Oxidation/Reduction (REDOX) Reference Table: Periodic Table & J PACKET #12: Oxidation/Reduction (REDOX) Reference Table: Periodic Table & J

Before we start, recall . . . each box on the periodic table has one or multiple oxidation numbers on the upper right hand corner of the box. The oxidation number is the charge that an atom takes on after either losing or gaining one or more electrons in order to most efficiently stabilize its outer most shell (act like a Noble Gas) using the least amount of energy. Any free element (not combined with any other element) has an oxidation number of 0. Ex: Na = Na0

Oxidation States Oxidation States are assigned to atoms to identify how many electrons are either gained or lost by an atom For example metals in Group 2 (like Ca) have a +2 oxidation state. Therefore Metals in Group 2 lose two electrons when they form compounds Changes in oxidation numbers indicate that a redox reaction has occurred. It is important to learn the rules for assigning oxidation states to atoms in order to determine whether oxidation or reduction has occurred.

(oxidation is loss, reduction is gain) JUST REMEMBER . . . “OIL RIG” (oxidation is loss, reduction is gain)

Oxidation loss of electrons by a molecule, atom, or ion Remember you learned that metals lose electrons to have a complete outer shell (more stable) For example Na will lose one electron to become Na+ ion Therefore Na is Oxidized When any atom loses electrons its oxidation number INCREASES Example: Na0  Na+ + e- -its oxidation # increases from 0 to + 1

Reduction The gain of electrons by a molecule, atom, or ion Nonmetals gain electron to have a complete outer shell Phosphorous will gain three electron to have 8 in its outer shell Because P gains 3 electrons P is Reduced When any atoms gains electrons (negative particles) the oxidation number DECREASES Example: P0 + 3e-  P-3 -its oxidation # decreases from 0 to - 3

REDOX Short-hand for an oxidation/reduction equation. In a single reaction there is both oxidation and reduction. REDOX reactions have conservation of matter, and conservation of charge. The numbers of elements are equal on both the reactant and the product side, and the total charge on both the reactant and product side equal 0. Na0  Na+ + e- P0 + 3e-  P-3 3Na + P  N3P (sodium phosphide) Na is being oxidized; P is being reduced Electrons to the left of the arrow means that they are being gained, to the right of the arrow means they are being lost.

REDOX In a REDOX reaction, when something is being oxidized, it is called the reducing agent, and when something is being reduced it is called the oxidizing agent Reducing Agent: an electron donor Oxidizing Agent: an electron acceptor

Oxidation State Rules 1) Free elements (not combined with any other element) have an oxidation number of zero; includes diatomic molecules. Ex: Na, O2, H2, Cl2 2) All metals in Group 1 have an oxidation number of +1. 3) All metals in Group 2 have an oxidation number of +2. 4) F (fluorine) always has an oxidation of -1

Oxidation State Rules 5) The oxidation of simple ions is equal to the charge on the ion. Ex: Mg+2 has an oxidation number of +2. 6) The sum of the oxidation numbers must equal 0 Examples: NaCl & MgCl2 7) In polyatomic ions, the sum of the oxidation numbers of all the atoms must equal the charge of the ion. Example: sulfate ion SO4-2 . Oxygen has an oxidation of -2, and therefore (-2) x (+4) = -8. Remember that the overall charge of this ion has to be -2. What is the oxidation # of sulfur?

Oxidation State Rules 8) In general, oxygen has an oxidation number of -2. Oxygen has an oxidation number of -1 in peroxides (O2-2) Example: H2O2. Oxygen has an oxidation number of +2 in compounds with fluorine Example: OF2 9) Hydrogen has an oxidation number of +1 in all compounds combined with a non-metal. Hydrogen has an oxidation number of -1 when it is a metal hydrides (metal and hydrogen. Example: LiH, and CaH2

Assigning Oxidation Numbers When you have a binary compound or polyatomic compound, assigning oxidation numbers is a little easier than when you have to assign oxidation numbers to compounds with more than two elements. EXAMPLES: NaCl CaO CO2 Li3N

Assigning Oxidation Numbers Dr. McGuiness’ “bookend” technique to assigning oxidation numbers to compounds with more than two elements. Identify the oxidation # of the last element (overall charge) Identify the oxidation # of the first element (overall charge) If there is no charge to the compound, then the overall charge must be 0, therefore you can determine the oxidation # of the element in the middle. LiMnO4 (+1) + (?) + (-8) = 0 -2 x 4 = -8 +1 x 1 = +1

Half-Reactions Once you can assign oxidation numbers, then you can take a REDOX reaction and break it down to half-reactions in order to figure out what is being oxidized and what is being reduced. Half-reaction: shows either oxidation or reduction. A redox reaction is made up of two half reactions (one oxidation, and one reduction). A redox reaction is when the oxidation number of the half-reaction and the reduction half-reaction occur simultaneously.

