Battery, Fuel cell, Galvanic cell, Voltaic cell.

School Science Lessons
2024-06-11

Batteries
(UNPh33.1) Table of contents

Batteries
33.1.1 Anode and cathode
33.1.2 Battery, batteries
33.1.3 Cells
33.1.4 Conventional current
33.1.3.3 Dry cell battery, Leclanché cells
33.1.6 Electrolytes
33.1.7 Primary cells non-rechargeable batteries
33.1.8 Secondary cells , rechargeable batteries
33.1.9 Weston Standard Cell

33.1.1 Anode and cathode
An anode is a positive electrode or terminal.
anode (Greek anodos way up)
A cathode is a negative electrode or terminal
cathode (Greek kathodos way down)
An electrode is a device to facilitate the conduction of electricity into or out of something.
electrode (Greek hodos way).
Electrodes are conductors through which electric current enters or leaves the electrolyte.
Cations are positive ions attracted to the negative cathode.
Anions are negative ions anions are attracted to the positive anode.
In both electrochemical cells and electrolytic cells:
cations (positive ions), Na+, move towards the cathode and anions (negative ions) e.g. Cl-, move towards the anode.
Battery terminals are labelled positive, +ve, and negative, -ve, according to voltage.
This is because the positive terminal remains at a positive voltage relative to the other terminal.
Oxidation, the loss of electrons from the ions in solution to the electrode, occurs at the anode.
Reduction, the gain of electrons from the electrode to the ions in solution, occurs at the cathode.
If the electrolyte is a salt consisting of a metal and non-metal, the metal precipitates at the cathode and the non-metal precipitates at the anode.
In both electrochemical cells and electrolytic cells, electrons carry current through the external circuit and ions carry current through the solutions.

33.1.2 Battery, batteries
33.1.2.1 Batteries , origin, use of the term "battery"
33.1.2.2 Battery comparison , common rechargeable batteries
33.1.2.3 Battery disposal
33.1.2.4 Battery uses , types of batteries
33.1.2.5 C-rates of batteries
33.1.2.6 Discharge curves of batteries
33.1.2.7 Energy density of batteries
Experiments
33.1.3.3 Dry cell battery, Leclanché cells
33.4.53.1 Zinc chloride battery

33.1.3 Cells
An electrochemical cell or "cell" is a vessel containing electrodes and electrolyte for electricity generation or electrolysis.
33.1.3.1 Aluminium foil chocolate wrapper cell in the mouth
33.4.51 Electricity from two coins
33.1.3.3 Dry cell battery, Leclanché cells
33.1.3.6 Magnesium / copper cell
33.3.14 Magnesium pencil sharpener electrodes
33.1.3.8 Potato cell
33.3.5 Prepare a simple chemical rectifier
33.3.1 Simple electric cell
33.1.3.11 Simple galvanic cell
33.1.3.12 Test a simple cell with different metals

33.1.4 Conventional current
If "current" is the conventional positive current, in metal wires, the current is carried by electrons moving in the direction opposite to the current.
In conventional current, the cathode of a battery is the terminal where current flows out and the anode is the terminal where current flows in.
Conventional current passes from the positive terminal, through a circuit, to the negative terminal of the battery, from cathode to anode.
In a discharging battery or a galvanic cell, the cathode is the positive terminal since that is where the conventional current flows out of the device.

33.1.3.3 Dry cell battery, Leclanché cells
Experiments Bring a dead battery to life: 33.1.5.1
Dry cell capacity: 33.1.5.2
Dry cell in electric circuit: 33.1.5.3
Dry cells in electric circuit, cells in series and parallel: 33.1.5.4
Dry cell, electric torch (flashlight) battery: 33.1.5.5
Zinc-carbon cell, dry cell torch battery, (flashlight battery): 33.1.5.6

33.1.6 Electrolytes
Electrolytes may be acids, bases or salts dissolved in water.
Electrolytes are liquids containing ions that may be decomposed by electrolysis, e.g. the solution in a car battery.
Electrolytes form ions when dissolves in water.
Living cells, blood and other tissues contain electrolytes.
The O.R.S. (oral rehydration salts), taken to treat diarrhoea and traveller's gastro-enteritis in children and adults, contains a balanced mix of glucose and electrolytes.

33.1.7 Primary Cells, non-rechargeable batteries
Primary cells are voltaic cells where the electrical energy comes from an irreversible chemical reactions of the components, and so cannot be recharged.
A primary cell is a voltaic cell in which the chemical reaction that produces the EMF cannot be reversed properly, so the cell cannot be recharged by electrical means.
These cells are light, small and easy to replace.
A primary battery is disposable, i.e. it can be discharged only once.
Primary cell, non-rechargeable batteries, include the following:
33.1.7.1 Daniell cell
33.1.7.2 Galvanic cell, Voltaic cell
33.1.7.3 Galvanic cell, Voltaic cell, with salt bridge
33.1.7.4 Leclanché cell
33.1.7.5 Lemon cell
33.1.7.6 Magnesium / copper battery
33.1.7.7 Mercuric oxide battery
33.3.14 Magnesium pencil sharpener electrodes
33.1.7.9 Silver-oxide battery (a button cell)
33.3.1 Simple electric cell
3.1.7.10 Test simple electric cell with copper and zinc in dilute sulfuric acid

33.1.8 Secondary cells, rechargeable batteries
Secondary cells are rechargeable batteries, accumulators, e.g. lead cell accumulator motor vehicle battery, are voltaic cells which can be charged and discharged with an electric current.
A secondary cell, accumulator, is a voltaic cell that can be recharged after discharged.
However, they are large and heavy and contain a lot of dangerous liquid electrolyte, e.g. Ni-Cd storage batteries with potassium hydroxide electrolyte can be loaded again up to1000 times.
In a circuit where ions conduct electricity, the positive and negative ions move at the same time as an ion electric current to positive or negative poles acted on by the electric field force.
A secondary battery is rechargeable, it can be discharged and later recharged, so it can be used many times, e.g. the lead cell accumulator motor vehicles.
The main types of rechargeable batteries:
33.1.8.1 Alkaline battery
33.1.8.2 Sealed lead-acid battery, SLA battery
33.6.16 Calomel, (Hg2Cl2). half cell
33.1.8.3 Hydrogen /oxygen fuel cell
33.1.8.4 Lithium-ion, Li-ion battery
33.1.8.5 Nickel-cadmium battery, NiCad battery
33.1.8.6 Nickel-metal hydride, NiMH battery

33.1.8.1 Alkaline battery
1. The positive electrode is manganese dioxide.
The negative electrode is a zinc rod.
The electrolyte is potassium hydroxide, which does not take part in the electrochemical reaction, but remains, because equal amounts of OH- are consumed and produced.
Zn (s) + 2OH- (aq) --> ZnO (s) + H2O (l) + 2e- (e = 1.28 V)
2MnO2 (s) + H2O (l) + 2e- --> Mn2O3 (s) + 2OH- (aq) (e = +0.15 V)
Zn (s) + 2MnO2 (s) <--> ZnO (s) + Mn2O3 (s) (e = 1.43)
An alkaline primary cell battery is small, lightweight and durable.
They are inexpensive and relatively efficient so they are currently the most common battery type.
A unit of battery capacity for smaller cells is the milliamp-hour, one thousandth of an amp for one hour.

33.6.13 Rechargeable alkaline-manganese battery, RAM battery
Rechargeable alkaline-manganese battery, RAM battery, can be recharged for repeated use
Cathode paste + barium sulfate + catalyst to recombine hydrogen + zinc oxide
Anode zinc powder suspended in gel
RAM batteries are low cost, low self-discharge, prefers shallow cycling.
They have no memory effect, but a short life cycle.
They are used for portable emergency lighting, toys, portable radios, cassette and CD players, test instruments.

33.1.8.2 Sealed lead-acid battery, SLA battery
Sealed paste-type lead-acid batteries, "gel cells", have a semisolid electrolyte that eliminates some of the drawbacks of regular lead acid batteries.
They are more expensive, but they can be used in any position and require little maintenance, e.g. laptop computers.
The two types of sealed batteries are as follows:
1. The gel cell that a silica additive in its electrolyte solution to stiffen it or gel, so eliminating some of the problems with venting and spilling.
2. The Absorbed Glass Mat (AGM) type has the electrolyte suspended near the electrodes.
These batteries are sealed and do not need maintenance, so are suitable for remote or enclosed locations.
SLA batteries are low cost, low self-discharge, float-charging, but shallow charging better.
Used for emergency lighting, solar power systems, wheel chairs.

