RULES Assigning Oxidation Numbers Examples 1. Each Uncombined Element has an 2Na + Cl 2 2NaCl oxidation number = 0 Na = 0 or written Na 0 Cl 2 = 0 or written Cl 0 Monatomic ions have an oxidation number 2 equal to the ionic charge. Al 3+ then Al = +3 2. The metals of Group 1 always have an oxidation number of +1 and the metals of Group 2 always have an oxidation number of +2. Group 1 = +1 Group 13 = +3 Group 2 = +2 In KCl: K=+1 In CaCl 2 : Ca=+2 In MgO; Mg=+2 In Li 2 O: Li=+1 3. Fluorine is ALWAYS = -1 The other halogens are also = -1 when they are the most electronegative element in the compound In HF: F = -1 In O 2 F: F = -1 In CaCl 2 : Cl = -1 In HClO; oxygen is more electronegative than Cl Cl=+1
RULES Assigning Oxidation Numbers 4. Hydrogen is +1 in a compound EXCEPT when it is combined with a METAL. When combined with a METAL, Hydrogen = -1 Examples In HF; H=+1 In H 2 SO 4 ; H=+1 In CaH 2 : H=-1 In LiH: H=-1 5. Oxygen is -2 EXCEPT when with Fluorine, then Oxygen = +2 EXCEPT in peroxide ion (O 2 2- ), then oxygen = -1 In H 2 O: O=-2 In OF 2 : O=+2 In Na 2 O 2 (sodium peroxide): O=-1 6. Most elements will be the charge on the periodic table. Zn = +2 7. The sum of the oxidation numbers in all compounds must be ZERO 7a. The sum of the oxidation numbers in polyatiomic ions must equal the charge on the ion In NaCl: (+1) + (-1) = 0 In CaCl 2 : (+2) + 2(-1) = 0 In Al 2 (SO 4 ) 3 : 2(+3)+3(-2) =0 In SO 4 2- : S + 4(-2) = -2 S must = +6
WHAT ARE. OXidation-REDuction Reactions (also called REDOX) OXIDATION REDUCTION Electrons are LOST (LEO = loss of electrons=oxidation) Charge goes up (The oxidation number on the atom increases) What Happens to the electrons? LEO the lion goes GER What Happens to the charge? Electrons are GAINED GER: Gain of Electrons=Reduction Charge goes down (the oxidation number on the Atom decreases) Ex: Mg+Cl 2 MgCl 2 Oxidation Half reaction is: Mg Mg 2+ + 2e - Ex: Hg 2+ + 2I - Hg + I 2 Oxidation Half reaction is: 2I - I 2 + 2e - NOTE: e - are on right side of arrow Write the reaction Include: Balanced # of atoms Balanced charge What happens to the electrons if you add the two half reactions? Ex: Mg+Cl 2 MgCl 2 Reduction Half reaction is: Cl 2 + 2e - 2Cl - Ex: Hg 2+ + 2I - Hg + I 2 Reduction Half reaction is: Hg 2+ + 2e - Hg NOTE: e - are on left side of arrow
WHAT ARE. OXidation-REDuction Reactions (also called REDOX) OXIDATION REDUCTION Also called the Reducing agent (the thing that gets oxidized is doing the reducing so..) Oxidizing Agent (the thing that gets reduced is doing the oxidizing so..) ELECTRONS LOST ELECTRONS GAINED EVERYTHING IS ELECTRICALLY NEUTRAL!! ALL CHARGES MUST BALANCE (AND CANCEL)! Identifying redox reactions 1. Assign oxidation numbers to both sides of the equation 2. Look to see if oxidation numbers change. If they do = redox (if not then not) HINTS: Single replacement reactions are REDOX Double replacement are NOT redox
How to Balance a Redox reaction: 1. Assign oxidation # s to all of the atoms. 2. Identify which are oxidized & which are reduced. 3. Use one bracketing line to to connect the atoms that undergo oxidation & another to connect those that undergo reduction. 4. Make the total increase in oxidation # equal to the total decrease in oxidation # by using the appropriate coefficient. 5. Make sure that the equation is balanced for both atoms & charges. Ex. -3 reduction 2 x (-3) = -6 +3-2 +2-2 0 +4-2 Fe 2 O 3 (s) + 3 2 3 CO(g) Fe(s) + CO 2 (g) +2 oxidation 3 x (+2) = +6
