Redox & Electrochemistry

You can’t have one… without the other! REDuction Oxidation= Redox Reactions that involve the transfer of electrons Reduction & Oxidation must occur simultaneously The total # of e- lost must be = to the total e- gained. • • • • +2 • • -2 • • • Ca + • O • Ca • O • • • • You can’t have one… without the other!

Oxidation Oxidation = lose e- Any chemical reaction in which an element loses electrons & the oxidation # increases (more positive) Oxidation # increases -4 -3 -2 -1 +1 +2 +3 +4 Oxidation = lose e-

Reduction Reduction = gain e- Any chemical change in which an element gains electrons & the oxidation # decreases (more negative) Oxidation # decreases -4 -3 -2 -1 +1 +2 +3 +4 Reduction = gain e-

Ger ! LEO says GER : Lose Electrons = Oxidation Sodium is oxidized Gain Electrons = Reduction Chlorine is reduced

Rules for Assigning Oxidation #’s: 1) Uncombined elements (not bonded with any other element) have an oxidation # of ZERO. Ex. H2 + Cl2 = 2HCl H2 = 0 Cl2 = 0

Rules for Assigning Oxidation #’s: 2) Sum of the oxidation #’s in a compound must equal ZERO. +2 -1 Ex. CaCl2 +2(1) - 1(2) +2 - 2 = 0 Be sure to multiply the oxidation # by the # of atoms indicated by the subscript.

Rules for Assigning Oxidation #’s: 3) All metals in Group 1( ) have and oxidation # of +1. Alkali metals 4) All metals in Group 2( ) have and oxidation # of +2. Alkali earth metals

Rules for Assigning Oxidation #’s: 5) Fluorine always has an oxidation # of –1. 6) Hydrogen has an oxidation # of +1 in all compounds except: in metal hydrides hydrogen is -1 (metal and hydrogen) Ex. +1 -1 +2 -1 LiH CaH2 +1(1) - 1(1) +2(1) - 1(2) +1 - 1 = 0 +2 - 2 = 0

Rules for Assigning Oxidation #’s: 7) In ionic compounds, monoatomic ions have oxidation #’s equal to their ionic charge. +1 -1 +1 -1 Ex. Ex. LiBr KBr

Rules for Assigning Oxidation #’s: 8) Oxygen has an oxidation # of –2 in all compounds except: +1 -1 in peroxides (H2O2) oxygen is –1. +1(2) - 1(2) +2 - 2 = 0 +2 -1 with Fluorine (OF2) oxygen is +2. +2(1) - 1(2) +2 - 2 = 0

Rules for Assigning Oxidation #’s: 9)The sum of the oxidation #’s of all the atoms must equal the charge of the ion. +6 -2 2- Ex. SO4 ? - 2(4) = -2 ? - 8 = -2 6 - 8 = -2 - 2 = -2 So sulfur is a +6.

Identifying Redox Reactions: 1) Inspect oxidation numbers from reactants to products 2) Oxidation numbers are located in the upper right hand corner & track the movement of electrons

Identifying Redox Reactions: 3) Single replacement reactions are always redox reactions ___ Zn + ___ HCl  4) Double replacement reactions are NEVER redox reactions ___ NaOH + ___ HCl 

Identifying Redox Reactions: 5) Changes in oxidation numbers oxidation Ex. +1 -1 2Na + Cl2 = 2NaCl reduction Na went from to : Na was +1 oxidized Cl went from to : Cl was -1 reduced

Half-Reactions: 1) Show the EXCHANGE OF ELECTRONS in a redox reaction. 2) 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. 3) Follow the LAW OF CONSERVATION OF CHARGE. In half reactions, the NET CHARGE must be the same on both sides of the reaction arrow.

