Galvanic Cell

Electrochemistry : 10.1 Galvanic Cell

Electrochemistry

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  Study of relationship between chemical change & electric work

  
  
  
  
  Oxidation
• Loss of electron by species accompanied by an increase in oxidation number

  
  

  Ex:
  
  
  
  
  
  
  

  
  

  Reduction
  
  
• Gain electron by a species accompanied by a decrease number of oxidation
  Ex:

      Redox reaction

• Process which there are net movement of electrons from one reactant to another
• Also called oxidation – reduction process.
• Oxidation and reduction occur at the same time.
  
  
  Ex:

Oxidizing agent

• Substance that accepts electron in redox reaction and undergoes decrease number of oxidation.
  
  
  Ex:
  
  
  
  
  
  
  
  
  
  oxidising agent is Fe₂O₃
  
  

  
  
  Reducing agent
• Substance that donate electron in redox reaction and undergoes an increase oxidation number.
  
  
  Ex:

CO is an reducing agent

    key point      

• Oxidation always accompany by reduction
• Oxidizing agent reduced
•  reducing agent oxidise

ELECTROCHEMICAL CELL

There are two type :

• Voltaic cell
• Electrolytic cell

VOLTAIC CELL

• Use spontaneous reaction to generate electric energy
• System does work on surrounding
• All batteries contains voltaic cell

ELECTROLYTIC CELL

• Use electrical energy to drive non-spontaneous reaction
• Surrounding do work on system
• Ex: electroplating & recovering metal from ores

ELECTRODES

• Object that conduct electricity between cell and surrounding
• 2 electrode (anode & cathode) are dipped into electrolyte
• ACTIVE ELECTRODES
• Involve in half reaction
• Ex: zinc (Zn), copper (Cu), Iron (Fe)

INACTIVE ELECTRODE

• If no reactant or product can be uses as electrode
• Ex: graphite (C), Platinum (Pt)
• The electrolyte solution contain all species involved in the half – reaction 

ANODE & CATHODE

• Oxidation half reaction takes place at cathode
• Electron given up by substance being oxidised (reducing agent) and leave the cell at anode
• Reducing half reaction takes place at cathode
• Electron are taken up by substance being reduced (oxidizing agent) & enters the cell at cathode

ELECTROLYTE

• Mixture of ions (usually in aqueos solution) that are involved in reaction or that carry charge.

ESSENTIAL IDEA OF VOLTAIC CELL

• Component of each half reaction (half – cell) are placed in a seperate container
• The two half cell are joined by circuit which consist a wire and a salt bridge

OXIDATION HALF CELL

• Metal bar (anode) is immersd in Zn²⁺ electrolyte 
  Ex: ZnSO₄
• Zn is reactant in oxidation half reaction
• Zn conduct electricity and release electron out of its half cell
  
  
  
  
  
• Mass of Zn electrode decrease

REDUCTION HALF CELL

• A Cu metal bar (cathode) is immersed in a Zn²⁺ electrolyte
• Ex: CuSO₄ solution
• Cu is a product in reduction half reaction
• Cu conducts electricity and released electron into itself.
  
  
  
  
  
• Mass of copper electrode increase

RELATIVE CHARGE ON ELECTRODES

• Anode: 

  

  
  
• Electron flow from anode to cathode through wire
• Cathode 
  
  
  
  
  
• Electron are continously generated at anode * consume at cathode

SALT BRIDGE

• Contains a solution of non-reacting ion such as KCl,KNO_3, Na₂SO₄ acts as liquid wire (allowing ions to flow & complete the circuit.

  
  

  HOW DOES THE CELL MAINTAIN ITS ELECTRICAL NEUTRALITY 
  
  

Anode Half Cell	Cathode Half Cell
Zn 2+ ions enters the solution causing the excess of positive charge.	Cu2+ ions leave the solution causing the excess of negative charge
Cl- ions from the salt bridge move into anode (Zn) half cell	K+ ions from salt bridge move into cathode Cu half cell

  

CELL NOTATION

• Ex : Zn(s) l Zn²⁺ (aq) l l Cu²⁺ (aq) l Cu(s)
• Anode left, cathode right
• 'l' represent phase boundary
• Use ',' for component that are in the same phase
  Graphite l I^- (aq) l I(s) l l H⁺ (aq), MnO₄^- (aq), Mn²⁺ (aq) l graphite
• Sometimes we specify the concentration of dissolved component
  Zn(s) l Zn²⁺ (1 M) l Cu²⁺ (1 M) l Cu(s)
• Electrode appear ait far right and left notation.
• Half-cell component usually appear in the same order as in the half-reaction
  
  
  
  
      Cu ²⁺ (aq) + 2^e- ------ Cu (s)        (reducing)
                      Zn(s) ------ Zn²⁺ (aq) + 2e^- (oxidizing)
  
  
  Cu ²⁺ (aq) + Zn(s) ------ Cu (s) + Zn²⁺ (aq) (overall)

Cell notation: the coefficient is not involved.

 SPONTANEOUS REACTION

• occurs as the result of different ability of metal to give up their electron to flow through the circuit.

 

CELL POTENTIAL (E_Cell)

• different in electrical potential of electrodes
• also called voltage or electromotive force (e.m.f)

 

E_cell > 0

• sponteneous reaction
  • The more positive E_cell
  • The more work the can do
• The further the reaction proceed to right

 

E_cell < 0

• Non spontaneous cell reaction

 

E_Cell = 0

• The reaction has reach equilibirum
• The cell can do no more work.

