Unit 9: Oxidation, Reduction, and Electrochemistry

TOPIC 1: Half Reactions Objective: Determine oxidation numbers for elements in a compound, identify the oxidized and reduced species, write half reaction, and balance ionic reactions

I. What is Electrochemistry? The study of electricity and how it relates to chemistry Electricity is caused by the movement of _____________ We will be focusing on how energy is produced/converted by electrochemical cells: Voltaic -_______________________________ Electrolytic - __________________________ charges (e-) Produces electrical energy (ie. – battery) Uses electrical energy to produce a chemical reaction (ie. – electroplating)

I. What is Electrochemistry? chemical reactions Within these cells there are _____________________ that involves a ____________________ These chemical reactions are called _________________ Reduction: _________________ and the ion becomes ______________________   Oxidation: ______________________ and the ion becomes ______________________ transfer of electrons redox reactions gain of e- more negatively charged loss of e- more positively charged

I. What is Electrochemistry? Way to remember it: “LEO says GER” LOSS of Electrons is OXIDATION GAIN of Electrons is REDUCTION

II. Assigning Oxidation Numbers The first step in determining which species was oxidized and which was reduced is knowing what the charge of each species starts at and ends at, so you can determine whether electrons were lost or gained. Why is that important? Ever put a battery in backwards? It doesn’t work! Circuits are designed to only let electrons through ONE way. It is possible to destroy an electronic device by putting the battery in backwards. We need to know which end is – (the end that loses electrons, oxidation) and which end is + (the end that gains electrons, reduction), so we need to know how to find the charges (oxidation numbers) of the species involved!

II. Assigning Oxidation Numbers Rules for assigning oxidation number (REVIEW): 1. Each uncombined element has an oxidation number of _________ Example: 2Na + Cl2 2NaCl __________________ has an oxidation number (charge) of _______ 2. If an element has only ______ oxidation number (charge) listed on the ___________________, then that is its oxidation number  Examples: Na has a charge of __________ Ca has a charge of __________ Al has a charge of __________ O has a charge of __________ zero Na and Cl2 zero one periodic table + 1 + 2 + 3 - 2

II. Assigning Oxidation Numbers Rules for assigning oxidation number (REVIEW): 3. If a nonmetal atom is the negative ion in an ionic compound, then the______________________ listed is its oxidation number Examples: In NaCl, Cl is the negative ion, so its charge is _____ In Li3N, N is the negative ion, so its charge is ______ In ZnS, S is the negative ion, so its charge is _____ first charge - 1 - 3 - 2

II. Assigning Oxidation Numbers Rules for assigning oxidation number (REVIEW): 4. The sum of all charges in an compound equals _____ Examples: Show that the sum of the charges equals zero for the following compounds: NaCl Li3N MgCl2 ZERO 1(+1) + 1(-1) = 0 3(+1) + 1(-3) = 0 1(+2) + 2(-1) = 0

II. Assigning Oxidation Numbers Rules for assigning oxidation number (REVIEW): 5. The sum of all charges in a polyatomic ion equals ___________________________ Examples: Identify the charges for the following PAI: SO4 NO3 S2O3 SCN the net charge listed on Table E - 2 - 1 - 1 - 2

II. Assigning Oxidation Numbers Rules for assigning oxidation number (REVIEW): 6. If an element has more than one charge listed, use the other charges to figure it out   ZnSO4 Fe(NO3)2 PO4-3 (+2) + 2(X) + 6(-2) = 0 X + 4(-2) = -3 1(+2) + X + 4(-2) = 0 2 + X – 8 = 0 2 + X – 12 = 0 P = X = +5 S = X = +6 N = X = + 5

III. Identifying the Oxidized and Reduced SPECIES If the charge becomes more positive going from left to right, the species is __________ and _______________ If the charge becomes more negative going from left to right, the species is _________ and _________________ Spectator Ions _________________________________ OXIDIZED LOSES ELECTRONS (OX # ↑) REDUCED GAINS ELECTRONS (OX # ↓) ION THAT DOES NOT CHANGE CHARGE