33.6.16 Calomel, (Hg2Cl2), half cell
A reference electrode consisting of a half cell with a mercury electrode in a solution of potassium chloride saturated with calomel, mercury (I) chloride, Hg2Cl2 or HgCl.
Standard electrode potential = -0.2415 V at 25oC.
HgCl (s) <--> Hg+ (aq) + Cl- (aq)
Hg+ (aq) + electron <--> Hg (s)
HgCl (s) + electron <--> Hg (s) + Cl- (aq)

33.1.8.3 Hydrogen / oxygen fuel cell
The fuel cell allows the reaction between hydrogen and oxygen to generate electricity.
The equipment is sold with a 2 V motor to generate power.
The reaction requires potassium hydroxide pellets and palladium (II) chloride.

33.1.8.4 Lithium-ion battery, Li-ion battery
Lithium batteries are energy dense and light weight, but relatively expensive.
Lithium cells are quickly becoming a popular battery choice.
Lithium ion is a newer, rechargeable lithium-based battery.
Lithium polymer is a lower cost, more stable version of the lithium ion cell.
+ve electrode: copper coated with carbon graphite
-ve electrode: aluminium coated with lithium-cobalt dioxide or other compounds
Electrolyte: Lithium salt, e.g. Lithium-phosphorus hexafluoride
Li-ion batteries have about twice the energy density of nickel batteries and they have no memory effect.
They must not be excessively overcharged or discharged.
They are used in laptop computers, and cell phones.
Lithium-ion batteries are very compact, low maintenance, low self-discharge, but must be charged with great care.
They are used for compact cell phones, notebook PCs, digital cameras, and any very small portable devices.
Lithium is the lightest metal and has the highest electrochemical potential, so it should provide a very high energy density.
However, it is highly reactive and could explode during recharging, so lithium compounds are used rather than lithium metal.
In 2013, a six year old girl died in Brisbane from lithium poisoning after she had swallowed a Li-ion battery button cell.
The medical authorities said there was no known antidote for lithium poisoning.
Authorities in Australia have since insisted on future child-safe packaging of button batteries.
Lithium metal and lithium ion/polymer batteries are prohibited in checked baggage for air travel.

33.1.8.5 Nickel-cadmium battery, NiCad battery
The types of battery include cylindrical cells like dry cells, button cells for miniature appliances, and battery packs for handheld radio transceivers and spacecraft.
+ve electrode: nickel / nickel hydroxide (nickel oxide hydroxide)
-ve electrode: cadmium / cadmium hydroxide
Electrolyte: potassium hydroxide
NiCad rechargeable batteries prefer deep cycling, good pulse capacity, but memory effect, high discharge rate.
Used in portable tools, appliances, model boats, data loggers, camcorders, portable transceivers, test equipment, electric toothbrushes, cordless power tools.
They have cost-effective energy storage and long working life.
During discharge:
At the cathode (+ve): nickel (IV) hydroxide + 2 electrons --> nickel (II) hydroxide (reduction).
At the anode (-ve): cadmium - 2 electrons --> cadmium (II) hydroxide (oxidation).
Recharging procedures
Nicads have a good flat discharge curve, but can potentially suffer from the "memory effect".
If consistently recharged before they are fully discharged, they "remember" the point of premature recharging and later runs down only to the capacity at last recharge.
The solution is to let the battery discharge almost completely before recharging, i.e. be deeply cycled.
Constant recharging after use for a short time may produce cause and change the form of cadmium crystals in the battery, causing lower electric release, lower voltage.
Nickel cadmium batteries can go through thousands of charging cycles.
Always follow local regulations when disposing of NiCad batteries, because they contain cadmium, an environmental pollutant.

33.1.8.6 Nickel-metal hydride battery, (NiMH, Ni-MH)
+ve electrode: nickel / nickel hydroxide (nickel oxide hydroxide)
-ve electrode: lanthanum-nickel, or zirconium-nickel, hydrogen storage alloys
Electrolyte: potassium hydroxide
NiMH batteries are a very compact energy source, but some memory effect, high self-discharging rate.
They are used in cell phones, cordless phones, compact camcorders, laptop computers, personal DVD and CD players.
NiMH batteries have 30% higher energy storage density than NiCad batteries.
Recharging procedures
They emit heat when charging so must be charged carefully.
They are more expensive than NiCad batteries, so are used for compact sources of power, but not deeply cycled, e.g. mobile phones, laptop computers.
NiMH batteries, are more environmentally friendly (no cadmium) and show a reduced memory effect.
Both nickel-based cells should never remain connected to a load when the current drops to zero, because their polarity will reverse.
So if a cell loses charge, remove it and recharge it without delay.

33.1.9 Weston Standard Cell
(Edward Weston, 1850-1936)
See diagram 3.33.01 : Weston Standard Cell.
The Weston Standard Cell is a glass "H" tube with platinum wires sealed into the bottom of each "H" in contact with the electrodes.
The positive electrode is liquid mercury under a mercurous sulfate paste.
The negative electrode is a cadmium / mercury amalgam.
Above the electrodes is cadmium sulfate solution.
The "saturated cell" has a layer of cadmium sulfate crystals above the electrodes and is used for very accurate measurements.
The "unsaturated cell" has packing to restrain the electrodes and is the most common, being portable.
The Weston Standard Cell became the international standard of voltage in 1908, 1.01830 International Volts at 20 oC.
So the cell must be kept at a constant temperature for accurate measurements.
It still can be used as a DC voltage reference.

33.1.3.12 Test a simple cell with different metals
See diagram 33.3.1 : Simple cell, magnesium and copper.
Experiment
Use alligator clips to connect a zinc strip to the negative terminal of a voltmeter and a copper strip to the positive terminal.
Dip the metals in dilute sulfuric acid separated from one another.
Note whether any reaction occurs.
Note the reading on the voltmeter.
The zinc dissolves in the acid.
Hydrogen comes from the copper.
From these reactions electrons flow through the external circuit producing a voltage of 1.1 V.
If a 1.5 V light bulb is connected in series, the light bulb glows for a short time then fades.
This happens because the hydrogen bubbles collecting on the copper strip reduces the flow of electric current.
Make the filament continue to glow by mechanically removing the bubbles or by adding the oxidizing agent potassium dichromate to the acid in the cell.
Repeat the experiment using magnesium, iron and lead in place of zinc.
The greater the difference in activity between the two metals the greater the voltage.
Test this by substituting magnesium, iron and lead for zinc, then record the voltages.

33.1.7.10 Test a simple electric cell with copper and zinc in dilute sulfuric acid
See diagram 3.84.3 : Voltaic cell.
Experiment
1. Put a clean piece of zinc and a clean piece of copper in separate test-tubes of dilute sulfuric acid.
Bubbles of hydrogen gas may come from the surface of the zinc.
No bubbles appear on the copper.
2. Put the zinc and copper in the same beaker containing dilute sulfuric acid.
Arrange the metals so that they touch, or connect them through an ammeter.
Many bubbles of hydrogen come from the copper and few or no bubbles come from the zinc.
The zinc is going into solution.
The electrons move towards the copper.
The zinc is the negative electrode.
The chemical energy of the zinc causes the electrons to flow.
Zn (s) --> Zn2+ (aq) + 2e-
At the surface of the copper, the electrons are transferred to the hydrogen ions of the sulfuric acid to form hydrogen gas.
The copper acts as a sort of catalyst here.
The copper is the positive electrode.
The copper is positively charged with respect to the zinc
2H+ + 2e- --> H2
or
2H3O+ (aq) + 2e- --> H2 (g) + 2H2O (l)
Overall reaction
Zn (s) + Cu2+ (aq) --> Cu (s) + Zn2+ (aq)

