Redox Reactions

 Have you ever drank from an aluminum can?  Ever used a flashlight?  Use your calculator on a test recently?  Enjoy exercising?  Are you alive?

What do all of these things have in common?

ENERGY!!! And all forms of energy harnessing require an understanding of Redox Reactions.

Redox Notes Part I : Define oxidation and reduction.

Oxidation-Reduction Reactions  Aka Redox Rxns  Chemical reactions that transfer electrons from one chemical to another at the same time  OXIDATION – the LOSS of electrons in a chemical rxn  REDUCTION – the GAIN of electrons in a chemical rxn

Redox is a tandem process  Oxidation cannot happen without reduction & vice versa  Thus, you will have two types of agents (chemicals)  OXIDIZING AGENT – substance that oxidizes another. It gets reduced.  REDUCING AGENT - substance that reduces another. It gets oxidized.

Why can oxidation not happen without reduction?  One substance that donates the electrons needs a place for the electrons to travel to.  Electrons don’t just vanish; they attach to another atom.  Therefore, one substance donates and the other accepts the electron.

LEO the Lion and his Little e’s Meet Leo’s family. He loves them very much.

LEO goes hunting and loses his little e’s.

LEO goes GER Listen to him ROAR!!

LEO’s pride searches everywhere…

And finally finds them playing at the watering hole

LEO gains his little ‘e’s back again.

LEO the lion goes GER L oss of E lectrons = O xidation G ain of E lectrons = R eduction

Example 1:  Identify the movement of the electron(s) and then label the reaction as oxidation or reduction:  Al  Al e -

Example 2:  Identify the movement of the electron(s) (If +  - …gained e -. If -  +…lost e - ) and then label the reaction as oxidation or reduction:  2Br - + Cl 2  Br 2 + 2Cl -

Atoms of elements, ions or compounds gain or lose outermost electrons during reactions that form a new set of elements, ions or compounds. REDOX

 SC.912.P.8.10: Describe oxidation- reduction reactions in living and non-living systems.

 A transfer of electrons occurs during all single replacement or Combustion reactions  Sometimes in Double Replacement and Decomposition reactions REDOX Reactions Energy

LEO the lion says GER  Loss of Electrons is Oxidation _____ Gain of Electrons is Reduction OIL RIG  Oxidation Is Loss of electrons causing the oxidation number to pump up to a higher value.  Reduction Is Gain of electrons causing a reduction in the value of the oxidation number. Review definition of oxidation and reduction.

 Photosynthesis and cellular respiration are biological examples of Redox reactions.  Write the chemical equation for these reactions. Living Systems

 Fires, rusting, and metals reacting in acid are also examples of RedOx reactions  Lets see why these reactions are classified as “REDOX.” Non-Living Systems

 When unbonded elements react to form compounds, one of the elements gains electrons (would that be the metal or the nonmetal?) while the other loses electrons (would that be the metal or the nonmetal?).  Think about the formation of NaCl. 1. Write your thoughts in your notebook. 2. Share your ideas with your partner. 3. Share with the class. Think-Pair-Share

 Metals undergo Oxidation: loss of e-  Ex Na  Na + +e -  Reduction: gain of e-  Ex Cl 2 + 2e -  2Cl -  The Oxidizing Agent-substance reduced: Cl 2  The Reducing Agent-substance oxidized: Na REDOX Reaction Example

 1) __ Mg + __ N 2  __ Mg 3 N 2 __  2) __ Fe + __ O 2 → __ Fe 2 O 3 __ (rust)   3) __ Ca + __ O 2  __ CaO __  4) __ H 2 + __ O 2 → __ H 2 O __ Identify the REDOX reactions.

 5) __CH 4 __+__ O 2  __CO 2 __ + __H 2 O__  6) __ Cl 2 + __ KI → __ KCl __ + __ I 2  7) __ CaCO 3 __ + __ HCl __ → __ CaCl 2 __ + __ H 2 CO 3 __ __ __ Identify the REDOX reactions.

 8) __ CH 4 __ + __ O 2 → __ CO 2 __ + __ H 2 O __  9) __ AgNO 3 __ + __ Cu  __ Cu(NO 3 ) 2 __ + __ Ag __ __  10) __ C 6 H 12 O 6 __ + __ O 2 → __ CO 2 __ + __ H 2 O __ __ __ Identify the REDOX reactions.