III. Identifying the Oxidized and Reduced SPECIES For the following reactions, identify the species (symbol AND charge) that undergoes oxidation, reduction or is the spectator ion (if available):  2Na + Cl2  2 NaCl Oxidized ______ Reduced _________ Spectator Ion ______ +1 -1 Nao Cl2o NONE

III. Identifying the Oxidized and Reduced SPECIES +2 -2 2CaO  2 Ca + O2   Oxidized _______ Reduced ________ Spectator Ion _______ 3 Mg + 2 Au(NO3)3  3 Mg(NO3)2 + 2 Au Oxidized _______ Reduced _________ Spectator Ion ________ O-2 Ca+2 NONE +3 -1 +2 -1 Mg0 NO3- Au+3

IV. Identifying the Oxidized and Reduced AGENTS reduction Reducing Agent: causes ___________ (or is _____________)   Oxidizing Agent: causes ___________ (or is _____________) For the following reactions, identify the oxidizing and reducing agent and spectator ion (if possible) 2 Li + F2  2 LiF Oxidizing Agent _____ Reducing Agent _____ Spectator Ion ______ oxidized oxidation reduced **AGENTS ALWAYS COME FROM THE REACTANT SIDE** +1 -1 ox. red. F20 Li0 NONE

IV. Identifying the Oxidized and Reduced AGENTS 2 MgO  2 Mg + O2 Oxidizing Agent_______ Reducing Agent _______ Spectator Ion ______ Ba + Sr(OH)2  Ba(OH)2 + Sr Oxidizing Agent_______ Reducing Agent _______ Spectator Ion ______ +2 -2 red. ox. Mg+2 O-2 NONE +2 -1 +2 -1 ox. red. Sr+2 Ba0 OH-1

V. Identifying Redox Reactions (Most) Redox reactions have a free element on one side of a reaction and the same element bonded to something else on the other side Examples: 2H2 + O2 --> 2H2O             Free element ---> Bonded element 2KClO3 --> 2KCl +3O2       Bonded oxygen ----> Free oxygen

CHECK FOR UNDERSTANDING Which equation represents an oxidation reduction reaction? (1) CH4 + 2O2 ==> CO2 + 2H2O (2) H2SO4 + Ca(OH)2  ==> CaSO4 + 2H2O (3) MgCrO4 + BaCl2  ==> MgCl2 + BaCrO4 (4) Zn(NO3)2 + Na2CO3  ==> 2NaNO3 + ZnCO3 Which balanced equation represents a redox reaction? (1) AgNO3 + NaCl ==>AgCl + NaNO3 (2) BaCl2 + K2CO3 ==> BaCO3 + 2KCl (3) CuO + CO ==> Cu + CO2 (4) HCl + KOH ==> KCl + H2O Conclusion Redox reactions can be combustion, synthesis, decomposition, and single replacement reactions, but NOT double replacements) If there is a change in oxidation # for a particular atom, then it is a redox reaction.

VI. Half-Reactions oxidation reduction product lost reactant gained A redox reaction may be split into 2 half-reactions, one for _______________ and the other for ____________________   Oxidation: X0  X+ + e- X- X0 + e- Reduction: X+ + e-  X0 X0 + e - X oxidation reduction (e- on the _____________ side shows that it has been _______) product lost reactant (e- on the ____________ side shows that it has been_______) gained IN GENERAL: e- go on the more positive side

VI. Half-Reactions +2 -2 red ox S -2  S0 + 2 e- Ca +2 + 2 e-  Ca0 For the following reactions, write the half-reactions:   CaS  Ca + S Oxidation Reduction +2 -2 red ox S -2  S0 + 2 e- Ca +2 + 2 e-  Ca0