33.1.3.11 Simple galvanic cell
Experiment
Add zinc to dilute hydrochloric acid.
Hydrogen forms and the test-tube becomes hot.
Use this reaction as a source of electrical energy.
Connect a strip of zinc and a strip of copper with wire to a 1.1 V light bulb.
Pour the dilute acid into the beaker.
Note the hydrogen bubbles liberated at the copper strip and note the glow of the light bulb.
The chemical energy liberated as heat in the test-tube reaction is liberated as electrical energy in this simple galvanic cell.
Zn (s) + 2HCl (aq) --> ZnCl2 (aq) + H2 (g)
The half reaction equations are as follows:
Oxidation: Zno (s) - 2e- --> Zn2+ (aq)
Reduction: 2H+ (aq) + 2e- --> H2 (g)
Zno (s) + 2H+ (aq) --> Zn2+ (aq) + H2 (g) (Net reactions of electron transfer)

33.4.53 Leclanché cell
Primary dry cell batteries of the LeClanche system or zinc chloride system are usually called "carbon zinc" batteries.
Primary dry cell batteries are not designed for recharging.
A Leclanché cell battery has:
1. cathode of carbon and manganese dioxide,
2. anode of zinc alloy,
3. slightly acidic electrolyte solution of ammonium chloride and zinc chloride in water.
2MnO2 + 2NH4Cl + Zn --> ZnCl2 + 2NH3 + Mn2O3 + H2O
Experiment
See diagram 4.53 : Leclanché cell.
Remove the outer covering from an old dry cell.
Use a saw to cut the cell in half and note its structure.
Note the carbon (+ve) pole in the centre.
The zinc container is the negative (-ve pole).
The material between the two poles is the electrolyte.
Note how the zinc has been eaten away by the chemical.

33.4.53.1 Zinc chloride battery
A zinc chloride battery has:
1. cathode of carbon and manganese dioxide,
2. anode of zinc alloy,
3. slightly acidic electrolyte solution of mostly zinc chloride and some ammonium chloride.
The open circuit voltage of a fresh zinc chloride battery is > 1.60 volts.
The electrolyte has a greater volume and is more acidic than the LeClanche battery.
8MNO2 + 4Zn + ZnCl2 + 9H2O --> 8MnOOH + ZnCl2 + 4ZnO +5H2O

33.1.3.6 Magnesium / copper cell
See diagram 33.3.1 : Simple cell, magnesium and copper.
Experiment
Attach two wires to a light bulb.
Attach a piece of magnesium ribbon to one wire.
Attach a strip of copper foil to the other wire.
Hold the bulb in your hand and dip the magnesium and copper into a beaker of dilute sulfuric acid or a solution of sodium hydrogen sulfate.
The bulb lights, and bubbles of hydrogen gas form on the copper strip.
The chemical reaction between the magnesium and the acid causes a current of electricity to flow along the wire and light the bulb.
The current also makes the hydrogen gas bubbles come out of the acid at the copper strip, although the copper itself does not react.

33.1.7.6 Magnesium / copper battery
Experiment
1. Connect the external circuit before adding the sodium sulfate solution.
Clean copper in dilute nitric acid and clean magnesium ribbon in 1 M hydrochloric acid.
Half cells:
Magnesium ribbon is in contact with 0.5 M sodium sulfate solution, in a jar.
Copper strip is in contact with 0.5 M copper (II) sulfate solution, in a dialysis tubing bag, then in the same jar.
2. Fasten a short piece of metal to the end of the magnesium rod with the hose clamp, then wrap on foam insulation.
Use the handle of a screwdriver to loosely wrap around 10 turns of copper wire, then slip over the foam on the magnesium rod.
Wrap the assembly with foam insulation.
Immerse the assembly in water (with no added electrolyte), up to the top of the foam.
The battery produces about1.5 V and 20-100 mA continuously.
Each year, the magnesium rod and copper wire should be sanded to remove build-up.
Equations
Mg (s) --> Mg2+ (aq) + 2e- (Eo 2.36 V at 25oC, 1 atmosphere pressure)
1/2 H2 --> H+ + 2e- (Eo = 0)
Cu (s) --> Cu2+ (aq) + 2e- (Eo = -0.337 V)
Mg (s) + Cu2+ (aq) --> Mg2+ (aq) + Cu (s)
EMF = Eo (oxidation) - Eo (reduction) so EMF = 2.36 - (-0.337) = 2.7 V
At the anode: Oxygen is liberated: 4OH- --> O2 + 2H2O + 4e-.
At the cathode: Hydrogen ions are not reduced to H2, because the Eo of reaction 2., where Eo = 0 V, is greater than the Eo of reaction 3. (where Eo = -0.337 V).

33.1.7.7 Mercuric oxide battery
Mercuric oxide button-cell batteries are similar to silver oxide batteries.
They were banned in 1996 in accordance with the "Mercury Containing and Rechargeable Battery Management Act", and are no longer sold in the United States.
Larger mercuric oxide batteries may still be used in industrial equipment.
Before these batteries were banned, the sludge contents of a run down battery could be distilled to retrieve the mercury.
Cathode: (mercuric oxide or mercuric oxide + manganese dioxide) + graphite to prevent large droplets of mercury pooling.
Anode: zinc
Separation: paper or porous material soaked with electrolyte.
Discharge reaction: zinc oxidizes to zinc oxide, mercuric oxide reduced to elemental mercury.
Extra mercuric oxide prevents formation of hydrogen gas when battery run down.
mercuric oxide --> mercury metal
HgO + H2O + 2e- --> Hg + 2OH-
zinc --> zinc oxide
Zn + 2OH- --> ZnO + H2O + 2e-
potassium hydroxide electrolyte
KOH --> K+ + OH-
Cell reaction
Zn + HgO --> ZnO + Hg

Batteries
1. Originally, a "battery" was a combination of pieces of artillery, but this military term was first used to describe the first use of linked Leyden jars.
The first electrochemical battery was Alessandro Volta's pile of copper and zinc plates, separated by paper discs soaked in salt solution, so called a voltaic pile.
A battery is a source of electrical energy with electromotive force, EMF, measured in volts, equal to the potential difference between its terminals, assuming no loss of internal energy in the battery.
A battery is an external electricity source applied to an electrolytic cell that forces electrons around the circuit through the cell away from its the negative terminal and towards its positive terminal.
The term "battery" refers to several joined electrical cells, but one dry cell is commonly called a "battery", e.g. a torch battery, (flashlight battery).
2. Battery current
The electric current is carried as ions in the electrolyte within the electrolytic cell.
A current whose direction does not change with time is called direct current.
The current whose current intensity is invariable in the circuit is called constant current.
3. Battery potential
The end of the resistor where current enters is the high potential end.
Current flows through a resistor from high potential to low potential.
The positive terminal of a battery is always the high potential terminal, assuming the internal resistance is small.
In the external circuit of the electrical source, the constant current flows from the high potential to the low potential.
In the internal circuit of the electrical source, the current flows from the low potential to the high potential.
4. Battery sizes Common batteries provide 1.5 volts are sold in sizes AAA very small, AA small, C medium, and D large.
So four cells in series gives 6 volts and six cells in series gives 9 volts.
A battery supplies direct current from two or more connected electrochemical cells.
5. Dry cell batteries
A dry cell has the electrolyte in paste form that should not leak out of the battery.
Dry cell "batteries" include the following:
* Alkaline batteries (1.5 V) in cassette players and portable radios,
* Cadmium cells, Weston standard cadmium cells, cadmium amalgam + mercury electrodes, cadmium sulfate electrolyte.
* Nickel metal hydride batteries, discharge at 30% per month, also ready to use low self-discharge batteries,
* Nickel cadmium batteries (1.4 V) that are rechargeable and can produce a high current,
* Silver-oxide cells (1.5 V) in calculators, cameras and watches,
* Lithium ion batteries, replaced the lithium cadmium batteries, Li+, move from -ve electrode to +ve electrode when discharging
* Zinc carbon batteries (1.5 V), low cost, but short life,

3.1.2.4 Battery uses, types of batteries
Button cell, coin cells, watch battery, is a small, button-like battery that provides power for watches and very small devices.
Button cells are usually disposable primary cells.
Coin cell batteies are dangerous to youn children who may swallow tham, thinking they are sweets. In Australia, young childen have died from the chemical in coin cell betteries. Anodes may be zinc or lithium.
Cathodes may be manganese dioxide, silver oxide, carbon monofluoride, cupric oxide or oxygen from the air.
Lantern batteries
Car batteries
Watch batteries
Hearing aid batteries
Nickel-cadmiumn (NICAD, nickel metal hydride (NiMH)
Transistor batteries, introduced for the early transistor radios are nine volt batteries.
They have a rectangular prism shape with rounded edges and connector at the top.
They are used in portable radios, smoke detectors, gas detectors.

