Oxidation- Reduction

Ms. Randall

Lesson 2: Recognizing Oxidation-Reduction

Reactions

Objective: To identify redox reactions based on the changes of oxidation states of atoms in a chemical reaction.

REDOX- REDuction – OXidation

• Reactions: reactions that involve the TRANSFER OF ELECTRONS; both reduction and oxidation must happen SIMULTANEOUSLY!

Reduction = GAIN OF ELECTRONS by an atom or ion

-OXIDATION NUMBER goes DOWN/REDUCES

-Nonmetals GAIN electrons 

Oxidation= LOSS OF ELECTRONS by an atom or ion

- OXIDATION NUMBER goes UP/INCREASES

-Metals LOSE electrons

2 ways to remember L E O the lion goes G E R

Lose e- oxidationgain e- reduction

  O I L R I G Oxidation is losing e-

reduction is gaining e-

*Oxidation and reduction happen because of the DESIRE for electrons in a chemical reaction. Species prefer to either LOSE or GAIN electrons in a chemical reaction.

 AND **Oxidation and reduction are

SIMULTANEOUS reactions and one cannot happen without the other. If one atom LOSES electrons, there must be another atom that will GAIN electrons.

Demo: Let’s make some copper!!

 Example: Al + CuCl2 Cu + AlCl3

 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

 NOTE: NOT ALL REACTIONS are REDOX reactions! We will now investigate how we can identify a redox reaction from other types of reaction.

IDENTIFYING REDOX REACTIONS

•  One way that we can begin to identify a redox reaction is to inspect the OXIDATION NUMBERS of each atom/ion from reactant to product side.

• Oxidation numbers are used to TRACK THE MOVEMENT OF ELECTRONS for each atom/ion (ELECTRON TRANSFER) FROM REACTANT TO PRODUCT SIDE OF RXN.

Oxidation Number (State) =

• POSITIVE, NEGATIVE, OR NEUTRAL (ZERO) VALUES that can be assigned to atoms; identify how many electrons are being lost or gained by an atom/ion WHEN THEY BOND

*top listed # to the upper right is the “most common” oxidation number for that element

Trick 1: SINGLE REPLACEMENT

REACTIONS are always REDOX!

Oxidation # 0 +1 -1 0 +2 -1

Example: Zn(s) + HCl(aq) H2(g) + ZnCl2 (aq)

 • *Zn/H2 are by themselves on one side and

bonded on the opposite side!

Trick 2: DOUBLE REPLACEMENT REACTIONS are NOT REDOX!

*charges stay the same for all elements in the rxn

Rules for assigning OXIDATION STATES (numbers):

1. UNCOMBINED ELEMENTS (ELEMENTS NOT BONDED TO ANY OTHER TYPE OF ELEMENT) have an oxidation number of ZERO.

• This includes diatomic elements.

Examples: Al(s)0 Na(s)0 Cl20

H20

2. The sum of the CHARGES in COMPOUNDS must ADD UP TO ZERO

Rules continued…3. Group 1 are always +1 (periodic

table) 4. Group 2 are always +2 (periodic

table) 5. F is always -1 (periodic table) 6. O is -2 (periodic table) except: -1 in

H2O2, and +2 in OF2 (fluorine is more electronegative)

7. H is +1 (periodic table) except: when combined with metals

8. Sum of the oxidation # in a polyatomic ion (table E) must equal the charge of the ion listed

 Ex: Cr2O7

2- Cr: 2(+6) = +12

O: 7(-2) = -14 -2

A reaction is REDOX if• OXIDATION NUMBERS CHANGE FOR 2

ELEMENTS WITHIN A REACTION• Oxidation = occurs where OXIDATION

NUMBER INCREASES for one specific element from reactant to product side of rxn; LEO or OIL

• Reduction = occurs where OXIDATION NUMBER DECREASES for one specific element from reactant to product side of rxn; GER or RIG

Check your understanding and practice

Lesson 3: Half reactions

Objective:To create half reactions and use them to balance redox equations

HALF REACTIONS

• Show the EXCHANGE OF ELECTRONS in a Reaction

• For each redox reaction, we can illustrate two half reactions. One shows OXIDATION and other illustrates REDUCTION.

• NOTICE: Always ADD ELECTRONS to the side of reaction that has a HIGHER TOTAL CHARGE (remember: electrons are NEGATIVE!)