VI. Half-Reactions +2 -1 +2 -1 ox red Mg 0  Mg+2 + 2 e- For the following reactions, write the half-reactions:   Mg + CuSO4  MgSO4 + Cu Oxidation Reduction +2 -1 +2 -1 ox red Mg 0  Mg+2 + 2 e- Cu +2 + 2 e-  Cu0

VI. Half-Reactions +1 -2 ox red H2 0  2H+1 + 2 e- O2 0 + 4 e-  2O -2 For the following reactions, write the half-reactions:   2H2 + O2  2 H2O Oxidation Reduction +1 -2 ox red H2 0  2H+1 + 2 e- O2 0 + 4 e-  2O -2

VI. Half-Reactions +1 -1 +1 -1 red ox Na 0  Na+1 + 1 e- For the following reactions, write the half-reactions:   2 Na + 2HOH  2 NaOH + H2 Oxidation Reduction +1 -1 +1 -1 red ox Na 0  Na+1 + 1 e- 2 H + + 2 e-  H20

VII. Balancing Ionic Redox Reactions A chemical reaction MUST HAVE: 1. Conservation of _________________ 2. Conservation of _________________ Ionic reaction: reaction with the spectator ions removed Balancing Ionic Redox Reactions: Li + Zn +2  Li +1 + Zn The reaction may looked balanced, because there is conservation of ___________________, but the reaction does not show a conservation of ______________ MASS (same # of atoms/ions on both sides) CHARGE MASS CHARGE

VII. Balancing Ionic Redox Reactions Examples: Balance the following ionic redox reactions 1. Determine how much charge changes for each species 2. Multiply each species by how much the other species changes charge  ____Li + _____ Zn +2  _____ Li +1 + _____ Zn 2 1 2 1 2 (0) + 1(+2) = 2(+1) + 1(0) +2 = +2 Mass and charge conserved

VII. Balancing Ionic Redox Reactions Examples: Balance the following ionic redox reactions 1. Determine how much charge changes for each species 2. Multiply each species by how much the other species changes charge  _____Cu + + _____ Fe +3  _____ Cu +2 + _____ Fe 1 3 1 3 3 (+1) + 1(+3) = 3(+2) + 1(0) +6 = +6 Mass and charge conserved

VII. Balancing Ionic Redox Reactions Examples: Balance the following ionic redox reactions 1. Determine how much charge changes for each species 2. Multiply each species by how much the other species changes charge  _____Al + _____ H +  _____ Al +3 + _____ H2 3 2 6 2 3 (0) + 6(+1) = 2(+3) + 3(0) +6 = +6 Mass and charge conserved

TOPIC 2: Electrochemistry Objective: Predict which redox reactions will be spontaneous based on their position on the activity series table, identify the parts of a voltaic and electrolytic cell and describe their purpose, explain why the changes in concentration in each half cell are occurring as the reaction proceeds, and design a voltaic and electrolytic cell of your own.

VIII. Table J (Revisited) TOP The metal closer to the _________ of Table J are ____________________ and, therefore, more likely to be _________________. The nonmetals that are more reactive are more likely to be ______________________ Any ____________ on Table J will react ________________ with a ______________________ that is ______________________ This means that the metal with a charge must be_______________ the ____________________ metal for there to be a reaction Activity Series ANIMATION MORE REACTIVE OXIDIZED (lose e-) REDUCED (gain e-) METAL SPONTANEOUSLY METAL ION BELOW IT BELOW NEUTRAL

VIII. Table J (Revisited) Example: According to Table J, which of the following will react spontaneously? X + H+  H2 + X+ A) Pb +2 B) Sn+2 C) Fe D) H2 Example: Which of the following will replace Ni+2 in the compound Ni(NO3)2? A) Sn +2 B) Pb+2 C) Sn D) Cr   Example: Which atom/ion pair will Co oxidize spontaneously under standard conditions? A) Co + Fe+2 B) Co + Pb+2 C) Co + Cu D) Co + Cr+3

