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Electrochemistry Chapter 9 Kulkarni. Electrochemistry: Study of chemical processes that involve the exchange of electrons. Reduction : Gain of electrons. The substance that gains electrons is said to be reduced Cl 2 + 2e - 2Cl - The process is reduction
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Electrochemistry: Study of chemical processes that involve the exchange of electrons.
Reduction: Gain of electrons. The substance that gains electrons is said to be reduced Cl2+ 2e- 2Cl- The process is reduction Chlorine is reduced to chloride.
Oxidation: Loss of electrons. The substance that loses electrons is said to be oxidized Ca 2e- + Ca+2 The process is oxidation Calcium is oxidized to calcium ion.
Redox Oxidation cannot occur without reduction and vice versa. That’s why they are called “redox” reactions! Ca + Cl2 CaCl2
“LEO the lion says GER” Loss Gain Electrons Electrons Oxidation Reduction
“OIL RIG” Oxidation Is Loss Reduction Is Gain
Oxidation Number • Used to keep track of electrons • Shows the general distribution of electrons NOT absolute charge!
RULES! • Uncombined elements: Oxid. # = zero. Ex. Cu, Cl2, P
Binary ionic compounds: Oxid. # = ionic charge Ex. CaCl2: Ca+2 o.#. Is +2 and Cl-1 o.#. is -1
3. Metals in compounds: Group 1 +1 Group 2 +2 Aluminum +3 Zn and Cd +2 Silver +1 Transitions ionic charge
Fluorine: -1 always • Hydrogen: • Usually +1 • With metals -1 Ex. NaH, CaH2 (hydrides)
Oxygen Usually -2 In peroxides -1 Ex: H2O2, group I oxides (Li2O2) With Fluorine +2 (OF2)
7. The sum of oxidation numbers in a neutral compound is zero. CaCl2: (+2) + 2(-1) = 0
Covalent Compounds The more electronegative element is assigned all of the electrons as if it were ionic. Ex. NO2 – Oxygens are each -2 so the nitrogen is +4. (x) + 2(-2) = 0, x = +4
The sum of oxidation numbers for a polyatomic ion equals the charge on the polyatomic ion. PO4-3: (+5) + 4(-2) = -3
MnO2 + 4HCl MnCl2 + 2H2O +Cl2 Redox Reactions In redox reactions, one or more atoms change oxidation numbers. +4 -2 +1 -1 +2 -1 +1 -2 0 Mn+4 is reduced Chloride (Cl-) is oxidized
Half-Reactions Two parts of a redox reaction written separately Mg + Cl2 MgCl2 • Mg Mg+2 + 2e- • Cl2 + 2e- 2Cl- Oxid. Red.
Reducing Agent: Helps another element become reduced. (it is oxidized in the process) Oxidizing Agent: Helps another element become oxidized. (it is reduced in the process)
NOTE: When discussing what is reduced/oxidized or what is the oxidizing/reducing agents we are always referring to REACTANTS with their OXIDATION NUMBER!
Balancing Redox – ½ rxn method • Write down the chemical equation. • Determine all oxidation numbers. • Write down balanced ½ rxns.
Find a common multiple for the number of electrons in each ½ rxn such that: # e-’s gained = # e-’s lost • Move coefficients from the ½ rxn to the complete rxn.
Complete the balancing process for the remaining elements. NOTE: If not all of an element is oxidized or reduced you may not be able to use the coefficient.
Balancing in Acidic Soln: • Assign Oxidation #s • Write the ½ reactions as before
Balance the charge with the addition of e-’s Find a common multiple for the number of electrons in each ½ rxn such that: # e-’s gained = # e-’s lost
Place coefficients into the original reaction. • Balance the oxygens with water. • Balance the hydrogens with H+. • Cancel of necessary
Balancing in Basic Soln: • Balance as if it were Acidic Solution. • Find a common multiple for the number of electrons in each ½ rxn such that: # e-’s gained = # e-’s lost
Neutralize the H+ with OH-. Add the same amount of OH- to the opposite side of the equation.
Cancel as needed H2O’s Hint: H+ + OH- = H2O That is why we neutralize the H+ in step number 2.
PRACTICE! In Class WKS
Electrochemical Cells System of two cells in which chemical energy is converted to electrical energy or vice versa.
1. Galvanic/Voltaic Cell SPONTANEOUS electro-chemical cell. The transfer of electrons produces energy Ex. batteries
2. Electrolytic Cells NON-SPONTANEOUS – energy must be added to bring about a chemical change. Ex. Electroplating
3. Components • ELECTRODES – conductor that establishes electrical contact with a nonmetallic, electrolytic part of the cell.
ANODE – where oxidation takes place. CATHODE – where reduction takes place “An ox and a red cat”
Figure 18.1: Schematic for separating the oxidizing and reducing agents.
POROUS BARRIER/ SALT BRIDGE Prevents half-reactions from mixing but allows ions to move freely to balance charges in both cells.
Figure 18.5: Schematic of a galvanic cell. Click on picture to go to http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf