _______/_______ up

Assign oxidation numbers for all elements and complete the tables: Example 1: C +

Assign oxidation numbers for all elements and complete the tables: Example 1: C + H 2 O Charge: Increases/Decreases increases decreases C 0 H+1 Example 2: Mn+4 CO + e-: Lost/Gained lost gained H 2 Oxidized/Reduced oxidized reduced Mn. O 2 + 4 HCl → Mn. Cl 2 + 2 H 2 O Charge: Increases/Decreases Cl-1 → increases decreases e-: Lost/Gained lost gained Oxidized/Reduced oxidized reduced

HALF REACTIONS Half reactions allow us to show the ________ Exchange of e- in

HALF REACTIONS Half reactions allow us to show the ________ Exchange of e- in a redox rxn. Half rxns. For each redox reaction, we can illustrate two _____. One half-reaction shows ______ oxidation and other shows ______. reduction Example of a Reduction Half Reaction: Fe 3+ + 3 e- → Fe gained *Electrons on left hand side, _____ in the rxn (____). GER Notice also how the charge for Fe goes down from left to right, ______ (____). reduction GER Charge goes down because Fe gained e-. _____

Example of an Oxidation Half Reaction: F e → Fe 3+ + 3 elost

Example of an Oxidation Half Reaction: F e → Fe 3+ + 3 elost LEO Notice *Electrons on left hand side, _____ in the rxn (____). also how the charge for Fe goes up from left to right, _______ (____). e -. oxidation LEO Charge goes up because Fe _______ lost NOTICE: Always _________ to the side of rxn that has Add electrons a _________ Higher total charge (remember: electrons are ______!) negative FOLLOWING THE LAW OF CONSERVATION: ü Half reactions follow the _______________. Law of conservation of mass This means that there must be the ____________ Same number of atoms on both sides of the reaction arrow. ü There must also be a _______________. In Conservation of charge half reactions, the __________________ Net charge must be the same on both sides of the equation, although it doesn’t necessarily need to equal zero.

RULES FOR SETTING UP HALF REACTIONS 1) Assign oxidation numbers to all elements in

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 1 st (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.

Assign oxidation numbers to all elements or polyatomic ions. Label the brackets for reduction

Assign oxidation numbers to all elements or polyatomic ions. Label the brackets for reduction (red) or oxidation (ox). Ex. 1: Mg + Zn. Cl 2 → Mg. Cl 2 + Zn OXIDATION Half Reaction: Mg 0 → ______ Mg+2 + ______ 2 e. REDUCTION Half Reaction: Zn+2 + ______ → ______ 2 e. Zn 0

Ex. 2 (balance masses): Hg + I 2 → Hg. I OXIDATION Half Reaction

Ex. 2 (balance masses): Hg + I 2 → Hg. I OXIDATION Half Reaction (make sure to balance the masses ): 2 e 2 Hg+1 + ______ 2 Hg 0 → ______ REDUCTION Half Reaction: I 20 2 e 2 I______ + ______ → ______

Ex. 3 (balance charges): Cu + Ag. NO 3 → Cu(NO 3)2 + Ag

Ex. 3 (balance charges): Cu + Ag. NO 3 → Cu(NO 3)2 + Ag OXIDATION Half Reaction: ______ → Cu 0 ______ 2 e. Cu+2 + ______ REDUCTION Half Reaction: ______ → ______ Ag 0 1 e. Ag+1 + ______ Now, we need to balance the charges: +1 + ___e 0 2 2 x (Ag+1 + 1 e- → Ag 0) = ___Ag 2 2 - → ___Ag ___

Table J and Spontaneous Reactions higher more General Rule: elements ______ on Table J

Table J and Spontaneous Reactions higher more 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.

Complete the following equations by writing in the products formed or “no rxn” Ex

Complete the following equations by writing in the products formed or “no rxn” Ex 1: Zn + Pb. Cl 2 → Zn. Cl 2 + Pb 0 Ex 2: Zn + Ba. O → No rxn Ex 3: Ca + Cr. F 2 → Ca. F 2 + Cr 0 Ex 4: Mn + Ni. S → Mn. S + Ni 0 Ex 5: Fe + Mg. I 2 → No rxn Co + Pb. Cl 2 → Co. Cl 2 + Pb 0 Ex 6:

TWO TYPES of ELECTROCHEMICAL CELLS 1. Voltaic (similar to a battery) 2. Electrolytic (similar

TWO TYPES of ELECTROCHEMICAL CELLS 1. Voltaic (similar to a battery) 2. Electrolytic (similar to alternator in cars) SIMILARITIES BETWEEN THE TWO: Redox • Both involve ______ reactions; chemical reactions which involve the flow of _______ electrons current • Both involve the flow of _________, or _____, Electrical energy measured in _____ volts • Both have _______ (conductive surfaces where oxidation or 2 electrodes anode and the _____ reduction occurs); called the _____ cathode • ________ or ________ occurs in each half cell oxidation reduction

RED CAT ___________ Reduction ALWAYS occurs at the cathode ___________ AN OX ___________ Oxidation

RED CAT ___________ Reduction ALWAYS occurs at the cathode ___________ AN OX ___________ Oxidation ALWAYS occurs at the anode ___________ (_________) Ions gain e- (_________) Metal loses e- anode to the • Electrons flow through the _______ wire from the _____ cathode _____.

