1Chemistry 2C Lecture 2: March 31 st, 2010 I.Acidic Redox Balancing II.Net Ionic Equations III. Balancing Redox rxs. in basic solutions IV. Electrochemistry Intro Lecture 2: Electrochemistry Should be able to balance the Thermite reaction Fe 2 O 3 (s) + Al (s) -> Fe (l) + Al 2 O 3 (s) OIL RIG

2Chemistry 2C Lecture 2: March 31 st, 2010 From last lecture: Balancing Redox Rxs. General method (for acidic aqueous rxs): 1. Write half reactions separately by first identifying the oxidized and reduced species 2. Balance element other than H or O, if any 3. Balance O by adding H 2 O molecules 4. Balance H by adding H + 5. Balance charge by adding electrons to the more positive side of the half-reaction 6. Multiple the half reaction by the whole number factors so that the number of electrons lost equal the number of electrons gained. 7. Add the two half reactions together and cancel common terms occurring on both sides of the arrow.

3Chemistry 2C Lecture 2: March 31 st, 2010 Balancing Redox Rxs. Fe 2+ (aq) + Cr 2 O 7 2- (aq)  Fe 3+ (aq) + Cr 3+ (aq) Unbalanced Eq: Step 1. Write half reactions separately by first identifying the oxidized and reduced species Fe is being oxidized (loss of electrons) Cr is being reduced (gain of electrons) Oxidation state Change: +2  +3 Oxidation state Change: +6  +3 Fe 2+ (aq)  Fe 3+ (aq) Cr 2 O 7 2- (aq)  Cr 3+ (aq) Eq. 1 Eq. 2

4Chemistry 2C Lecture 2: March 31 st, 2010 Balancing Redox Rxs. Fe 2+ (aq) + Cr 2 O 7 2- (aq)  Fe 3+ (aq) + Cr 3+ (aq) Unbalanced Eq: Step 2. Balance element other than H or O, if any Fe 2+ (aq)  Fe 3+ (aq) Cr 2 O 7 2- (aq)  Cr 3+ (aq) Eq. 1 Eq. 2 (old) Cr 2 O 7 2- (aq)  2Cr 3+ (aq) Eq. 2 (new)

5Chemistry 2C Lecture 2: March 31 st, 2010 Balancing Redox Rxs. Fe 2+ (aq) + Cr 2 O 7 2- (aq)  Fe 3+ (aq) + Cr 3+ (aq) Unbalanced Eq: Step 3. Balance O by adding H 2 O molecules Fe 2+ (aq)  Fe 3+ (aq) Eq. 1 Cr 2 O 7 2- (aq)  2Cr 3+ (aq) Eq. 2 (old) Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 0 Eq. 2 (new)

6Chemistry 2C Lecture 2: March 31 st, 2010 Balancing Redox Rxs. Fe 2+ (aq) + Cr 2 O 7 2- (aq)  Fe 3+ (aq) + Cr 3+ (aq) Unbalanced Eq: Step 4. Balance H by adding H + Fe 2+ (aq)  Fe 3+ (aq) Eq. 1 Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 0 Eq. 2 (old) Eq. 2 (new) 14H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 0

7Chemistry 2C Lecture 2: March 31 st, 2010 Balancing Redox Rxs. Fe 2+ (aq) + Cr 2 O 7 2- (aq)  Fe 3+ (aq) + Cr 3+ (aq) Unbalanced Eq: Step 5. Balance charge by adding electrons to the more positive side of the half-reaction Fe 2+ (aq)  Fe 3+ (aq) Eq. 1 (old) Fe 2+ (aq)  Fe 3+ (aq) + e - Eq. 1 (new) Eq. 2 (old) 6e H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 0 Eq. 2 (new) 14H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 0

8Chemistry 2C Lecture 2: March 31 st, 2010 Balancing Redox Rxs. Fe 2+ (aq) + Cr 2 O 7 2- (aq)  Fe 3+ (aq) + Cr 3+ (aq) Unbalanced Eq: Step. 6 Multiple the half reaction by the whole number factors so that the number of electrons lost equal the number of electrons gained Fe 2+ (aq)  Fe 3+ (aq) + e - Eq. 1 (old) 6e H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 0 Eq. 2 6Fe 2+ (aq)  6Fe 3+ (aq) + 6e - Eq. 1 (new) The number of electrons in oxidized half reaction is equal to the number of electrons in the reduced half reaction (6).