33.1.2.5 C-Rates of batteries
"C-rate" is a measure of the rate at which a battery is being discharged relative to its maximum capacity.
C is the nominal capacity in Ah or mAh.
At the "1 C rate", a 1.8 Ah battery should be able to deliver 1.8 A in one hour.
If discharged at 0.5 C rate, the nominal charge lasts two hours.
If discharged at 2C rate, the nominal charge lasts thirty minutes.
When C-rate is applied to charging:
> 0.5 C is fast charging,
0.2 C to 0.5 C is normal charging,
0.05 C to 0.2 C is trickle charging.
Too fast charging or discharging may generate heat inside the battery.
Too slow charging or discharging lowers effective capacity, because all batteries have some internal self-discharging.

33.1.2.6 Discharge curves of batteries
See diagram 33.86 : Discharge curves of good and bad batteries.
See diagram 33.86.1 : Discharge capacity and terminal voltage.
A discharge curve is the chart of the voltage level of a battery as it drops off during use.
Internal resistance causes the drop in voltage.

33.1.2.7 Energy density of battery energy storage density
Energy stored per kilogram of battery weight, watt-hours per kilogram, Wh /kg.

33.1.2.3 Battery disposal
In USA, battery labels must include the battery chemistry, the " three chasing arrows" and instruction that user must recycle or dispose of the battery properly.
Primary batteries and secondary batteries that no longer function should be disposed of according to local environmental regulations on disposal of batteries.
The Australian Battery Recycling Initiative (ABRI) is formed by battery manufacturers, recyclers, retailers, government bodies and environment groups to promote the collection, recycling and safe disposal of all batteries.

33.1.7.9 Silver-oxide battery
A silver-oxide battery uses silver oxide as the positive electrode (cathode), zinc as the negative electrode (anode) plus an alkaline electrolyte, NaOH or KOH.
The silver is reduced at the cathode from Ag (I) to Ag and the zinc is oxidized from Zn to Zn (II).
The chemical reaction that takes place inside the battery is the following:
Zn + Ag2O --> ZnO + 2Ag ( in the presence of KOH and NaOH).

33.3.1 Simple electric cell
Copper and zinc in dilute sulfuric acid produce electricity.
See diagram 33.3.1 : Simple electric cell.
Experiment
1. Put a piece of zinc metal and a piece of copper metal into a large beaker 3 / 4 full of dilute sulfuric acid.
Connect a galvanometer between the zinc metal and the copper metal.
Observe the deflection of the galvanometer needle, many hydrogen gas bubbles on the copper surface, but few bubbles on the zinc surface.
Zinc atoms transfer their electrons to the copper, so zinc atoms become zinc ions into the solution.
The copper transfers the electrons to hydrogen ions to form hydrogen gas by the contact surface of the copper with sulfuric acid.
The copper acts as catalyst.
2. Use a beaker containing dilute (5%) sulfuric acid, copper and zinc electrodes and a galvanometer to show current flow.
Note which electrode the bubbles gather on and how this affects the reading on the galvanometer?
Stir the liquid to dislodge the bubbles and read the galvanometer again.
3. Put some strong aqueous copper (II) sulfate solution in a beaker.
Connect copper foil to the positive terminal of a voltmeter and a zinc rod or foil to the other terminal.
Dip the two metals briefly into the copper (II) sulfate solution.
Note the readings on the voltmeter.
The voltage falls to zero after a short time, because copper deposited on the zinc and caused the reaction to stop.
Note what happens at the copper rod and at the zinc rod.
Determine the direction of electron flow.
4. Pour concentrated copper (II) sulfate solution into a beaker.
Insert a sheet copper and sheet zinc into the solution.
Measure the voltage drop by connecting the copper to the +ve terminal of a voltmeter and the zinc to the -ve terminal.
Record the reading.
Observe that the sheet zinc dissolves and hydrogen gas bubbles form on the surface of the copper.
Observe that the voltage will decrease to zero with increasing of the copper on the sheet zinc.
5. Repeat the experiment with magnesium ribbon, an iron nail or lead foil instead of the zinc.
Record the voltage each time.
he larger the difference in activity between two metals, the larger the voltage.
When copper deposits on the zinc electrode, it prevents more zinc from entering the solution.
This causes the voltage fall to zero after a short time and the cell becomes "dead".
You can add some subsidiary devices to prevent this happening.
6. Use a beaker containing dilute sulfuric acid, pieces of copper and zinc as electrodes, a switch, conducting wire and a galvanometer.
Note on which electrode the bubbles gather.
Note the effect of their formation on the galvanometer.
Stir the liquid to dislodge the bubbles and note what happens.
Instead of zinc and copper strips use two zinc strips and, later, two copper strips.
Note any current flow in either case.
7. Use an aluminium patty pan, the small disposable pan used in the oven for making cup cakes.
Put a salt solution in the pan and test different metals and objects, e.g. coins, hair pins, wires, as electrodes.
Attach one terminal of a galvanometer to the test electrodes and attach the other terminal to the aluminium pan.
Which combination makes the best electric cell?

33.1.7.1 Daniell cell
See diagram 33.3.2a : Daniell cell.
Daniell cell, porous pot, Zn in ZnSO4 solution/ copper in CuSO4 solution
The Daniell cell uses a porous pot to prevent copper depositing on the zinc.
Zn (s) --> Zn2+ (aq) + 2e-
Cu2+ (aq) + 2e- --> Cu (s)
Zn (s) + Cu2+ (aq) --> Cu (s) + Zn2+ (aq)
The Daniell Cell has a zinc plate in zinc sulfate solution and a copper plate in copper sulfate solution, with the two solutions separated by a porous earthen pot.
It produces about 1.08 volts.
It was once used as a voltage standard, but it has a short life and the EMF it produces in not constant.
Also, the electrodes must be removed from the electrolyte when the cell is not in use.
Experiment
1. The Daniell cell has EMF about 1.1 volts, and low internal resistance.
It is a primary cell whose EMF is constant for a considerable period of time and maintains a steady, small current.
However, since the copper (II) sulfate solution slowly diffuses through the clay porous pot to attack the zinc rod, the cell must be emptied and washed after use.
The Daniell cell uses a battery jar, clay porous pot, zinc electrode cylinder, copper electrode rod, 10% copper (II) sulfate solution, 10% zinc sulfate solution.
Put sheet zinc in a beaker containing zinc sulfate solution or dilute sulfuric acid solution.
Put a sheet copper in another beaker containing saturated copper (II) sulfate solution.
Connect the copper to the positive terminal of a voltmeter and the zinc to the negative terminal of the voltmeter.
Observe the voltmeter reading is zero.
Make a simple salt bridge by soaking filter paper in a concentrated solution of an electrolyte, e.g. sodium chloride or potassium nitrate.
Fix the filter paper to dip into the zinc sulfate and copper (II) sulfate solutions.
Observe the voltmeter and record the reading.
The voltmeter shows that current is flowing.
Disconnect the voltmeter and substitute a 1.5 volt bulb, ammeter, and conducting wire.
Examine the electrodes after two minutes.
Note the zinc corrodes and new copper has deposited on the copper electrode.
Observe whether the copper (II) sulfate solution loses some of its blue colour.
To prolong the life of the salt bridge, make a permanent salt bridge from a glass U-tube filled with a 1 M aqueous potassium nitrate solution.
You can mix the solution with agar gel to keep it in the U-tube.
Put cotton wool plugs at each end of the U-tube.
2. Put a porous pot in a beaker.
Pour 0.5 M zinc sulfate solution into the porous pot.
Pour concentrated copper (II) sulfate solution into the beaker and fill to the same level as the zinc sulfate solution.
Make a cylinder shape with copper foil and place it in the beaker to surround the porous pot.
Connect the copper foil to the positive terminal of a 1 to 5 V voltmeter.
Connect a zinc rod to the negative terminal of the voltmeter and lower the zinc rod into the zinc sulfate solution.
Note the reading of the voltmeter.
Insert a 1.5 V light bulb in place of the voltmeter and note whether it lights.
Insert an ammeter into the circuit to find the current flowing.
Note whether you can vary the current by moving the copper electrode nearer to the zinc electrode, or by altering the surface area of the copper foil by raising and lowering it in the solution.
Oxidation occurs at the zinc anode, losing electrons to the electric circuit.
Reduction occurs at the copper cathode, gaining electrons from the electric circuit.
At the anode, zinc atoms lose two electrons to become zinc ions, Zn2+.
At the cathode, copper ions, Cu2+ receive two electrons from the electric circuit to become copper atoms.
The zinc ions in solution gather around the zinc anode.
The copper ions are removed from solution when they are reduced to copper metal and join the copper cathode.
In the solution, sulfate ions move towards the zinc anode and copper (II) ions move towards the copper cathode.
Positive ions, cations, move towards the cathode and negative ions, anions, move towards the anode.
So charge is carried by ions in solution and carried by electrons in the electric circuit.
The reactions continue until all the zinc in the anode joins the solution or all the copper ions are plated onto the cathode.
3. Instead of a porous pot, repeat the experiment with a salt bridge made from a U-tube containing M potassium nitrate solution and agar gel.
4. Pour concentrated copper (II) sulfate solution into a clay porous pot in a large beaker.
The solutions should be at the same level.
Bend a sheet copper into a cylinder shape and put it in the beaker to surround the porous pot.
Put sheet zinc into the porous pot.
Connect the copper to the positive terminal of a voltmeter.
Connect the zinc to the negative terminal of the voltmeter.
Record the reading and observe whether there are changes in colour of the sheet copper, the sheet zinc and the solution.
Disconnect the voltmeter and substitute: 1.5 volt bulb, ammeter; conduction wire.