Voltaic Cells chemical • Cells that ___________ convert _______ energy spontaneously into ________ energy

Voltaic Cells chemical • Cells that ___________ convert _______ energy spontaneously into ________ energy or electric _______. electrical current • __________ batteries CATHODE Less active • The ___________ of the 2 metals (Table J) • ____________________ to it Spontaneously attracts electrons positive Voltaic cell • the _______ electrode in a ___________ reduction • electrode where ________ occurs (_______) RED CAT ANODE More active • The ___________ of the 2 metals (Table J) Spontaneously loses electrons • ____________________ to cathode • the _______ electrode in a ___________ negative Voltaic cell AN OX • electrode where ________ occurs (_______) oxidation

Example 1: Wet Cell Car batteries are a form of _______ Lead storage battery

Example 1: Wet Cell Car batteries are a form of _______ Lead storage battery • _______, • Consists of LEAD ANODE and LEAD OXIDE CATHODE • Both electrodes immersed in a ________ solution Sulfuric acid • Advantage: process is readily ______ reversible (by alternator) • Disadvantage: very ____, dangerous heavy somewhat _______ Example 2: Dry Cell • ____________ are the type of batteries in a portable Dry cell batteries 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 Flow of ions • prevents the _____________ Build-up of charge

Voltaic Cells (a. k. a Galvanic Cells) → 1. Use Table J to predict

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? ______________________________ Pb is negative electrode and Ag is positive electrode Ag is gaining electrons and therefore reduced 3. Explain your answer to #2. __________________________ which is the cathode and +, Pb is losing electrons and therefore ____________________________________________________________________________ oxidized which is the anode and -.

4. At which electrode or in which half-cell does reduction occur? _____ 2 5.

4. At which electrode or in which half-cell does reduction occur? _____ 2 5. At which electrode or in which half-cell does oxidation occur? _____ 1 2 6. Which electrode is the cathode? _____ 1 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

Electrolytic Cells nonspontaneous • Cells that use _________ Electrical energy to force a ________________________

Electrolytic Cells nonspontaneous • Cells that use _________ Electrical energy to force a ________________________ to occur. Chemical reaction electrolysis electroplating • This process is for _________ and __________ Example: _____________ (keeps the car battery Alternator in car replenished with energy)

Electrolysis Experiment Animation (w/ tutorial): http: //media. pearsoncmg. com/bc/bc_0 media_chem/chem_sim/electrolysis_fc 1_gm_11 -26 -12/main. html

Electrolysis Experiment Animation (w/ tutorial): http: //media. pearsoncmg. com/bc/bc_0 media_chem/chem_sim/electrolysis_fc 1_gm_11 -26 -12/main. html Standard Hydrogen Electrode (Zinc) http: //www. chem. iastate. edu/group/Greenbowe/sections/projectfolder/animations/SHEZn. V 7. html Standard Hydrogen Electrode (Copper) http: //www. chem. iastate. edu/group/Greenbowe/sections/projectfolder/animations/SHECu. html

Electrolytic Cells CATHODE • electrode where _______ are _____ electrons sent • the ________

Electrolytic Cells CATHODE • electrode where _______ are _____ electrons sent • the ________ electrode (opposite of voltaic cell) negative • electrode where ________ occurs (______) RED CAT reduction ANODE • electrode where _______ are ____________ Drawn away from electrons • the ________ electrode (opposite of voltaic cell) positive oxidation • electrode where ________ occurs (______) AN OX NOTICE: • There is __________. This is a forced chemical reaction. No salt bridge • You will always see a ________ Power source hooked up to an electrolytic cell which drives the __________ Forced rxn

Compare and contrast the two types of electrochemical cells: GALVANIC/VOLTAIC ELECTROLYTIC Flow of e(spontaneous

Compare and contrast the two types of electrochemical cells: GALVANIC/VOLTAIC ELECTROLYTIC Flow of e(spontaneous or forced) spontaneous forced (+) electrode cathode anode cathode Anode → cathode (-) electrode *Direction of eflow Reduction ½ cell Oxidation ½ cell Cathode anode *Direction of e- flow is either “Anode → Cathode” or “Cathode → Anode”

 Nonspontaneous Reactions on Electrolytic Cells Electroplating (see electroplating video)

Nonspontaneous Reactions on Electrolytic Cells Electroplating (see electroplating video)