9Chemistry 2C Lecture 2: March 31 st, 2010 Balancing Redox Rxs. Fe 2+ (aq) + Cr 2 O 7 2- (aq)  Fe 3+ (aq) + Cr 3+ (aq) Unbalanced Eq: Step 7: Add the two half reactions together and cancel common terms occurring on both sides of the arrow 6e H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 0 Eq. 2 6Fe 2+ (aq)  6Fe 3+ (aq) + 6e - Eq. 1 Electrons are never a species in a balanced redox equation!!!! Sum 6e H + (aq) + Cr 2 O 7 2- (aq) + 6Fe 2+ (aq)  2Cr 3+ (aq) + 7H Fe 3+ (aq) + 6e -

10Chemistry 2C Lecture 2: March 31 st, 2010 Net Ionic Equations What is a net ionic equation? An equations that concentrates on the true reacting species and ignores non-interacting species that are just “hanging out” Suppose Mg metal is treated with hydrochloric acid, H 2 gas is evolved and MgCl 2 can be isolated from the evaporated solution (not precipitated). The balanced reaction may look like this: Mg (s) + 2HCl (aq) -> MgCl 2 (aq) + H 2 (g) Oxidized (loss of electrons) Reduced (gain of electrons)

11Chemistry 2C Lecture 2: March 31 st, 2010 Net Ionic Equations But HCl is a strong acid, (hence a strong electrolyte) so it really should be written as H + (aq) + Cl - (aq) Similarly MgCl 2 is soluble in H 2 0, so this should be written as Mg +2 (aq) and Cl - (aq) So overall the balance reaction shouldn’t be Mg (s) + 2HCl (aq) -> MgCl 2 (aq) + H 2 (g) But really Mg (s) + 2H + (aq) + 2Cl - (aq) -> Mg +2 (aq) + Cl - (aq) + H 2 (g)

12Chemistry 2C Lecture 2: March 31 st, 2010 Net Ionic Equations In this equation Mg (s) + 2H + (aq) + 2Cl - (aq) -> Mg +2 (aq) + Cl - (aq) + H 2 (g) The chloride ion is not participating with the redox reaction (it is neither being reduced nor being oxidized). This is commonly referred to as a “spectator ion” and the “net ionic equation” is really Mg (s) + 2H + (aq) -> Mg +2 (aq) + H 2 (g)

13Chemistry 2C Lecture 2: March 31 st, 2010 Electrochemistry Premise: We know that a large number of reactions fall into the redox class, which involve the transfer of electrons from a species (atom) that is being oxidized to a species (atom) that is being reduced. Such charge transfer reactions are ubiquitous in nature! -Photosynthesis (light to chemical energy) -Photovoltiacs (light to electric energy) -Muscles (energy released by redox rx. causes motion) -Dynamite (redox rx. that generate a lot of energy) -Combustion of fuels (oxidation) -Batteries (redox rx. To make electricity) -Corrosion, chrome plating, “tin” cans (sacrificial metals with boats)

14Chemistry 2C Lecture 2: March 31 st, 2010 Electrochemistry Premise: We know that a large number of reactions fall into the redox class, which involve the transfer of electrons from a species (atom) that is being oxidized to a species (atom) that is being reduced. Terms: Oxidizing agents: Species that take electrons and are themselves reduced (e.g. they oxidize species) Reducing agents: Species that give electrons and are themselves oxidized (e.g. they reduce species)

15Chemistry 2C Lecture 2: March 31 st, 2010 Electrochemistry From the list above, it is clear that many redox reactions don’t occur in the simplified acidic aqueous environment that was discussed last lecture. For example, we may require the balancing of a redox reaction in a basic solution. How? The easiest way is to follow the protocol for balancing redox rxs. for acidic aqueous solutions, but you must use OH - to neutralize the H + species (after step 4).

16Chemistry 2C Lecture 2: March 31 st, 2010 Balancing basic Aquaeous redox rxs. Following the protocol with acidic aqueous rxs. S 2 O 3 2- (aq) + Cl 2 (g) -> HSO 4 - (aq) + Cl - (aq) +2 The sulfur in S 2 O 3 2- is being oxidized (loses e - ) The chlorine in Cl 2 is being reduced (gains e - ) 0+6 The strength of a oxidizing agent (or reducing agent) varies depending on the nature (and thermodynamics) of the species. Cl 2 is a stronger oxidizing agent than I 2 which was used in your lab. Step 1. Write half reactions separately by identifying the oxidized and reduced species

17Chemistry 2C Lecture 2: March 31 st, 2010 Balancing basic Aquaeous redox rxs. S 2 O 3 2- (aq) + Cl 2 (g) -> HSO 4 - (aq) + Cl - (aq) Cl 2 (g) -> Cl - (aq) S 2 O 3 2- (aq) -> HSO 4 - (aq) Eq. 1 (reduction): Eq. 2 (oxidation): Step 1. Write half reactions separately by identifying the oxidized and reduced species Step 2. Balance element other than H or O, if any Cl 2 (g) -> 2Cl - (aq) S 2 O 3 2- (aq) -> 2HSO 4 - (aq) Eq. 1 (reduction): Eq. 2 (oxidation):