33.4.51 Electricity from two coins
1. Take two coins made of different metals.
Clean them well with steel wool or fine sand paper.
Fold some paper into a pad so that it is larger than the coins.
Soak the paper in salt water.
Place one coin on top of the pad and the other underneath.
Hold them between your thumb and finger.
Connect both leads of a sensitive galvanometer or multimeter to the coins and note the deflection.
2. Soak absorbent cotton in salt water.
Obtain two coins made of different metals.
Place the cotton between the coins.
The thickness of the cotton should be more than 2 mm.
Connect both leads of a sensitive galvanometer to the coins at the same time.
Observe the current flowing at the galvanometer.
Clean the coins with steel wool or fine sandpaper beforehand and make sure that no salt water remains on the leads.
3. Put aluminium foil with a copper coin on it in water for a day.
The water appears cloudy and the aluminium foil is perforated where the coin was lying on it.
The water becomes cloudy due to dissolved aluminium.
4. Make a pile of copper coins alternating with pieces of sheet zinc.
Between each pair of metals insert newspaper soaked in sodium chloride solution.
Wind thin, insulated copper wire 50 times around a plotting compass.
Press each end of the wire against each end of the pile.
A current causes a deflection of the compass needle.
5. Make a simple cell with two coins.
Use two coins made of different metals.
Clean them well with steel wool or fine sand paper.
Fold some paper hand towel or absorbent paper into a pad so that it is larger than the coins.
Soak the absorbent paper in salt water.
Put one coin on top of the pad and the other underneath.
Hold them between the thumb and finger.
Connect both leads of a sensitive galvanometer to the coins and watch the deflection.
6. Use two coins made of different metals.
Clean them well with steel wool or fine sand paper.
Fold some paper into a pad so that it is larger than the coins.
Soak the paper in salt water.
Place one coin on top of the pad and the other underneath.
Hold them between your thumb and finger.
Connect both leads of a sensitive galvanometer or multimeter to the coins and note the deflection.

33.1.7.5 Lemon cell
Lemon battery, lemon cell, electricity from lemons
"Lemon Clock", connects electronic clock to lemon with Cu and Zn electrodes, (toy product).
See diagram 32.149 : Lemon cell.
Experiments
1. Connect a wire to a piece of zinc.
Use zinc cut from the can of a used dry cell, torch battery.
Connect another wire to a piece of copper.
Roll a lemon on the table with your hand to break up some tissue inside.
Push the zinc and copper strips through the skin of the lemon so that they do not touch.
Connect both leads of a sensitive galvanometer or multimeter to wires and note the deflection.
Repeat the experiment using a potato.
Note whether the distance between the metal strips affects the galvanometer reading.
Oxidation: Zn --> Zn2+ + 2e-
Reduction: 2H+ + 2e- --> H2
2. Connect copper and zinc electrodes in a lemon to a digital voltmeter.
Connect a wire to a piece of zinc.
Use zinc cut from the can of a used dry cell, torch battery.
Connect another wire to a piece of copper.
Roll a lemon on the table with your hand to break up some tissue inside.
Push the zinc and copper strips through the skin of the lemon so that they do not touch.
Connect both leads of a sensitive galvanometer or multimeter to wires and note the deflection.
Electrons flow from the zinc to the copper electrode.
The citric acid in the lemon acts as a salt bridge.
Note whether the distance between the metal strips affects the galvanometer reading.
Repeat the experiment with potatoes, oranges, grapefruit or soft drinks.
3. Gently press or roll a lemon on the table to squash the tissue inside.
Connect one terminal of a galvanometer to a piece of zinc and connect the other terminal to a piece of copper.
Push the two pieces of metal through the skin of the lemon.
The metals must not touch.
Observe the deflection of the galvanometer needle to see whether current flows.
The lemon juice acts as an electrolyte.
Does the distance between the metals affect the deflection of the galvanometer needle?
Repeat the experiment with a potato.
There is almost no deflection.
4. Use a lemon, orange or any juicy fruit, a copper wire, a coated iron wire a centre zero galvanometer out of the storeroom, a voltmeter, 0 to 10 volts.
Put the different kinds of wire in turn into the lemon the acidified water Add a few drops of dilute sulfuric acid, and acidified hydrogen gas peroxide solution, 5 mL of hydrogen peroxide + a few drops of dilute sulfuric acid.
Connect the wires to a galvanometer in each case and observe the pointer when the wires are in each electrolyte or solution, e.g. fruit juice, or H2O2 solution.
Leave the wires in the acid and see that the zinc dissolves off the iron beneath it.
The copper cathode seems unchanged, but is covered with bubbles of hydrogen gas.
The peroxide removes these bubbles.
Try to light a 11 / 2 volt globe.
The a series connection means copper of one cell to zinc of the next cell ...
5. Connect one terminal of a galvanometer to a piece of zinc and connect the other terminal to a piece of copper.
Use the hand to roll a lemon on the table to squash the tissue inside.
Push the two pieces of metal through the skin of the lemon.
The metals must not touch.
Note any deflection of the galvanometer needle.
The lemon juice acts as an electrolyte.
Does the distance between the metals affect the deflection of the galvanometer needle?
Repeat the experiment with a potato.
There is almost no deflection.
6. Make two slits in the skin of a lemon and push a piece of copper (positive terminal), and a zinc washer or piece of aluminium foil, (negative terminal.
Attach wires to make a circuit.
A chemical reaction takes place between the metals and the acid in the lemon juice, causes the current to flow and light a 1.5 volt bulb.
7. To make a lemon battery / voltaic cell, "Lemon screamer lasagne cell", stick copper and galvanized steel electrodes into a lemon.
Attach a voltmeter or a galvanometer.
Repeat the experiment with a same size potato.
8. Squeeze lemon juice over pieces of absorbent paper so that they become damp, but not dripping wet.
Make an alternating pile of 2p and 10 p (UK) coins with a piece of wet absorbent paper between each coin.
The larger coin should be at the top and the bottom of the pile.
Wet the tips of the index finger and the thumb, pick up the pile of coins and squeeze the coins tightly.
Fell the slight electric shock.
This experiment may not work, but it is worth trying with your local coins.
The British coins have change composition.
Since 1992, the two pence 2.03 mm thickness coin is made of copper-plated steel.
Since 2012, the ten pence 2.05 mm thickness coin is made of nickel-plated steel.
9. Make a lemon-powered clock
Use a digital clock without a plug, powered by two AA batteries, two galvanized nails, e.g. 16d (3.5 inches), two lengths of bare copper wire + 3 crocodile clips.
Cut a large lemon in halves to form half lemon A, and half lemon B.
Push the galvanized nails into one end of each half lemon.
Push the ends of the lengths of bare copper wire into the other ends of each half lemon.
Remove the batteries from the clock and note the positive and negative terminals.
Clip a length of copper wire to the positive terminal of the clock.
Use a crocodile clip to connect the wire from half lemon A to the positive terminal of the clock.
Use a crocodile clip to connect the galvanized nail in wire half lemon B to the negative terminal of the clock.
Use a crocodile clip to join the galvanized nail in half lemon A to the wire in half lemon B.
Citric acid in the lemon juice electrolyte dissolves some zinc on the galvanized nail so the nail loses electrons and become positive.
The copper wire gains some of the electrons and become negative.
So a circuit is formed through the clock and the clock can go again.
10. Cut a lemon into four pieces.
Cut the peel in each piece and insert a "copper" coin half way in.
Insert a nail into the side of each piece.
Cut nine pieces of copper wire and use wire cutters to crimp two test clips to each end.
Connect the negative side of a 4V LED, (the flat edge on the bulb of the LED) to each of the coins with an alligator clip.
Attach a second wire to the nail of this slice and to the coin of another piece and so link all three of the four pieces.
Connect the last wire end to the nail of this slice and to the positive side of the LED light.
Observe light in the LED.
Electrons leave the coins so the charge on them becomes positive.
The nails receive the free electrons and so the charge on them becomes negative.
Electrons flow from negative to positive across the LED so it lights up.
Repeat the experiment using different coils as electrodes.