 Single Replacement Reactions are always redox reactions! Combustion reactions are always redox reactions! Any time an oxidation number changes (which means electrons are gained or lost) during the reaction, a redox reaction is occurring. END OF CHEMISTRY 1 PPT

 You must assign oxidation numbers to all elements in the reaction.  Identify which elements oxidation number changed from reactants to products. Some electron transfers are not as easy to predict as metal with nonmetal. To figure out what is oxidized and what is reduced you can follow a plan. ( Honors )

 Oxidation number = number of e- gained or lost by an atom when it forms an ion  K + Br -  EX: 2K + Br 2  2KBr  Potassium is oxidzed from 0 to +1  Bromine is reduced from 0 to -1 Changes in oxidation number

 1. The oxidation # of an uncombined atom = 0.  2. The oxidation number of a monatomic ion is equal to the charge on the ion.  3. The oxidation number of the most electronegative element in a molecule is equal to the charge it would have if it were a ion with noble gas configuration.  4. F is always -1  5. O is always -2 (except in peroxides and when attached to F) Here are the Rules

 6. H is always +1 (except when attached to more electronegative metals, Li, Na, Ca, and Al  7. Group 1A, 2A, and 3A always have an oxidation number equal to the group number (equal to the charge it would have if it were a ion with noble gas configuration.)  8. Sum of all oxidation numbers in a neutral compound is 0  9. If not neutral, sum of all oxidation numbers is equal to the overall charge on ion Rules cont…

 Neutral elements that are not bonded to any other element have oxidation number of zero. Examples: Na (s), Cl 2(g), Hg (l) all have oxidation numbers of zero.  Group 1 metals that are bonded to other elements have an oxidation number of +1 (positive one). Examples: Na in NaCl is +1, Li in LiOH is +1  Group 2 metals that are bonded to other elements have an oxidation number of +2 (positive two). What would Mg in MgO be? Examples

 Group 2 metals that are bonded to other elements have an oxidation number of +2 (positive two). Examples: Ca in CaCl 2 is +2, Ba in Ba(OH) 2 is +2, Therefore, Mg in MgO would be +2  Oxygen that is bonded to other elements has an oxidation number of -2 (negative two) unless it is in a peroxide or bonded to F. Examples: O in CO 2 is -2, O in LiOH is -2, O in Na 3 PO 4 is -2  BUT WAIT… THERE ARE EXCEPTIONS Examples

 Oxygen that is bonded to other elements has an oxidation number of -2 (negative two) unless it is in a peroxide or bonded to F. Examples: O in CO 2 is -2 O in LiOH is -2 O in Na 3 PO 4 is -2 What is O in O 2 ? What is O in H 2 O?  Exception: O in HOOH, hydrogen peroxide, is -1 Examples

 Halogens that are bonded to other elements have an oxidation number of -1 (negative one) unless they are bonded to a more electronegative element such as a halogen closer to the top of the periodic table. WAIT! What happens to oxygen when it is bonded to fluorine?  That’s right, oxygen must be positive in this super rare case! Fluorine is always -1, Why? What is O in O 2 ? zero. What is O in H 2 O? -2

 Halogens that are bonded to other elements have an oxidation number of -1 (negative one) unless they are bonded to a more electronegative element such as a halogen closer to the top of the periodic table. Examples: Cl in NaCl is -1 Cl in PCl 5 is -1 Cl in CaCl 2 is -1  Exception: Cl in ClF 5 is +5,  Cl in ClBr 6 is -6 while Br in ClBr 6 is +1 OK – Let’s face it, the larger halogens are only easily predictable when bonded with a metal, otherwise, much thinking is required! Examples

 A compound has an overall charge of zero, which means all the negative charges have to equal the positive charges.  Examples: When calculating the oxidation number of N in NO 2, use the rules above to help you. You see that oxygen normally has an oxidation number of -2 and there are two oxygen atoms. 2(-2) = -4. The total number of negative charges is 4 negatives. The only other atom that is present is nitrogen. That means the nitrogen is responsible for all for the positive charge. X + -4 = 0. X = +4. Therefore, the oxidation number on N in NO 2 is +4.  The oxidation number of C in CO (carbon monoxide) is +2.  The oxidation number of C in CO 2 (carbon dioxide) is +4.  The oxidation number of P in PCl 3 (phosphorous trichloride) is +3.  The oxidation number of P in P 2 O 5 (diphosphorous pentoxide) is +5.  Mn in MnO 2 is +4. Compounds are Neutral