Writing Half Rxns: Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s) oxidation Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s) reduction 1) Assign oxidation numbers 2) Draw brackets to identify oxidation & reduction 3) Write : Ox: Red:

Writing Half Rxns: Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s) oxidation Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s) reduction 4) Draw : Ox: Mg(s) Mg2+(aq) + 2e- 5) Copy what is on each side of arrow:

Writing Half Rxns: Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s) oxidation Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s) reduction 4) Draw : Red: 2Ag+(aq) + 2e- 2Ag(s) 5) Copy what is on each side of arrow:

Writing Half Rxns: Given: Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s) oxidation Mg(s) + 2Ag+(aq)  Mg2+(aq) + 2Ag(s) reduction 7) Figure out how many e- you have. 8) Balance charge. (Oxid. e- are lost and go on right side; Red. e- are gained and go on the left side) **Mass & charge are balanced. Lost e- leo (Oxid # ) Ox: Mg0(s)  Mg+2(aq) + 2e- Red: 2Ag+(aq) + 2e-  2Ag(s) (Oxid # ) ger gain e-

Example 1: Leo Ger Ox: Al0  Al+3 + 3e- Red: Fe+2 + 2e-  Fe0 oxidation __Al + __Fe+2 __Al+3 + __Fe reduction 1) Write ½ rxns. Lost e- Ox: Al0  Al+3 + 3e- Red: Fe+2 + 2e-  Fe0 gain e-

__Al + __Fe+2 __Al+3 + __Fe oxidation __Al + __Fe+2 __Al+3 + __Fe reduction 2) The e- lost MUST = the e- gained. Lost e- 2 [Al0  Al+3 + 3e-] Ox: Red: 3 [Fe+2 + 2e-  Fe0] gain e-

__Al + __Fe+2 __Al+3 + __Fe oxidation __Al + __Fe+2 __Al+3 + __Fe reduction 3) Distribute the coefficient. 2 [Al0  Al+3 + 3e-] Ox: Red: 3 [Fe+2 + 2e-  Fe0] [2Al0  2Al+3 + 6e-] [3Fe+2 + 6e-  3Fe0]

__Al + __Fe+2 __Al+3 + __Fe leo 2 3 2 3 __Al + __Fe+2 __Al+3 + __Fe ger 4) Enter the coefficients into original equation. [2Al0  2Al+3 + 6e-] [3Fe+2 + 6e-  3Fe0]

Table J and Spontaneous Reactions Shows how active each metal and nonmetal is. Elements at the top on Table J are more reactive than the elements below them and replace elements below them in compounds.

Yes Li K Li + KCl Cl + Li K - Is Li above K in Table J ? Example 1: Li + KCl Cl + Li K - Let’s look at the metals. - Is Li above K in Table J ? Yes Li K - is more active than . The Li will replace K in cmpds. This rxn is Spontaneous ! No energy needed.

Yes F2 Cl F2 + 2NaCl 2Na + F Cl2 - Is F2 is above Cl in Table J ? Example 2: F2 + 2NaCl 2Na + F Cl2 - Let’s look at the nonmetals. - Is F2 is above Cl in Table J ? Yes F2 Cl - is more active than . The F2 will replace Cl in cmpds. This rxn is Spontaneous ! No energy needed.

Examples 3 & 4: Ca + MgCO3 CO3 + Ca Mg F Cl2 + 2NaF + 2Na Cl2 This rxn is Spontaneous F Cl2 + 2NaF + 2Na Cl2 This rxn is Not Spontaneous

Trends in Oxidation & Reduction Active metals:   Lose electrons easily   Are easily oxidized Active nonmetals:   Gain electrons easily   Are easily reduced

Electrochemistry There are 2 types of electrochemical cells, voltaic (galvanic) and electrolytic.