33.3.5 Prepare a simple chemical rectifier
Use a glass container filled with saturated solution of borax and electrodes of aluminium and lead.
Use a low volt direct current power source.
You can use this half wave rectifier as a battery charger.
Connect four such cells to form a bridge rectifier.
Check "+" and "-" terminals with a voltmeter.

33.1.3.1 Aluminium foil chocolate wrapper cell in the mouth
When the aluminium foil from a chocolate wrapper ("silver paper") touches an amalgam filling, mainly tin, two metals are in contact in the saliva electrolyte.
A current is generated, the aluminium tending to dissolve and the current of electrons sends an unpleasant pulse along the nerves.
Also a metallic taste may be due to dissolved aluminium ions, Al3+.

33.1.3.8 Potato cell
Noisy potato cell
Push 2 cm of copper wire and zinc wire one into a raw potato.
Hold an earphone connected to the wires to hear a crackling sound caused by a weak electric current.

33.3.9 Ionic migration
Migration of coloured ions in an electrolyte can be shown with a flat chamber the size of a microscope slide and coloured ions, e.g. MnO4-, Cu2+, Cr2O2.
Dissolve some sodium sulfate in sufficient water to half fill a U-tube.
Add drops of universal indicator that is red in acidic solutions and blue purple in alkaline solutions.
The colour of the indicator should be green showing that the solution is neutral.
To an equal volume of water add 1 gram agar agar gel for each 100 mL of water.
Warm until the gel dissolves and then mix the two solutions.
Pour this solution into the U-tube until the arms are about half full.
When the gel has set, pour dilute sulfuric acid into one arm and dilute sodium hydroxide into the other.
Insert platinum or carbon electrodes into the solutions.
Connect the electrode in contact with sulfuric acid to the positive terminal of a battery.
Connect the electrode in contact with the sodium hydroxide solution to the negative terminal of the battery.
Allow the current to pass for some time and observe the colour changes produced in each arm.
The violet colour in the gel below the sodium hydroxide solution is because of the movement of hydroxide ions into it under the influence of the electric field.
The red colour in the gel below the sulfuric acid solution is because of the movement of hydrogen ions into it.
So there is evidence for a two way flow of ions.

33.1.5.1 Dry cell, Bring a dead battery to life
Warm a used 1.5 v torch battery.
The bulb may light again.
The zinc container of a battery cell becomes corroded by the ammonium chloride solution as a paste.
This creates an excess of electrons in the zinc and an electron loss in the carbon rod.
The carbon rod is coated with manganese dioxide to prevent the build-up of hydrogen gas that would stop the reaction.
The bulb's incandescent filament gives out light when enough electrons flow through it.
When the chemical reactions in the battery slow and flow of the electrons are not enough to make the filament glow.
However, warming the battery accelerates the chemical reaction so that the filament can briefly give out light again.
No chemical reaction can occur.
When the zinc has corroded entirely and turned into white powder, zinc chloride.

33.1.5.2 Dry cell capacity
Dry cells are voltaic cells with the electrolyte in the form of a paste, usually ammonium chloride, to avoid spilling of the electrolyte.
The total charge output from the cell is called the capacity of the cell.
You can measure capacity of a cell or accumulator in the number of ampere hours of charge it can deliver.
The capacity of the cell and its work principle depend on the volume.
The more the capacity the cell has, the greater work current appliance can be used and the longer the times it is used.
Experiment
Use different dry cells.
Connect a dry cell with a light bulb to form a circuit.
Observe the normal brightness of the bulb and record the time.
Then insert in parallel other bulbs to the bulb on the circuit in turn and observe the variation of the brightness of the bulbs.
When the bulbs no longer light, stop the experiment and record the number of the bulbs connected to the circuit and time after first connection in the circuit.
If the bulbs still give light, wait until they all no longer light and record the time from the first connection in the circuit.
Repeat the experiment with some different size dry cells.
Note the numbers of the bulbs increased until they go out and stop doing the experiment.
Compare the difference of the discharge time between the dry cells.
Use a millivoltmeter to measure the voltage of dry cells before and after discharge.
The voltage of new cells is about 1.5 V.
The voltage of "no charge" cells is about 0.75 V, the discharge stop voltage.

33.1.5.3 Dry cell in electric circuit
See diagram 4.54 : Dry cell in electric circuit.
Experiment
Connect an electric light bulb, e.g. 4.4 volts, V, 0.5 amps, A, and a lampholder, to the +ve and -ve terminals of a dry cell,
or a lead cell accumulator or a low voltage power supply.
Notice the filament made of tungsten carbide.
Passage of the electric current through the tungsten carbide wire causes it to become very hot and give off light.
Reverse the connections to the source of electricity and the lamp still operates although the electricity is flowing in the opposite direction.
Draw a diagram of this simple circuit to show the path of the current through the light bulb and around to the other end of the cell.

33.1.5.4 Dry cells in an electric circuit, cells in series and parallel
See diagram 32.1.6.4 : Dry cells in series and parallel.
See diagram: 32.151.1 : Simple electric circuit.
See diagram: 32.151.2 : Torch battery electrical experiments.
Experiment
1. Observe the effect on current of increasing potential difference.
Use an ammeter to record the electric current flowing when 1, 2 and 3 of the 1.5 volt dry cells are connected in series in the circuit.
The greater the rate at which the electrons pass, the further the needle moves in the ammeter.
Increasing the potential difference increases the current that flows through the wire.
2. Observe the current through an electric jug element when voltage drop changes.
Sretch out and cut off about 15 cm of the jug element, screw it firmly across the terminals of the voltmeter.
Connect your ammeter, switch and four dry cells, all in series.
Record the voltage when 4, 3, 2, 1 of the 1.5 volt dry cells are connected in series.
3. Connect an electric bulb, e.g. 2.4 V, 0.5 A, and lampholder, to the +ve and -ve terminals of a dry cell, or lead cell accumulator, or low voltage power supply.
Notice the filament made of tungsten carbide.
Passage of the electric current through the tungsten carbide wire causes it to become very hot and give off light.
Reverse the connections to the source of electricity and the lamp still operates although the electricity is flowing in the opposite direction.
Draw a diagram to show the path of the current through the bulb and around to the other end of the cell.
This is a simple electric circuit.
Use circuit diagrams to represent the electrical components in a circuit.

