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Ms. Reynolds

MacArthur High School: (516) 434-7225

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REDOX
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​​​​​​Unit 9: Oxidation-Reduction Reactions (Redox)​


Monday

Tuesday

Wednesday

Thursday

Friday

4/20

Lesson 1:Assigning Oxidation Numbers

1. Watch Power Point video

2. Copy notes

3. Q1 Review due – complete on TestWizard.com

4/21

Lesson 1:Assigning Oxidation Numbers

1. Read pages 158-160

2. Complete questions 1-10 submit answers on Edmodo (picture of text or list of answers)

4/22

Lesson 2:RedOx Reactions

1. Watch Power Point video

2. Copy notes

4/23

Lesson 2:RedOx Reactions

1. Read pages 161-162

2. Complete questions 11-12, 17-20, 22-24 submit answers on Edmodo (picture of text or list of answers)

 

4/24

Lesson 3:Half-Reactions

1. Watch Power Point video

2. Copy notes

 

4/27

Lesson 3:Half-Reactions

1. Read pages 163-164

2. Complete questions 32-36 submit answers on Edmodo (picture of text or list of answers)

4/28

Lesson 4:Electrochemical Cells – Terminology

1. Copy notes

4/29

Lesson 5:Electrochemical Cells – Spontaneous Reactions

1.  Watch Power Point video

2.  Copy notes

4/30

Lesson 5:Electrochemical Cells – Spontaneous Reactions

1.  Read pages 165-166

2.  Complete questions 37-38, 40, 42-44 submit answers on Edmodo (picture of text or list of answers)

 

5/1

Lesson 6:Electrochemical Cells – Nonspontaneous Reactions

1.  Watch Power Point video

2.  Copy notes

5/4

Lesson 6:Electrochemical Cells – Nonspontaneous Reactions

1.  Read pages 167-168

2.  Complete questions 39, 41, 45-46 submit answers on Edmodo (picture of text or list of answers)

 


 

 

 

 

 

























Learning Targets: At the end of this unit you should be able to do the following:

1.      I can assign oxidation numbers to elements alone or in compounds/polyatomic ions.

2.      I can define oxidation as the loss of electrons and reduction as the gain of electrons in a system.

3.      I can comprehend that oxidation and reduction MUST go TOGETHER.

4.      I can identify redox reactions as those that show a change in oxidation states of the substances.

5.      I can determine the number of electrons being lost or gained because of CONSERVATION.

6.      I can write and balance half reactions.

7.      I can differentiate between spontaneous reactions and nonspontaneous reactions.

8.      I can draw and label a voltaic cell to represent a spontaneous reaction in which chemical energy is converted to electrical energy.

9.      I can draw and label an electrolytic cell to represent a nonspontaneous  reaction in which electrical energy is converted to chemical energy.

10.  I can determine the flow of electrons which is ALWAYS from anode to cathode!

11.  I can explain that the purpose of the salt bridge  in a voltaic cell is for the flow of ions in a spontaneous reaction.

12.   I can explain that a battery is needed in an electrolytic cell because it is a nonspontaneous reaction.​





Monday 4/20/2020 Assignment: 1.  Watch Power Point video of lesson.  2.  Copy the notes   3. Q1 Review due on www.TestWizard.com

Tuesday 4/21/2020 Assignment: 1. Read pages 158-160; 2. Complete questions 1-10 and submit a picture of your answers on Edmodo

Power Point video of notes and examples. Please play slideshow from beginning: Lesson 1.pptxLesson 1.pptx  


Lesson 1: Assigning Oxidation Numbers 

Oxidation Numbers:  are assigned to atoms or ions as a way to keep track of electron transfers.

 Rules for Assigning Oxidation Numbers (top right hand corner of periodic table element)

Free standing elements (elements alone) = 0 (ex: H2, Na, S8)
Ions (in ionic compounds) carry their charge (group 1 elements =+1, group 17 elements = -1…)
Hydrogen = +1 unless it is with a metal (metal hydride) then -1
Oxygen = -2 in most compounds

Compounds total 0 (or charge if a polyatomic ion)


Finding Oxidation Numbers

Situation 1:  Find the oxidation numbers of N and O in N2O.

Solution:  

*Oxygen has an oxidation number of –2

*The total for compounds must = 0

           so Nitrogen must total +2, since there are 2 of them each N has a +1 oxidation number.

 

Situation 2:  Find the oxidation numbers of K, Mn and O in KM​​nO4.

Solution:  For compounds composed of more than two elements, draw a chart.

Element

K

Mn

O4

Each

+1

?