IX. Spontaneous Reactions on Voltaic Cells Voltaic Cell - a redox reaction whose two half-reactions are carried out separately, and the electrons given off by the oxidation half-reaction are used to power a device, and then given to the reduction half-reaction. ALSO CALLED A BATTERY (9v) OR CELL (AAA, AA, C, D) Voltaic Cell Animation and another set of animations: Complete Cell, Anode, Cathode, Salt Bridge How a Voltaic Wet Cell Works: Converts _____________ energy to ______________ energy by the use of a _______________________________  __________________ move from the metal that is ____________ to the metal that is _____________ Flow of electrons: ________________________________ CHEMICAL ELECTRICAL SPONTANEOUS REDOX REACTION ELECTRONS OXIDIZED REDUCED ANODE  CATHODE

IX. Spontaneous Reactions on Voltaic Cells Voltaic Cell - a redox reaction whose two half-reactions are carried out separately, and the electrons given off by the oxidation half-reaction are used to power a device, and then given to the reduction half-reaction. ALSO CALLED A BATTERY (9v) OR CELL (AAA, AA, C, D) Voltaic Cell Animation and another set of animations: Complete Cell, Anode, Cathode, Salt Bridge

IX. Spontaneous Reactions on Voltaic Cells Parts of a Voltaic Wet Cell:  Electrode:   Anode: Cathode: REMEMBER: A PIECE OF METAL WHERE OX. OR RED. OCCURS ELECTRODE WHERE OXIDITION OCCURS (SIZE ↓ ) ELECTRODE WHERE REDUCTION OCCURS (SIZE ↑ ) “RED CAT and AN OX” REDUCTION OCCURS ON CATHODE ANODE HAS OXIDATION

IX. Spontaneous Reactions on Voltaic Cells Salt Bridge: External Circuit: + ions from bridge move into cathode half cell -- ions from bridge move into anode half cell Ions from the salt brdige flow to the solutions completing the circuit Keeps the overall charge neutral in the half cells WIRE THAT CONDUCTS THE e-

X. Nonspontaneous Reactions on Electrolytic Cells Electrolytic Cell Animation

X. Nonspontaneous Reactions on Electrolytic Cells Parts of an Electrolytic Cell: Anode:   Cathode Battery: Electrolyte: Flow of electrons: Metal that is used to plate an object (+) Metal that is being plated (--) Produces energy to force nonspontaneous reaction to occur Allow ions to flow between electrodes e – start at anode and move toward the object being plated

X. Nonspontaneous Reactions on Electrolytic Cells Parts of a Electrolytic Cell: Uses a ________________ redox reaction (the reaction _____________ occur on its own) Uses ____________ to____________________ This process is called ________________________ Used for: 1) 2) 3)   nonspontaneous will NOT electricity force a chemical reaction to occur electrolysis electroplating Isolation of an element in a compound (NaClNa+Cl2) Purification of an element

X. Nonspontaneous Reactions on Electrolytic Cells Electroplating (see electroplating video)  

X. Nonspontaneous Reactions on Electrolytic Cells The Car Battery (or rechargeable batteries) use both types of cells Spontaneous Reaction to start the car Nonspontaneous Reaction to charge the battery VIDEO  

XI. Similarities and Differences Between Voltaic and Electrolytic Cells   Similarities “AN OX and RED CAT” e – flow from anode to cathode (+) ions flow toward the cathode, (--) ions flow toward the anode Anode gets smaller, cathode gets larger

XI. Differences Between Voltaic and Electrolytic Cells ____________________________ _____________________________ Anode is ____________________ Cathode is ___________________ Anode is ___________________ Cathode is ____________________ ____________________ electricity __________________ electricity _________________ REACTION __________________ REACTION ½ Reactions are for… requires a battery is an electric cell negative positive positive negative USES PRODUCES SPONTANEOUS NONSPONTANEOUS Two different elements (two different solutions) The same element (one solution both electrodes)