33.1.5.5 Dry cell, electric torch, (flashlight), battery
See diagram 4.53 : Leclanché cell.
Investigating a dry cell: A Plastic wrapper, B Zinc case (cathode) (-ve) C Ammonium chloride paste
(electrolyte) D Manganese dioxide + carbon particles, E Carbon rod (anode) (+ve)
Experiment
Remove the outer covering from an old dry cell.
Use a saw to cut the cell in half and note its structure.
Note the carbon (+ve) pole in the centre.
The zinc container is the negative (-ve pole).
The material between the two poles is the electrolyte.
Note how the zinc has been eaten away by the chemical.

33.1.5.6 Zinc-carbon cell, dry cell torch battery, (flashlight battery)
See diagram 32.150 : Torch battery cut vertically.
See diagram 32.154.1 : Electric torch.
A battery supplies direct current from two or more connected electrolytic cells.
The term "battery" refers to several joined electrical cells, but one dry cell is commonly called a "battery", e.g. a torch battery, flashlight battery.
A dry cell has the electrolyte in paste form that should not leak out of the battery.
Dry cell "batteries" include the following:
1. Alkaline batteries (1.5 V) in cassette players and portable radios,
2. Nickel metal hydride batteries, discharge at 30% per month, also ready to use low self-discharge batteries,
3. Nickel cadmium batteries (1.4 V) that are rechargeable and can produce a high current,
4. Silver-oxide cells (1.5 V) in calculators, cameras and watches,
5. Lithium ion batteries, replaced the lithium cadmium batteries, Li+ move from -ve electrode to +ve electrode when discharging, in consumer electronics,
6. Zinc carbon batteries (1.5 V), low cost, but short life,
7. Cadmium cell, Weston standard cadmium cell, cadmium amalgam + mercury electrodes, cadmium sulfate.
Experiment
1. Remove the outer covering from an old dry cell, e.g. electric torch, 2.4V, 0.5A.
Use a saw to cut the cell in half and observe its structure.
Note the carbon (+ ve pole) in the centre.
The zinc container is the negative (- ve pole).
The material between the two poles is the ammonium chloride electrolyte.
Note how the zinc has been eaten away by the chemical.
Draw a circuit diagram.
Note the directions of insertion of batteries.

33.1.7.2 Galvanic cell, Voltaic cell
See diagram 3.2.84 : Copper and zinc foil in a voltmeter.
See diagram 3.85.2 : Voltaic cell
See diagram 3.85.1 : Voltaic cell
Experiment
1. Zinc + copper in concentrated copper sulfate solution
* Put concentrated copper (II) sulfate solution in a beaker.
Connect copper foil to the positive terminal, red wire, of a voltmeter and a zinc foil to the negative terminal, black wire.
Simultaneously dip the two metals briefly into the copper sulfate solution.
Record the readings on the voltmeter.
The voltage falls to zero after a short time, because black copper deposited on the zinc and caused the reaction to stop.
When copper deposits on the zinc electrode, it prevents more zinc from entering the solution.
This causes the voltage to fall to zero after a short time and the cell becomes "dead".
You can separate the electrolytes to prevent the voltage fall by using 1. a Daniell Cell that has a porous pot or 2. a salt bridge.
* Pour concentrated copper (II) sulfate solution into a beaker.
Connect a copper rod to the positive terminal of a voltmeter and a zinc rod to the negative terminal.
Dip the two metals briefly into the copper (II) sulfate solution.
Zinc dissolves and hydrogen bubbles form on the surface of the copper.
The voltmeter reads 1.1 V, so electrons are moving from the zinc to the copper.
2. Zinc + Copper in dilute sulfuric acid
See diagram 3.84.3 : Voltaic cell.
* Put a clean piece of zinc and a clean piece of copper in separate test-tubes of dilute sulfuric acid.
Bubbles of hydrogen gas may come from the surface of the zinc.
No bubbles appear on the copper.
* Put the zinc and copper in the same beaker containing dilute sulfuric acid.
Arrange the metals so that they touch, or connect them through an ammeter.
Many bubbles of hydrogen come from the copper and few or no bubbles come from the zinc.
The zinc is going into solution.
The electrons move towards the copper.
The zinc is the negative electrode.
The chemical energy of the zinc causes the electrons to flow.
Zn (s) ---> Zn2+ (aq) + 2e-
At the surface of the copper, the electrons are transferred to the hydrogen ions of the sulfuric acid to form hydrogen gas.
The copper acts as a sort of catalyst here.
The copper is the positive electrode.
The copper is positively charged with respect to the zinc.
2H+ + 2e- ---> H2
or
2H3O+ (aq) + 2e----> H2 (g) + 2H2O (l)
3. Use alligator clips to connect a zinc strip to the negative terminal of a voltmeter and a copper strip to the positive terminal.
Dip the metals in dilute sulfuric acid separated from one another.
Note whether any reaction occurs.
Note the reading on the voltmeter.
The zinc dissolves in the acid.
Hydrogen comes from the copper.
From these reactions electrons flow through the external circuit producing a voltage of 1.1 V.
If a 1.5 V light bulb is connected in series the light bulb glows for a short time then fades.
This happens because the hydrogen bubbles collecting on the copper strip reduces the flow of electric current.
Make the filament continue to glow by mechanically removing the bubbles or by adding the oxidizing agent potassium dichromate to the acid in the cell.
Repeat the experiment using magnesium, iron and lead in place of zinc.
The greater the difference in activity of the two metals the greater the voltage.
Test this by substituting magnesium, iron and lead for zinc, then record the voltages.
4. Put zinc + copper in dilute hydrochloric acid.
Add zinc to dilute hydrochloric acid.
Hydrogen forms and the test-tube becomes hot.
Use this reaction as a source of electrical energy.
Connect a strip of zinc and a strip of copper with wire to a 1.1 V light bulb.
Pour the dilute acid into the beaker.
Note the hydrogen bubbles liberated at the copper strip and note the glow of the light bulb.
The chemical energy liberated as heat in the test-tube reaction is liberated as electrical energy in this simple galvanic cell.
Zn (s) + 2HCl (aq) ---> ZnCl2 (aq) + H2 (g)
The half reaction equations are as follows:
Oxidation: Zno (s) - 2e- ---> Zn2+ (aq)
Reduction: 2H+ (aq) + 2e- ---> H2 (g)
Zno (s) + 2H+ (aq) ---> Zn2+ (aq) + H2 (g) (Net reaction of electron transfer)