FOLLOWING THE LAW OF CONSERVATION:

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.

 There must also be a CONSERVATION OF CHARGE. In 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

2. Assign oxidation numbers to all elements in reaction

1. Determine if a reaction is redox (look for change in oxidation #of 2 elements.

3. Determine which species is oxidized and which is reduced (use brackets).

4. Then break the overall reaction into oxidation and reduction reactions called HALF REACTIONS. ***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. 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 the 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.

Balancing Redox Reactions

Now, we need to balance the charges: ***(Remember that the same # of moles of electrons that are lost in oxidation must be gained in reduction)

  Multiply each half reaction so the electron #

is equal in both reactions. (Remember electrons lost equals electrons gained)

DiatomicsExample: Balance the following equation

H2 + O2 H20

Not balanced:H2

0 H+1 + 1e-

O20 + 1e- O-2

Balanced for mass in the half reaction, then check to see if the electrons lost = electrons gained.

H20 2H+1 + 2e-

O20 + 4e- 2O-2

Then you would multiply the appropriate half reaction to balance electrons lost= electrons gained. Now balance the overall reaction for mass

2(H20 2H+1 + 2e-)

O20 + 4e- 2O-2

2H2 + O2 2 H20

Check your understanding and practice

Lesson 4: Spontaneous Reactions & Reference Table J

Objective:Use table J to determine if a reaction occurs spontaneously.

Table J and Spontaneous Reactions (TELLS OF IF A REDOX REACTION WILL TAKE PLACE

• General Rule: elements HIGHER on Table J are MORE 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.

 Ex 1: Zn + PbCl2 ZnCl2 +

Pb    Ex 2: Zn + AlCl3 No rxn

Check your understanding and practice

Lesson 5: Electrochemical cells

Objective:To identify the oxidation and reduction reactions occurring in an electrochemical and voltaic cells.

TWO TYPES of ELECTROCHEMICAL CELLS

1. Voltaic (similar to a battery)2. Electrolytic (similar to alternator in

cars)

Voltaic Cells (a.k.a Galvanic Cells)

• VOLTAIC CELLS SPONTANEOUSLY (no extra energy needed) convert CHEMICAL energy into ELECTRICAL. Basically, a voltaic cell is a BATTERY.

•  • Electrons flow through the WIRE from

the ANODE to the CATHODE.

ANODE

• the NEGATIVE electrode in a VOLTAIC CELL

• SPONTANEOUSLY LOSES ELECTRONS to cathode

• electrode where OXIDATION occurs

CATHODE

• the POSITIVE electrode in a VOLTAIC CELL

• SPONTANEOUSLY ATTRACTS ELECTRONS to it

• electrode where REDUCTION occurs

*SALT BRIDGE

• connects the two half cells and provides a path for the flow of IONS between the two containers (this is only part of what completes the circuit)

Galvanic Cell Animation

Practice

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? Cu is +, Pb is –

3. Explain your answer to #2. If e- spontaneously flow to Cu, then it has to be positive COMPARED to Pb. If Pb supplies e-, then it must be negative COMPARED to Ag.

7. At which electrode or in which half-cell does reduction occur? Cu

8. At which electrode or in which half-cell does oxidation occur? Pb

9. Which electrode is the cathode? Cu10. Which electrode is the anode? Pb*Electrons don’t flow to the cathode;

they flow through it to the ions in solution. That’s why the cathode never becomes negative.

• http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

Electrolytic Cells

• Cells that use ELECTRICAL ENERGYto force a NONSPONTANEOUS CHEMICAL REACTION to occur.

• This process is also known as ELECTROLYSIS.

CATHODE

• the NEGATIVE electrode (opposite of voltaic cell)

• electrode where ELECTRONS are SENT

• electrode where REDUCTION occurs (RED CAT)

ANODE

• the POSITIVE electrode (opposite of voltaic cell)

• electrode where ELECTRONS are DRAWN AWAY FROM

• electrode where OXIDATION occurs (AN OX)

NOTICE:

• There is NO SALT BRIDGE. This is a forced chemical reaction.

• You will always see a POWER SOURCE hooked up to an electrolytic cell which drives the FORCED RXN

WHAT YOU MUST KNOW ABOUT THE TWO CELLS

 

Check your understanding and practice