-2

Total

+1

+7

-8

*Oxygen has an oxidation number of –2

*Potassium is a group 1 element with an oxidation number of +1

*The total for compounds must = 0​

          so Manganese must be +7

 



Wednesday 4/22/2020 Assignment: 1.  Watch Power Point video of lesson    2.  Copy the notes     *You can also watch this quick video for additional review https://www.youtube.com/watch?v=W4-LMBx9AFs     

Thursday 4/23/2020 Assignment: 1. Read pages 161-162; 2. Complete questions 11-12, 17-20, 22-24 and submit a picture of your answers on Edmodo

Power Point video of notes and examples. Please play slideshow from beginning: Lesson 2 RedOx Reactions.pptxLesson 2 RedOx Reactions.pptx



Lesson 2: RedOx Reactions  

Redox Reactions: reactions in which both oxidation and reduction occur.

A.   Oxidation: the loss of electrons by an atom or an ion, increasing the oxidation number.

B.   Reduction: the gain of electrons by an atom or an ion, reducing the oxidation number.

*During Redox there must be an exchange of electrons; one substance must lose electrons (be oxidized) and one substance must gain electrons (be reduced)*

  ​        Lose Electrons Oxidation                               

       Gain Electrons Reduction 

                                                                                               redox leoger.jpg

 

*Sometimes you will see the acronym OIL RIG which means Oxidation ILosing electrons Reduction IGaining electrons

 






Not all reactions are RedOx reactions. In order to decide if a reaction is a RedOx, you must:

         1. Assign oxidation numbers using the rules listed in lesson 1

         2. If a substance has a change in oxidation number then it is RedOx

 

  Quick identification of RedOx reactions:

         1. If an uncombined element is on one side of the equation and in a compound on the other side it is RedOx

         2. Single Replacement reactions are ALWAYS RedOx

         3. Double Replacement reactions are NEVER RedOx


Examples: 

Single Replacement:   Li + NaCl --> Na + LiCl         

                                Li goes from 0 (it's alone) to +1 (in the compound) 

                                Na goes from +1 (in the compound) to 0 (it's alone)

                                Cl is -1 on both sides (called a spectator ion)


Double Replacement:  MgCl​2  +  CaBr2  --> MgBr2  + CaCl2

                                 Mg is +2 on both sides because it is in a compound and it is a group 2 element

                                 Ca is +2 on both sides because it is in a compound and it is a group 2 element

                                 Cl is -1 on both sides because it is in a compound and it is a group 17 element

                                 Br is -1 on both sides because it is in a compound and it is a group 17 element

There is no change in oxidation numbers.

Spectator Ions: are not changed during a RedOx reaction



                                
Friday 4/24/2020 Assignment: 1.  Watch Power Point video of lesson.  2.  Copy the notes   

Monday 4/27/2020 Assignment: 1. Read pages 163-164; 2. Complete questions 32-36 and submit a picture of your answers on Edmodo

Power Point video of notes and examples. Please play slideshow from beginning:​ Lesson 3 Half Reactions.pptxLesson 3 Half Reactions.pptx


​Lesson 3: Half Reactions

A.   Electronic equations or equations in which only the atoms/ions being oxidized or reduced are shown.

Let's look at a complete reaction from Lesson 2

Li + NaCl --> Na + LiCl

       Li goes from 0 to +1  (How did it do this? It lost 1 negatively charged electron)

       Na goes from +1 to 0   (How did it do this? It gained 1 negatively charged electron)

       Cl is a Spectator Ion: ions that are not changed during a redox reaction.


B.   Half Reactions break the electronic equation into an oxidation half and a reduction half.

                 Oxidation ½: Li0 → Li+1 + 1e-

Reduction ½: Na+1 + 1e- → Na0

 

C.   Balancing Half Reactions

   Must conserve the number of electrons! Number of electrons lost equals number of electrons gained

   We will not always find RedOx reactions in which the oxidized and reduced substances lose/gain the same number of electrons - we will need to balance them.  

  Example:  AlCl3 + Ba --> Al + BaCl2

                  Al goes from +3 (in a compound) to 0 (alone)    How did it do that? It gained 3 electrons

                  Ba goes from 0 (alone) to +2 (in a compound)   How did it do that? It lost 2 electrons

Half Reactions:    Ox 1/2: Ba0 --> Ba+2 + 2e-         Charges:   0 = +2 + -2

                        Red 1/2: Al+3 + 3e- --> Al0           Charges:   +3 +-3 = 0  

1.Notice that the element with the 0 oxidation state stays alone, the electrons go on the side of the ion.

2. Notice that if you add the charges on both sides of the equation they will equal each other  

If each Ba loses 2 electrons but each Al gains 3 electrons what is the least common multiple of 2 and 3?  ​6

So multiply the top reaction by 3      3Ba0 --> 3Ba+2 + (3)2e-

and the bottom reaction by 2           2Al+3 + (2)3e- --> 2Al0 

When you balance the initial reaction you get:

                                                     ​2AlCl3 + ​3Ba + 6e- --> ​2Al + 3BaCl2 + 6e- ​

Mass, energy & charge are ALWAYS conserved!