Voltaic Cells (Galvanic) Cells that spontaneous convert chemical energy into electrical energy or electrical current Examples: wet cell batteries-lead storage dry Cell Batteries-remotes, radios

Voltaic Cells Electrodes: Cathode REDuction = CAThode Less active of the 2 metals (Table J) Spontaneously attracts electrons to it Positive electrode in a voltaic cell Reduction takes place at the cathode REDuction = CAThode

Voltaic Cells Electrodes: Anode ANode = OXidation More active of the 2 metals (Table J) Spontaneous loses electrons to cathode Negative electrode in a voltaic cell Oxidation takes place at the anode ANode = OXidation In a voltaic cell anode =

Description of Voltaic Cells Half Cells: - 2 half cells - rxns occur in separate vessels Half cell 1 Half cell 2

Description of Voltaic Cells Half Cells: - electrode – a metal strip - wire connects to electrodes - allows the flow of electrons wire electrode electrode Half cell 1 Half cell 2

Description of Voltaic Cells Half Cells: - voltmeter – measures electric current - salt bridge - allows the flow of ions - prevents mixing of 2 solutions switch wire voltmeter IONS electrode electrode Salt bridge Half cell 1 Half cell 2

Description of Voltaic Cells 1) 2 half cells Complete Circuit 2) Wire & electrodes 3) Salt bridge 4) voltmeter switch wire voltmeter IONS electrode electrode Salt bridge Half cell 1 Half cell 2

Anode to cathode Let’s see what happens in a voltaic cell: - A rxn between Zn & Cu - How will e- travel? Anode to cathode RED CAT e- e- e- e- AN OX e- e- ger leo Cu Zn red. oxid. CATHODE ANODE Zn(NO3)2 Cu(NO3)2

According to Table J, which metal is oxidized? Zn Which metal is reduced? Cu e- e- e- e- e- e- + + + CATHODE + ANODE + + + + + + + + + + + + Zn(NO3)2 Cu(NO3)2

Electrolytic Cells Cells that use ELECTRICAL ENERGY to force a NONSPONTANEOUS CHEMICAL REACTION to occur. There is NO SALT BRIDGE. You will always see a POWER SOURCE (battery) hooked up to an electrolytic cell which drives the FORCED REACTION

Electrolytic Cells CATHODE electrode where ELECTRONS are SENT the NEGATIVE electrode (opposite of voltaic cell) electrode where REDUCTION occurs (RED CAT) ANODE electrode where ELECTRONS are DRAWN AWAY FROM the POSITIVE electrode (opposite of voltaic cell) electrode where OXIDATION occurs (AN OX)

Electrolytic Cells Used for ELECTROLYSIS and ELECTROPLATING Examples: alternator (keeps the car battery replenished with energy) Used for ELECTROLYSIS and ELECTROPLATING

Electrolytic Cells Electrolysis: the decomposition of a substance with an electric current. Examples: water

Electrolytic Cell Electroplating-the process of adding a layer (plate) of metal on the surface of another object. Example: Gold plated jewelry & chrome bumpers

Electroplating 1) Oxidation occurs at the anode. 2) The metal (Ag) bar loses e-. e- e- e- e- 3) e- flow from the metal to the object (spoon) e- e- Cathode is always attached to the neg. terminal of battery. e- anode cathode

Electroplating 4) The spoon gains e- and becomes negative. e- e- e- e- e- 5) The spoon attracts the positive Ag+ ions in the solution. e- e- e- e- 6) The Ag+ ions stick to the negative spoon – plating it with silver. e- e- anode cathode

Electroplating 7) The Ag bar anode loses e- and will eventually disappear. e- e- e- e- e- e- e- At the anode: Ag     Ag+ + e- e- e- + e- At the cathode: Ag+ + e-     Ag + + e- anode + + cathode

Comparison of Voltaic and Electrolytic Cells Similarities Both use redox reactions The anode is the site of oxidation The cathode is the site of reduction The electron flow through the wire is from anode to cathode Differences Voltaic cells use spontaneous reactions to produce energy (voltmeter) Electrolytic use non-spontaneous reactions that requires energy (power source) Voltaic cells the anode is negative and the cathode is positive Electrolytic cells the anode is positive and the cathode is negative.