Electro-chemistry

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CHAPTER 15
Electro-chemistry
15.4 Electrochemical Cells
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The lemon battery
A chemical reaction between magnesium and the
citric acid generates electrons Now we know it
is an oxidation reaction
Electron flow
Magnesium
Citric acid
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The lemon battery
A chemical reaction between magnesium and the
citric acid generates electrons Now we know it
is an oxidation reaction This means that
somewhere else, a reduction reaction must also
take place to absorb the electrons
LED
Electron flow
Copper
Citric acid
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LED
Oxidation and reduction reactions are separated
Electron flow
Electron flow
Copper
Magnesium
Citric acid
electrochemical cell a device in which redox
reactions take place.
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Three main components of an electrochemical
cell 1) Two electrodes 2) The electrolyte 3) A
conducting path that connects the electrodes
externally
electrochemical cell a device in which redox
reactions take place.
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An electrochemical cell
Three main components of an electrochemical
cell 1) Two electrodes
The anode electrode is where the oxidation
occurs The cathode electrode is where the
reduction occurs
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An electrochemical cell
Three main components of an electrochemical
cell 1) Two electrodes 2) The electrolyte
Electrodes are immersed in a conductive
solution. This solution, or electrolyte,
contains free ions.
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An electrochemical cell
Three main components of an electrochemical
cell 1) Two electrodes 2) The electrolyte 3) A
conducting path that connects the electrodes
externally
There must be a connection between oxidation and
reduction Electrons must be able to flow from the
anode to the cathode
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There are two types of electrochemical
cells Voltaic cells Spontaneous chemical
reactions at the electrodes generate electrical
current(A battery is a voltaic
cell) Electrolytic cells An external electric
current drives nonspontaneous reactions between
the electrodes and the electrolyte
(or galvanic cells)
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There are two types of electrochemical
cells Voltaic cells Spontaneous chemical
reactions at the electrodes generate electrical
current(A battery is a voltaic
cell) Electrolytic cells An external electric
current drives nonspontaneous reactions between
the electrodes and the electrolyte
(or galvanic cells)
In both cases, the cells are constructed with two
half-cells
Anodewhere oxidation happens
Cathodewhere reduction happens
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A voltaic cell
The flow of electrical current is completed
inside the battery Negative ions move from the
cathode to the anode Positive ions move from the
anode to the cathode to maintain electroneutrality
inside the battery
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Chemistry in a voltaic cell
Consider the following setup
Mg loses electrons more easily than Cu so
electrons would flow from the Mg side
e flow
e flow
Oxidation (anode) Mg(s) ? Mg2(aq)
2e Reduction (cathode) Cu2(aq) 2e ? Cu(s)
Mg(s)
Cu(s)
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Chemistry in a voltaic cell
Consider the following setup
e flow
e flow
Why does the flow of electrons only happen for a
short time?
Mg(s)
Cu(s)
LED no longer lights up
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Chemistry in a voltaic cell
Consider the following setup
Oxidation (anode) Mg(s) ? Mg2(aq) 2e
e flow
e flow
Mg(s)
Cu(s)
Mg2 ions have nowhere to go Mg2 and e are
attracted to each other The reaction stops as
soon as it starts
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Chemistry in a voltaic cell
Oxidation (anode) Mg(s) ? Mg2(aq) 2e
Reduction (cathode) Cu2(aq) 2e ? Cu(s)
e flow
e flow
Mg(s)
Cu(s)
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Chemistry in a voltaic cell
The solution provide a path for positive ions to
move from the anode half-cell to the cathode
half-cell
e flow
e flow
Mg(s)
Cu(s)
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salt bridge an electrical connection between the
oxidation and the reduction half-cells of an
electrochemical cell.
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Notation
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Electromotive force
Electrons that flow from the anode to the cathode
are pushed by the electromotive force
electromotive force (emf) the difference in the
electrical potential between the anode and the
cathode of an electrochemical cell.
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Electromotive force
Based on the emf, we can calculate the amount of
energy that the cell can provide
cell emf (the driving force of the cell reaction)
electrical energy output
moles of electrons
Faraday constant 96,500 C/mole
electromotive force (emf) the difference in the
electrical potential between the anode and the
cathode of an electrochemical cell.
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Electromotive force
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Electromotive force
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Electromotive force
Asked Wcell, energy from this cell with 2.43 g
of Mg Given The cell reaction, Ecell 2.71
V, and the amount of reactant Relationships H
alf-reactions Solve
Mg(s) ? Mg2(aq) 2e Cu2(aq) 2e ? Cu(s)
From the oxidation half-reaction, 1 mole Mg 2
moles e So n 0.0200 moles e
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Electromotive force
Asked Wcell, energy from this cell with 2.43 g
of Mg Given The cell reaction, Ecell 2.71
V, and the amount of reactant Relationships H
alf-reactions Solve Answer With
2.43 g of Mg, this cell will release 5,230 J of
energy.
Mg(s) ? Mg2(aq) 2e Cu2(aq) 2e ? Cu(s)
From the oxidation half-reaction, 1 mole Mg 2
moles e So n 0.0200 moles e
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The total cell voltage (Ecell) is a combination
of the potential at each half-cell
Ecell reduction potential oxidation
potential
For the oxidation half-reaction Mg(s) ? Mg2(aq)
2e Eoxidation 2.37 V For the reduction
half-reaction Mg2(aq) 2e ? Mg(s) Ereduction
2.37 V
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The total cell voltage (Ecell) is a combination
of the potential at each half-cell
Ecell reduction potential oxidation
potential
For the oxidation half-reaction Mg(s) ? Mg2(aq)
2e Eoxidation 2.37 V For the reduction
half-reaction Mg2(aq) 2e ? Mg(s) Ereduction
2.37 V
Scientists decided to only keep track of
Ereduction
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Standard reduction potentials
Mg2(aq) 2e ? Mg(s) Ereduction 2.37 V
The cell voltage under standard conditions is
called the standard reduction potential, Eocell
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Standard reduction potentials
We cant measure the potential of single
electrodes We need a reference half-cell!