 An ion has an overall charge equal to the charge of the ion. That means the positive charges will NOT equal the negative charges, but instead, when you add all the charges together the sum will be equal to the charge of the ion.  Example: The Mn in permanganate ion, MnO 4 -, is =7 (Here is how: X + 4( - 2)= - 1) X = - 1 add 8 to both sides.  X = 7 Polyatomic Ions

H in HCl is +1 H in BH 3 is -1 Hydrogen bonded to a metal is assigned -1, and hydrogen bonded to nonmetal is +1.

a) HCl b) KNO 3 c) OH - d) Mg 3 N 2 e) I 2 Assign oxidation numbers to each element in the element, compound or ion. DO NOW!

f) ClO 3 - g) Al(NO 3 ) 3 h) S 8 i) H 2 O 2 j) PbO 2 Assign oxidation numbers to each element DO NOW!

k) NaHSO 4 l) SO 3 2- m) O 2 n) KMnO 4 o) LiH Assign oxidation numbers to each element. More Practice!

p) Fe 2 O 3 q) SO 3 r) NH 4 + s) H 2 SO 4 t) Na Assign oxidation numbers to each element in the compounds listed below. More Practice!

Have you MASTERED THIS SKILL? Try it in the context of a chemical equation!! Assign oxidation numbers to each element in the compounds listed below. Assess yourself!

 1) __ Mg + __ N 2  __ Mg 3 N 2 __  2) __ Fe + __ O 2 → __ Fe 2 O 3 __  3) __ Ca + __ O 2  __ CaO __  4) __ H 2 + __ O 2 → __ H 2 O __ Assign oxidation numbers to each element

 5) __CH 4 __+__ O 2  __CO 2 __ + __H 2 O__  6) __ Cl 2 + __ KI → __ KCl __ + __ I 2  7) __ CaCO 3 __ + __ HCl __ → __ CaCl 2 __ + __ H 2 CO 3 __ __ __ Assign oxidation numbers to each element

 8) __ CH 4 __ + __ O 2 → __ CO 2 __ + __ H 2 O __  9) __ AgNO 3 __ + __ Cu  __ Cu(NO 3 ) 2 __ + __ Ag __ __ Assign oxidation numbers to each element

 Fe + CuSO 4  Cu + Fe 2 (SO 4 ) 3  -omit spectator ions-ions that don’t change their O# Example: SO 4 stays at -2  Fe + Cu +2  Cu + Fe +3  Split into half reactions and balance electron transfer  Fe  3e -1 + Fe +3  Cu +2 +2e -1  Cu Example: 6 electrons is the least common multiple. Balance using Half Reactions

 Split into half reactions and balance electron transfer  2(Fe  3e -1 + Fe +3 )  3(Cu +2 +2e -1  Cu) The iron lost a total of six electrons as it was oxidized  2Fe  6e Fe +3  3Cu +2 +6e -1  3Cu The balanced reaction is  2Fe + 3Cu +2  2 Fe Cu Balance using Half Reactions Fe + CuSO 4  Cu + Fe 2 (SO 4 ) 3

1) __ Mg + __ N 2  __ Mg 3 N 2 __ 2) __ Fe + __ O 2 → __ Fe 2 O 3 __ 3) __ Ca + __ O 2  __ CaO __ 4) __ H 2 + __ O 2 → __ H 2 O __ Balance each reaction

5) __CH 4 __+__ O 2  __CO 2 __ + __H 2 O__ 6) __ Cl 2 + __ KI → __ KCl __ + __ I 2 7) __ CaCO 3 __ + __ HCl __ → __ CaCl 2 __ + __ H 2 CO 3 __ __ __ Balance each reaction.

8) __ CH 4 __ + __ O 2 → __ CO 2 __ + __ H 2 O __ 9) __ AgNO 3 __ + __ Cu  __ Cu(NO 3 ) 2 __ + __ Ag __ __ Balance each reaction

Self Assessment.  What happens during oxidation and reduction?  What types of reactions are also redox reactions?  (Honors Extension) How can you balance redox equations?