We just don't usually "see" the electrons.

                                                                             ​2AlCl3 + ​3Ba --> ​2Al + 3BaCl2​



Tuesday 4/28/2020 Assignment: 1.  Copy the notes (definitions)   


Lesson 4: Electrochemical Cells

Terminology:

1. Half cell: produced when a metal is placed in a salt solution. Provides metals to be oxidized and metal ions   to be reduced.   

2. Anion: a negatively charged ion (because it gained electrons)    

3. Cation: a positively charged ion (because it lost electrons)   

4. Electrode: the site at which oxidation or reduction occurs (usually a metal bar)   

5. Anode: the electrode where oxidation occurs    

6. Cathode: the electrode where reduction occurs    

7. Electrochemical cell: a cell which converts either electrical energy to chemical energy OR chemical energy to electrical energy

8. Voltaic Cell: a spontaneous electrochemical cell that converts chemical energy to electrical energy 

9. Electrolytic Cell: a nonspontaneous electrochemical cell that requires electrical energy to produce a chemical reaction.    





 Wednesday 4/29/2020 Assignment: 1.  Watch Power Point video of lesson.  2.  Copy the notes   

Thursday 4/30/2020 Assignment: 1. Read pages 166-166; 2. Complete questions 37-38, 40, 42-44 and submit a picture of your answers on Edmodo

Power Point video of notes and examples. Please play slideshow from beginning:​Lesson 5 Spontaneous Rxns.pptxLesson 5 Spontaneous Rxns.pptx


Lesson 5: Spontaneous Reactions and Voltaic Cells

​(REVIEW) Table J - Reactivity Series of Selected Metals & Nonmetals 

The element that is higher on table J can replace the other element spontaneously. ​

Spontaneous Reaction: a reaction that proceeds to completion without adding energy sources once it begins.

1.    You must have something to oxidize and something to reduce.

*One metal must be neutral, the other an ion!!!

              

2.     Metals: 

        The metal being oxidized (the neutral one) must be higher on Table J than the ion being reduced.

2Cr+3 + 3Mn(s) = 3Mn+2​ + 2Cr(s) 

*Mn is neutral and it is higher on Table J so it will be oxidized​

Nonmetals: 

The nonmetal being reduced (the neutral one) must be higher on Table J than the ion being oxidized.

Cl2 + 2Br- --> 2Cl- + Br2

*Cl  is neutral and it is higher on Table J so it will be reduced.


Voltaic Cells (aka Galvanic Cells): Half cells of 2 different metals connected by wires to produce electricity. 

· These are batteries which redox is spontaneous.

· Electrons flow from the metal being oxidized(one losing electrons) to the ion being reduced (one gaining electrons).

· Salt Bridge:  functions to complete the circuit by allowing ions to flow from reducing side to oxidizing side.

In the diagram below we see Zinc & Copper.                                                             

Since Zinc is higher on Table J it will be oxidized (lose electrons)

Electrons ALWAYS FLOW FROM ANODE TO CATHODE (AnOx fat RedCat)

redcatanox.jpg 

The half reactions will be:

                 Ox: Zn0 --> Zn+2 + 2e-

               Red: Cu+2 + 2e- --> Cu      

                                                                            

                   voltaic cell.jpg







PLEASE DRAW THIS DIAGRAM






 
Friday 5/1/2020 Assignment: 1.  Watch Power Point video of lesson.  2.  Copy the notes   

Monday 5/4/2020 Assignment: 1. Read pages 167-168; 2. Complete questions 39, 41, 45-46 and submit a picture of your answers on Edmodo

Power Point video of notes and examples. Please play slideshow from beginning:​​Lesson 6 nonspontaneous reactions.pptxLesson 6 nonspontaneous reactions.pptx


Lesson 6: Nonspontaneous Reactions and Electrolytic Cells  

Electrolytic Cells: electricity (a power source) is used to produce a nonspontaneous reaction

· Electricity is required!  This is not spontaneous!!  Requires the use of a Battery.

- Uses: Electroplating, electrolysis

*The only way to get a group one metal in its elemental form is by electrolysis of their fused salts.

Example:  KBr = K(s) + Br2(s)                        

 electrolytic cell.jpg
 

​Feature ​Voltaic ​Electrolytic
​Reation ​Type ​Spontaneous ​Nonspontaneous
​Oxidation ​Anode ​Anode
​Reduction ​Cathode ​Cathode
​Electron Flow ​Anode to Cathode ​Anode to Cathode
​Anode Charge ​Negative ​Positive
​Cathode Charge ​Positive ​Negative
​Uses  ​Batteries ​Electroplating & electrolysis