Single Replacement Reactions are ALWAYS REDOX Double Replacement Reactions are NEVER REDOX Synthesis and Decompostion may or may be REDOX – must assign the oxidation numbers to check.

Half-Reactions Cu + Ag(NO3) → Cu(NO3)2 + Ag Example: Assign oxidation numbers to everything See which oxidation numbers change from reactant side to product side. Determine which half-reaction is oxidation (loses e-), and which is reduction (gains e-) OX: Cu0 → Cu2+ + 2e- RED: Ag+ + 1e- → Ag0

Balancing REDOX reactions Write half reactions for oxidation and reduction. Balance the half reactions separately. Equalize the number of electrons gained and lost. Add the half reactions together and cancel the electrons to get the final balanced equation. The electrons lost in a REDOX reaction have to equal the electrons gained

Balancing REDOX reactions Cu + 2Ag(NO3) → Cu(NO3)2 + 2Ag OX: Cu0 → Cu2+ + 2e- RED: 2(Ag+ + 1e- → Ag)  2Ag+ + 2e- 2Ag0 When copper is oxidized, it loses 2 electrons, which are gained by the silver ion. Copper = reducing agent Silver = oxidizing agent

Let’s Practice  For the following REDOX reactions complete the following: Write the correct half-reactions Identify what is oxidized/what is reduced Identify the oxidizing/reducing agents Balance the equation

1. Ce4+ + Sn2+ → Ce3+ + Sn4+ 2. Fe + SnBr4 → FeBr2 + Sn 3. Cu + Ag(NO3) → Cu(NO3)2 + Ag

Activity Series – Table J Table J compares how active each metal and nonmetal is. Metals higher up are more active, and replace metals from below them from compounds (remember single replacement???) In a single replacement, the free element has to be more reactive than the element in compound in order for the reaction to be spontaneous. If it isn’t – the reaction does NOT GO!!

Spontaneous or Not?? Li + KCl  Ca + MgCO3  K + LiCl  F2 + 2NaCl  Cl2 + 2NaF 

Recall that metals lose electrons, and therefore undergo oxidation Metals that are more reactive oxidize easier Recall that non-metals gain electrons, and therefore undergo reduction Non-metals that are more reactive reduce easier

Electrochemical Cells There are two types of electrochemical cells, Voltaic & Electrolytic These cells rely on REDOX reaction in different ways to either generate energy or to separate elements in compound that would normally not exist on their own in nature.

(chemical energy  electrical energy) Voltaic Cell A voltaic cell uses REDOX reactions that are spontaneous to produce electricity (chemical energy  electrical energy) A battery is an example of a voltaic cell.

Voltaic Cell There are two half-cells in a voltaic cell (anode & cathode) Each half-cell is a metal strip called an electrode. There is a wire connecting the two electrodes in which electrons travel through. Electrons always travel from the anode  cathode

Voltaic Cell There is a salt-bridge connecting the two half cells in order to permit ions to flow between the two half-cells. Ions travel through the salt bridge that are connecting the two half cells (a complete or closed circuit)

Wire Electrodes Half-Cells Salt Bridge Half-Cells

Voltaic Cell The anode is the half-cell where oxidation ALWAYS occurs. In a voltaic cell, the anode is negatively charged ( ANODE/OXIDATION – An Ox) The anode will always be the half-cell with the metal electrode that is most reactive (Table J)

Voltaic Cell The cathode is the half-cell where reduction ALWAYS occurs. In a voltaic cell, the cathode is positively charged (CATHODE/REDUCTION - Red Cat)

Voltaic Cell Zn + Cu+2  Cu + Zn+2 Let’s consider this REDOX reaction. After writing the half-reactions, we can determine which half cell is the anode, and which is the cathode. Also, Zn is more reactive than Cu on Table J.