18Chemistry 2C Lecture 2: March 31 st, 2010 Balancing basic Aquaeous redox rxs. Cl 2 (g) -> 2Cl - (aq) 5H 2 0 (l) + S 2 O 3 2- (aq) -> 2HSO 4 - (aq) Eq. 1 (reduction): Eq. 2 (oxidation): Eq. 1 (reduction): Eq. 2 (oxidation): Step 3. Balance O by adding H 2 O molecules Step 4. Balance H by adding H + Cl 2 (g) -> 2Cl - (aq) 5H 2 0 (l) + S 2 O 3 2- (aq) -> 2HSO 4 - (aq) + 8H + (aq)

19Chemistry 2C Lecture 2: March 31 st, 2010 Balancing basic Aquaeous redox rxs. Step 4New. Add OH- to both sides to neutralize any H+ Eq. 1 (reduction): Eq. 2 (oxidation): Cl 2 (g) -> 2Cl - (aq) 8OH - (aq) + 5H 2 0 (l) + S 2 O 3 2- (aq) -> 2HSO 4 - (aq) + 8H + (aq) + 8OH - (aq) Eq. 2 (oxidation):8OH - (aq) + 5H 2 0 (l) + S 2 O 3 2- (aq) -> 2HSO 4 - (aq) + 8H 2 O (l) The new step of adding OH - to each side of the equation to neutralize H + species. This is the primary difference between balancing redox equations in acidic or basic aqueous environments!

20Chemistry 2C Lecture 2: March 31 st, 2010 Balancing basic Aquaeous redox rxs. Eq. 1 (reduction): 2e - + Cl 2 (g) -> 2Cl - (aq) Eq. 2 (oxidation):8OH - (aq) + S 2 O 3 2- (aq) -> 2HSO 4 - (aq) + 3H 2 0 (l) + 8e - Step 5. Balance charge by adding electrons to the more positive side of the half-reaction Step. 6 Multiple the half reaction by the whole number factors so that the number of electrons lost equal the number of electrons gained Multiple Eq. 1 by 4 to get equal number of electrons gained and lost. Eq. 1 (reduction):8e - + 4Cl 2 (g) -> 8Cl - (aq)

21Chemistry 2C Lecture 2: March 31 st, 2010 Balancing basic Aquaeous redox rxs. Eq. 1 (reduction): 8e - + 4Cl 2 (g) -> 8Cl - (aq) Eq. 2 (oxidation):8OH - (aq) + S 2 O 3 2- (aq) -> 2HSO 4 - (aq) + 3H 2 0 (l) + 8e - Step 7: Add the two half reactions together and cancel common terms occurring on both sides of the arrow Total rx: 8OH - (aq) + S 2 O 3 2- (aq) + 8e - + 4Cl 2 (g) -> 2HSO 4 - (aq) + 8Cl - (aq) + 3H 2 0 (l) + 8e -

22Chemistry 2C Lecture 2: March 31 st, 2010 Disproportionation Rxss Special reactions in which the same substance is both oxidized and reduced (obviously not the same specific atom) 2H 2 O 2 (aq) -> 2H 2 O (l) + O 2 (g) 2Cu + (aq) -> Cu (s) + Cu +2 (aq) Exercise: Balance in a basic solution ClO 2 -> Cl0 3 - (aq) +Cl - (aq) H 2 O 2 (aq) -> H 2 O (l) H 2 O 2 (aq) -> O 2 (g) Eq. 1 (reduction): Eq. 2 (oxidation): Cu + (aq) -> Cu (s) Cu + (aq) -> Cu +2 (aq) Eq. 1 (reduction): Eq. 2 (oxidation):

23Chemistry 2C Lecture 2: March 31 st, 2010 Electrochemistry Is an old, but still important science!!! 1786: Many say that eletrochemistry had its start with Luigi Galvani who plugged in poor froggies with different metals. The froggies muscles moved and thus he came to the conclusion that electricity was a “life force.” Mid-18 th century: Ben Franklin had many theories about positive and negative electricity. His famous kite expt. was in He pushed the approach that electricity was like a fluid that passed from one body to another

24Chemistry 2C Lecture 2: March 31 st, 2010 Electrochemistry Late 1700’s: Luigi Galvani friend Alessandro Volta realized that the animal electricity was really a reaction from different metals. He made the first battery by stacking Ag and Zn disks separated by a wet cloth containing a salt or weak acid solution. These people gave us their names: “Galvanic cell”, “Voltaic cell”, Volts, galvanization. Question for next lecture: Why was is important to use wet, salty clothes in Volta’s experiment?