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Standard reduction potentials
We cant measure the potential of single
electrodes We need a reference half-cell!
Use this reaction as the reference where
Ereduction 0 V
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Standard reduction potentials
The potential of all other cells is measured with
respect to the reference cell
Reference half-cell
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Using the standard reduction potentials,
calculate the cell voltage (Eocell) of the cell
Zn(s) Zn2(aq) Cu2(aq) Cu(s).
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Using the standard reduction potentials,
calculate the cell voltage (Eocell) of the cell
Zn(s) Zn2(aq) Cu2(aq) Cu(s).
Asked Eocell of the reaction Zn(s) Cu2(aq) ?
Zn2(aq) Cu(s) Given The cell reaction,
standard reduction potentials Relationships Eoce
ll Eoreduction Eooxidation
Look up Table 15.4
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Using the standard reduction potentials,
calculate the cell voltage (Eocell) of the cell
Zn(s) Zn2(aq) Cu2(aq) Cu(s).
Asked Eocell of the reaction Zn(s) Cu2(aq) ?
Zn2(aq) Cu(s) Given The cell reaction,
standard reduction potentials Relationships Eoce
ll Eoreduction Eooxidation
Zn(s) ? Zn2(aq) 2e Eooxidation 0.76
V Cu2(aq) 2e ? Cu(s) Eoreduction 0.34 V
reverse reaction
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Using the standard reduction potentials,
calculate the cell voltage (Eocell) of the cell
Zn(s) Zn2(aq) Cu2(aq) Cu(s).
Asked Eocell of the reaction Zn(s) Cu2(aq) ?
Zn2(aq) Cu(s) Given The cell reaction,
standard reduction potentials Relationships Eoce
ll Eoreduction Eooxidation
Zn(s) ? Zn2(aq) 2e Eooxidation 0.76
V Cu2(aq) 2e ? Cu(s) Eoreduction 0.34 V
Solve Eocell Eoreduction Eooxidation Eocel
l 0.34 V 0.76 V 1.10 V Answer The cell
voltage is 1.10 V
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Reaction spontaneity
The sign of Eocell also tells you whether the
reaction is spontaneous or nonspontaneous
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Determine whether the reaction Zn(s) Ni2(aq) ?
Zn2(aq) Ni(s) is spontaneous under standard
conditions.
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Determine whether the reaction Zn(s) Ni2(aq) ?
Zn2(aq) Ni(s) is spontaneous under standard
conditions.
Asked The spontaneity of the reaction under
standard conditions Given The reaction and the
table of standard reduction potentials Relationsh
ips Eocell Eoreduction Eooxidation The
half-cells and their half-cell potentials
Zn(s) ? Zn2(aq) 2e Eooxidation 0.76
V Ni2(aq) 2e ? Ni(s) Eoreduction 0.23 V
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Determine whether the reaction Zn(s) Ni2(aq) ?
Zn2(aq) Ni(s) is spontaneous under standard
conditions.
Asked The spontaneity of the reaction under
standard conditions Given The reaction and the
table of standard reduction potentials Relationsh
ips Eocell Eoreduction Eooxidation The
half-cells and their half-cell potentials
Zn(s) ? Zn2(aq) 2e Eooxidation 0.76
V Ni2(aq) 2e ? Ni(s) Eoreduction 0.23 V
Solve Eocell 0.23 V 0.76 V 0.53 V gt
0 Answer Since the total cell voltage is a
positive number the reaction is spontaneous and
proceeds as indicated.
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Three main components of an electrochemical
cell 1) Two electrodes 2) The electrolyte
3) A conducting path that connects the electrodes
externally (like a salt bridge)
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cell emf (the driving force of the cell reaction)
electrical energy output
moles of electrons
Faraday constant 96,500 C/mole
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Nernst equation
Batteries (a voltaic cell) run out of energy
Nernst equation
moles of electrons
standard cell potential
Under standard conditions Q 1 so that Ecell
Eocell
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Nernst equation
Batteries (a voltaic cell) run out of energy
Nernst equation
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Electrolytic cells
The nonspontaneous reaction that is induced is
called electrolysis
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Electrolytic cells
When we use a battery it acts as a voltaic cell
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Electrolytic cells
When we recharge a battery it acts as an
electrolytic cell
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Electrolytic cells
Decomposition of water
2H2O(l) ? 2H2(g) O2(g)
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Electrolytic cells
Decomposition of water
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Electrolytic cells
Decomposition of water
Reduction 2H2O(l) 2e ? H2(g) 2OH(aq) Eored
0.83 V Oxidation 2H2O ? O2(g) 4H(aq)
4e Eoox 1.23 V
Eocell 2.06 V
nonspontaneous reaction
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Electrolytic cells
Decomposition of water
Hydrogen (H2) is being considered as a source of
energy for the future. The production of
hydrogen from water is the cleanest but also the
most expensive way of producing hydrogen.
Eocell 2.06 V
nonspontaneous reaction