 

SI UNIT CELL POTENTIAL

• unit = volt (V)
• 1V = 1J/C

C = coulumb (SI unit of electrical charge)

 

STANDARD CELL POTENTIAL (E⁰_cell)

• Different in electrical potential of electrodes measured at a specified temperature (usually 298k) with all components in their standard states.
• standard state 
• 1 atm for gaseous
• 1 M for solution
• Pure solid for electrodes

STANDARD ELECTRODE (HALF CELL) POTENTIAL (E⁰_half-cell)

• potential associate with a given half-reaction (electrode compartment) when all component are in their standard states.

E⁰_{half cell =} E⁰_anode or E⁰_cathode

• also call standard reduction potential.
• example : 

                         Zn²⁺(aq) + 2e^- ------ Zn(s) E⁰_zinc (E⁰_anode)

                        Cu²⁺(aq) + 2e^- ------ Cu(s) E⁰_copper(E⁰_cathode)

                    Zn(s) + Cu²⁺ (aq) ------ Zn²⁺ (aq) + Cu(s)

*changing the balancing coefficients of a half-reaction does not change E⁰value because electrode potential are intensive properties –does not depend on amount

E⁰_cell AND E⁰_{half cell}

• Example : 
  

Half cell reation

                            Cu²⁺(aq) + 2e^- ------ Cu(s)

                         Zn(s) + Cu²⁺(aq) ------ Zn²⁺(aq) + Cu(s)

                         Zn(s) + Cu²⁺(aq) ------ Zn²⁺(aq) + 2e^-

 

• E⁰_cell = E⁰_cathode – E⁰_anode

 

STANDARD HYDROGEN ELECTRODE

• specially prepared platinum electrode immersed in a 1M aqueous solution of a strong acid,H⁺(aq) or H₃O⁺(aq), through which H₂ gas at 1 atm is bubled

DETERMINING E⁰_{half cell}

• Use standard hydrogen electrode (SHE)
• standard reference half-cell
  
  
  2H⁺ (aq, 1m) + 2e^- ------ H₂ (g,1atm) E⁰ ref= OV

         Zn²⁺ (aq, 1m) + 2e^- ------ Zn (s, 1atm) E zinc= ?

Zn(s) │ Zn²⁺ (1M) ││ H⁺ (1M) │ H₂ (1atm) │ Pt (s)

E⁰ cell  = E⁰ cathode – E⁰ anode 

              = E⁰ ref – E⁰ zinc 

0.76 V  = 0.00 V – E⁰ zinc 

E⁰ zinc  = -0.76 V

Pt(s) │ H₂ (1atm) │ H⁺ (1M) │ Cu²⁺ (1M) │ Cu (s) 

  

Anode : H₂ (1atm) 2H⁺ (1M) +2e^-

Cathode : 2e^- + Cu²⁺ (1M)     Cu (s) 

H₂ (1atm) + Cu²⁺ (1M)     Cu (s) + 2H⁺ (1M)

 

RELATIVE STRENGH OF OXIDIZING AND REDUCTING AGENT

• Example:

Cu²⁺(aq) + 2e^- ------ Cu (s) E⁰ = 0.34 V

2H⁺ (aq) + 2e^- ------ H₂ (g) E⁰ = 0.00 V

Zn²⁺ (aq) + 2e^- ------ Zn (s) E⁰ = -0.76 V

• The more positive E⁰ value, the more tendency to be reduced
• Strength of oxidizing agent (reactant)
  Cu²⁺ > H⁺ > Zn²⁺
• Strength of reducing agent (product)
  Zn > H₂ > Cu
  
  

STANDARD REDUCTION POTENTIAL 

• All value are relative to hydrogen electrode
• Strength of oxidizing agent – increase up
• Strength of reducing agent – increase down
• Half-cell component usually appear in the same order as in the half-reaction
                                               Cu²⁺ (aq) + 2e^- ------ Cu (s)                        (reducing)
                                                                Zn (s) ------ Zn²⁺ (aq) + 2e^-        (oxidizing)
                                           Cu²⁺ (aq) + Zn (s) ------ Zn²⁺ (aq) + Cu(s)    (overall)

  
  

Cell notation : the coefficient is not involve

 

WRITING SPONTANEOUS REDOX REACTION

Zn(s) + Cu²⁺ (aq) ------ Zn²⁺ (aq) + Cu(s)

 

Zn     - stronger reducing agent

Cu²⁺ - stronger oxidizing agent

Zn²⁺  - weaker (0.9)

Cu     - weaker (0.9) 

• stronger oxidizing agent : E⁰ larger (more positive)
• stronger reducing agent : E⁰ smaller (more negative)

Cu²⁺ (aq) + 2e^- Cu(s) E⁰: 0.34 V

Zn²⁺ (aq) + 2e^- Zn(s) E⁰: - 0.76 V

• How to determine anode (oxidation) and cathode (reduction) for spontaneous reaction, E⁰_cell > 0 ?
• Strong reducing agent = E⁰ larger (positive)

                                                 = E⁰ cathode (reduction)

Reduction: Cu²⁺ (aq) + 2e^- ------ Cu (s) E⁰ =0.34V

Oxidation :                  Zn (s) ------ Zn²⁺ (aq) + e^- E = - 0.76V

• Stronger reducing agent     = E⁰ smaller 

                                                         = E⁰ anode (oxidation)

PREDICTING SPONTANEOUS REDOX REACTION USING DIAGONAL RULE

• Under standard – state condition, any species on the left of a given half-cell reaction will react spontaneously with a species that appear on the-right of any half-cell reaction located below it,
• diagonal rule !!! 

Cu²⁺ (aq) + 2e^- ------ Cu (s)     E⁰ = 0.34V

 Zn²⁺ (aq) + 2e^- ------ Zn (s)      E⁰ = - 0.76V