ELECTROCHEMISTY IS an application of REDOX reactions Energy is produced or applied from REDOX reactions Two different ways ELECTROCHEMICAL Chemicals PRODUCE electricity (I.e., metals & solutions) HALF CELL /Half Reaction:»Show only 1/2 half of the reaction Either oxidation rxn OR reduction rxn» BOTH mass and charge are conserved!! (equal on both sides) Strip of metal Al ELECTOLYTIC Electricity FORCES a chemical change (apply the electricity) CELL POTENTIALS:»Voltages that are produced by each cell If E 0 = 0.00, have DEAD BATTERY (equilibrium) Solution that Contains that Metal ions Al 3+
ELECTROCHEMISTY Spontaneous Reactions: USE TABLE J TO PREDICT The neutral metal must be higher than the metal ion with which it s reacting Zn + Cu 2+ Cu + Zn 2+ On table J, Zn is higher than Cu = spontaneous Cu + Mg 2+ Cu 2+ + Mg: NOT SPONTANEOUS, Cu not higher than Mg NEEDS BATTERY! Electrons naturally flow from higher metal to lower metal
ELECTROCHEMICAL CELL (aka: Voltaic cells or Galvonic cells) Chemical rxn produces electricity Chemistry electricity SPONTANEOUS: - G i.e. making a battery + volts and +E 0 How do we tell if it is spontaneous? Metals as electrodes Zn, Cu Electrons flow from higher metal to lower metal on TABLE J Zn higher than Cu ANODE: oxidation occurs (Zn) Negative (-) Zn 0 gives up electrons Salt Bridge (U-tube) Salt Bridge: ions migrate in both directions CATHODE: reduction occurs (Cu) Positive (+) Cu +2 gains electrons Zn mass Ion Concentration AN OX Anode = oxidation RED CAT Reduction = cathode Decrease mass /Decrease e - Increase mass/increase e - Mass of Cu Ion Concentration
ELECTROLYTIC CELL (NEEDS A BATTERY) Apply electrical energy to make a chemical rxn happen. Need a battery NOT SPONTANEOUS!! + G (-) volts TWO MAJOR TYPES Electrolysis Separating a compound Into its elements EXAMPLE OF THE TWO MAJOR TYPES: +1-2 0 0 ELECTROLYSIS: 2H 2 O 2H 2 + O 2 Electroplating Plating or coating an Element onto an object BRINE SOLUTION (NaCl): 2NaCl 2Na + Cl 2(g) + Cl 2(g) H 2(g) - - + Anode Oxidation O -2 O 2(g) + 2e - e - e - Cl - H + Anode OH - Na + Ox. 2Cl - Cl 2 + 2e - Cathode Red. 2H + +2e - H 2 Battery Cathode Red. 2H + +2e - H 2
ELECTROPLATING (2 nd type) ELECTROLYTIC CELL (NEEDS A BATTERY) FOR ALL ELECTROLYTIC CELLS RED CAT AN OX Reduction Oxidation @ Cathode @ Anode Neg. term (-) Pos. term. (+) Metal used to plate- Always anode (oxidation) Always attached to positive term. Object to be plated- Always cathode (reduction) Always attached to negative term.
Similarities & Differences: Cu + Pb +2 ==> not spontaneous. Zn+ Ag+ ==> Spontaneous! Table J Metal is Higher than the ion (compound) Electrochemical/ Voltaic Cell Chemicals producing electricity- EXO (IS Spontaneous) + volts Summary: What s going on? Electrolytic Cell Electricity forcing a chemical reaction to occur Endo (NOT spontaneous)
Charge: Positive or Negative? Cathode + lower on Table J Anode - Higher on Table J Cathode -Attached to - Anode + Attached To + Reduction Oxidation Reduction Oxidation An Ox & Red Cat IS a Battery! IS a battery or NEEDS a battery? NEED a Battery!
e- flow anode cathode Salt bridge- allows IONS to migrate in both directions e- flow out- in + opposites attract object to be plated is ALWAYS attached to the negative!!!
Also known as Voltaic Spontaneousproduces electricity Makes a battery Electron flow predicted by Table J. --From higher to lower metal. Anode is the (-) terminal Cathode is the (+) terminal Both involve the transfer and movement of e- to make something happen LEO & GER Oxidation occurs at the anode An Ox Reduction occurs at the cathode Red Cat NOT Spontaneous Needs a battery to force e- to flow in wanted direction Anode is connected to (+) terminal of battery and is NOW (+) electrode Cathode is connected to (-) terminal of battery and is now (-) electrode Used for : Electrolysis-separating compounds into elements Electroplating coating items