(electrical energy  chemical energy) Electolytic Cell The REDOX reaction in an electrolytic cell is non-spontaneous, and therefore electrical energy (battery) is required to induce a chemical reaction (electrical energy  chemical energy)

Anode (+) Cathode (-) Battery NaCl  Na+ + Cl-

Electrolytic Cell Electrolysis: the process in which electricity breaks down a compound. Example: NaCl The negative end of the battery is attached to the cathode (reduction) which makes it negatively charged; the positive end of the battery is attached to the anode (oxidation) which makes it positively charged. Positive ions (Na+) in solution are attracted to the cathode which is negatively charged. Negative ions (Cl-) in solution are attracted to the anode which is positively charged.

CATHODE (reduction): Na+ + e-  Na0 We have created neutral sodium which is never found in nature (so reactive). ANODE (oxidation) Cl-  e- + Cl0 We have created neutral chlorine which is never found in nature (so reactive)

Voltaic Electrolytic Non-spontaneous Electrical  Chemical Chemical  Electrical Energy Anode – oxidation (-) Cathode – reduction (+) Example: BATTERY Electrons travel from anode to cathode (wire) More reactive metal ALWAYS the site of oxidation Salt bridge is for the flow of ions from one half-cell to another Non-spontaneous Electrical  Chemical Requires a battery as a source of energy Anode – oxidation (+) Cathode – reduction (-) (+) charged ion moves toward the cathode (-) charged ion moves toward the anode

Electroplating Another example of an electrolytic cell. The process of electroplating requires a layer of metal such as silver or copper, coating or covering any object to be plated (spoon or fork) The item being plated is the cathode - reduction/(-) The electrode must be the same metal that you are plating the object in. Cathode (-)

Review Questions 1) In which substance does chlorine have an oxidation number of +1? A) HClO2 B) HClO C) Cl2 D) HCl 2) What is the oxidation state of nitrogen in NaNO2? A) +1 B) +2 C) +3 D) +4

Mg(s) + 2H+(aq) + 2Cl-(aq)  Mg+2(aq) + 2Cl-(aq) + H2(g) 3) Which balanced equation represents a redox reaction? A) AgNO3 + NaCl  AgCl + NaNO3 B) BaCl2 + K2CO3  BaCO3 + 2KCl C) CuO + CO  Cu + CO2 D) HCl + KOH  KCl + H2O 4) Given the reaction: Mg(s) + 2H+(aq) + 2Cl-(aq)  Mg+2(aq) + 2Cl-(aq) + H2(g) Which species undergoes oxidation? A) Cl-(aq) B) Mg(s) C) H+(aq) D) H2(g)

Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) 5) Given the reaction: Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) Which statement correctly describes what occurs when this reaction takes place in a closed system? A) There is a net gain of mass. B) Atoms of Zn(s) lose electrons and are oxidized. C) Atoms of Zn(s) gain electrons and are reduced. D) There is a net loss of mass. 6) Given the reaction that occurs in an electrochemical cell: Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s) During this reaction, the oxidation number of Zn changes from A) 0 to +2 B) 0 to -2 C) +2 to 0 D) -2 to 0

__Br2 + __Sn  __Br- + __Sn+2 7) Given the unbalanced equation: __Br2 + __Sn  __Br- + __Sn+2 When the equation is correctly balanced using the smallest whole-number coefficients, the coefficient of Br- is A)1 B) 2 C)3 D) 4 8) What is the reducing agent in the reaction: Pb + 2AgNO3  Pb(NO3)2 + 2Ag? A) Pb B) NO-3 C) Ag+ D) Ag 9) According to the Activity Series chemistry reference table, which molecule is most easily reduced? A) I2 B) Br2 C) Cl2 D) F2

10) Given the reaction: 4Al(s) + 3O2(g)  2Al2O3(s) (a) Write the balanced oxidation half-reaction for this oxidation-reduction reaction. (b) What is the oxidation number of oxygen in Al2O3?

11) The diagram below represents a voltaic cell at 298 K and 1 atmosphere. When the switch is closed, electrons flow from A) Mg(s) to Ag(s) B) Ag(s) to Mg(s) C) Ag+(aq) to Mg+2(aq) D) Mg+2(aq) to Ag+(aq)

Questions 12 and 13 refer to the following: The diagram below shows a spoon that will be electroplated with nickel metal. 12) Does the chemical cell diagram represent a voltaic or an electrolytic cell? [Give one reason to support your answer.] 13) Does the spoon represent the anode or the cathode in this electrochemical cell? [Give one reason to support your answer.]