33.1.7.3 Voltaic cell, galvanic cell, with salt bridge
See diagram 3.84.5 : Voltaic cell with salt bridge
Experiment
1. Zinc and copper galvanic cell
Put a zinc rod in a beaker containing zinc sulfate solution and put a copper rod in a beaker containing copper (II) sulfate solution.
Connect the copper to the positive terminal of a voltmeter and the zinc to the negative terminal of the voltmeter.
The reading of the voltmeter is zero.
Make a salt bridge by filling a glass U-tube with 1 M aqueous potassium nitrate solution and agar gel,
or by soaking filter paper in a concentrated solution of an electrolyte, e.g. sodium chloride or potassium nitrate.
Fix the filter paper to dip into the zinc sulfate and copper (II) sulfate solutions.
The voltmeter shows that current is flowing, so read the voltmeter.
Disconnect the voltmeter and substitute: 1.5 V light bulb, ammeter, conducting wire.
Examine the electrodes after 2 minutes.
The zinc corrodes and new copper has deposited on the copper electrode.
The copper (II) sulfate solution loses some of its blue colour.
Make a more permanent salt bridge with a glass U-tube filled with a 1 M potassium nitrate solution.
The solution may be mixed with agar gel to keep it in the U-tube.
Put cotton wool plugs at each end of the U-tube.
Note the voltage, the current and whether the light bulb glows.
2. Zinc and tin galvanic cell
The zinc half cell has zinc metal in a solution containing Zn2+ (aq) ions.
The tin half cell has tin metal in a solution containing Sn2+ ions.
The half cells are connected together with a wire and a salt bridge.
Oxidation occurs in the Zn2+ (aq) / Zn (s) half cell, so zinc is the anode, which loses electrons and enter the solution as Zn2+ (aq),
leaving behind electrons.
So this electrode becomes negatively charged.
Reduction occurs in the Sn2+ (aq) / Sn (s) half cell, so tin is the cathode, which gains two electrons from the electrode as Sn (s).
So this electrode becomes positively charged.
Electrons travel from the negatively charged electrode to the positively charged electrode through the wire, so they flow from the anode to the cathode, i.e. from the zinc electrode to the tin electrode.
The half cells are linked by a salt bridge separator that allows transfer of ions, but not water.
The salt bridge contains a solution of a salts, e.g. NaNO3, KNO3, NaCl or KCl.
They do not take part in the reaction, but when the half cells are connected ions move between the salt bridge and the half cells.
So that the whole cell remains neutral and to complete the circuit.
In the anode half cell, Zn2+ (aq) ions enter the solution as oxidation occurs, leading to a build up of positively charged cations in the solution.
However, anions in the salt bridge, e.g. NO3- (aq) or Cl- (aq), flow into the half cell to neutralize the solution.
In the cathode half cell, Sn2+ (aq) ions leave the solution as reduction occurs, leading to a loss of positively charged cations in the solution.
However, cations in the salt bridge, e.g. Na+ (aq) or K+ (aq), flow into the half cell to neutralize the solution.
Using the standard cell potentials.
E = -0.76 V for Zn2+ / Zn (s),
E = -0.14 V for Sn2+ / Sn (s),
So value for the Sn2+ / Sn (s) half cell is less negative:
Sn2+ (aq) + 2e- --> Sn (s),
Zn (s) --> Zn2+ (aq) + 2e-,
Zn2+ (aq) + Zn (s) --> Cu (s) + Zn2+ (aq)
In the zinc and tin galvanic cell, the standard cell potential is the reading on the voltmeter when the cells are first connected:
Ecell = E reduction half cell - E oxidation half cell = (-0.14 V) - (-0.76 V) = + 0.62 V

33.3.14 Magnesium pencil sharpener, electrodes
1. Magnesium pencil sharpeners have steel blades.
When the body and blade are separated and placed in cola, they act as electrodes in a galvanic cell producing almost 2 volts.
The iron blade acts as the cathode.
The magnesium body acts as the anode.
2. Use a screwdriver to remove the steel blade from a magnesium pencil sharpener.
Use crocodile clips to connect the positive wire from a voltmeter to the steel blade and the negative wire to the magnesium.
Dip the steel and magnesium electrodes into a dilute sodium chloride solution in a Petri dish.
Record the voltage measured by the voltmeter (about 1.3 V).
Repeat the experiment by substituting cola for the sodium chloride solution.
Record the voltage measured by the voltmeter (about 1.9 V).
Observe tiny bubbles on the steel blade electrode.
This experiment shows that cola would function better as a salt bridge then sodium chloride solution!
This may be because acids in the cola contribute hydrogen ions that were reduced to hydrogen gas.
At the anode, oxidation occurs: Mg --> Mg2+ + 2e-
At the cathode, reduction occurs: 2H+ + 2e- --> H2
Repeat the experiment with more tan one galvanic cell (with cola) in series to try to light a LED (light emitting diode).

33.1.7.4 Leclanché cell
Leclanché cell, dry cell, electric torch (flashlight) battery
See diagram 3.88 : Leclanché cell.
See diagram 32.150 : Torch battery cut vertically.
A Leclanché cell (Georges Leclanché 1839-1882, France) is a primary voltaic cell with a carbon rod anode, zinc cathode, dilute ammonium chloride solution electrolyte and e.m.f. approximately 1.5 volts.
Zn + H2SO4 --> (discharge) ZnSO4 + H2O + H2 (g) The term "battery" refers to several joined electrical cells, but one dry cell is commonly called a "battery", e.g. a torch battery, (flashlight battery).
A torch "battery" is the dry cell version of the Leclanché cell.
It has manganese dioxide [manganese (IV) oxide] around the carbon rod to oxidize hydrogen and so depolarize the anode.
The electrolyte is in the form of a water paste so the dry cell is not really "dry".
2MnO2 + H2 --> Mn2O3 + H2O
manganese dioxide + hydrogen --> dimanganese trioxide + water.
1. Remove the outer covering from an old dry cell. Use a saw to cut the cell in half and note its structure.
Note the carbon (+ve terminal) conducting rod that runs along the axis.
The zinc canister is the -ve terminal.
The material between the two terminals is the electrolyte..
Note how the zinc has been eaten away by the chemical.
In older batteries, the canister leaked if two much of the zinc canister dissolved.
However, nowadays the canister has an outer jacket to ensure that the battery is "leak proof".
However, batteries may still link if they are left in the sun and a build-up of hydrogen gas ruptures insulating seals or the outer jacket.
Alkaline batteries containing potassium hydroxide electrolyte may rupture to form white feather-like crystals of potassium carbonate when the potassium hydroxide reacts with the carbon dioxide in the air.
2. Add 1 mL of phenolphthalein indicator to 2 cm of ammonium chloride solution in a shallow dish.
Attach a crocodile clip and conducting wire to a carbon rod from a dry cell battery and a piece of zinc foil.
Join the conducting wires.
Dip the carbon rod and zinc foil into the solution to act as electrodes.
Note any changes around the electrodes.
3. The voltage from a single "battery" is usually about 1.5 V, however maximum current varies with the type of battery.
Batteries for radios produce a small current for long periods, e.g. 4 amps.
Batteries for torches or flashlights produce a large current for short periods, e.g. 6 amps.
Connect an ammeter directly across the terminals of different types of batteries.
4. Use a saw to cut through a used dry cell lengthways.
The central carbon rod is the positive electrode.
The zinc container is the negative electrode.
The moist paste contains the electrolyte ammonium chloride that acts like an acid in dissolving zinc and the black manganese (IV) oxide that acts as a depolarizer by oxidizing hydrogen.
It slowly oxidizes hydrogen to water so that any hydrogen formed cannot block the flow of electric charge.
If the power from a dry cell fades after continual use, the power may be restored after some hours, because of the depolarizing action of the manganese (IV) oxide.
In a used battery the zinc appears eaten away, because some of it has gone into solution.
5. Cut down one side of a dry cell battery with a hacksaw.
Scoop out the contents and put them in a beaker half full of water, then stir.
Put cotton wool in the neck of a funnel to act as a filter.
Pour the contents of the beaker into the funnel.
The black substance remaining in the cotton wool is mainly manganese dioxide.
Evaporate the clear solution that had passed through the filter.
White ammonium chloride and some zinc chloride remain after evaporation.
Test for ammonia by adding solid sodium bicarbonate and heat until you can smell the ammonia given off.
Also, a white precipitate of zinc hydroxide forms that dissolves again to form the zincate ion [Zn(OH)4]2-.
6. Mix 4 g of carbon black, powdered carbon, with 10 g manganese (IV) oxide, the oxidant.
Stir in ammonium chloride solution to make a thick paste.
Make a zinc can made from zinc foil rolled into a cylinder or a cleaned zinc can from an old dry cell.
Cut absorbent paper to make a cylinder to line the zinc can.
Place the mixture on the absorbent paper, compress it into a cylinder and wrap the absorbent paper around the cylinder so that it just fits inside the zinc can.
Pour ammonium chloride solution between the paper and the zinc to ensure good contact.
Press the mixture into the zinc can firmly and tightly.
Attach a crocodile clip and conducting wire to a carbon rod from an old dry cell.
Press the carbon rod down the centre of the mixture so that it almost touches the bottom of the zinc can.
Connect another crocodile clip and conducting wire to the zinc container.
This cell should light a 1.5 V light bulb and run a small 1.5 V electric motor.
The carbon lowers the internal resistance of the cell.
Test the voltage and the current from the cell.
Zn (s) ---> Zn2+ (aq) + 2e- (oxidation)
2NH4+ (aq) + 2MnO2 (s) + 2e- ---> Mn2O3 (s) + 2NH3 (aq) + H2O (aq) (reduction)

33.91 Dry cell batteries comparison experiment
Connect different dry cell batteries to a light bulb or a clock.
Note the length of time the bulb remains lit or when the clock stops.
Researchers have found that the battery lasting the longest was "Duracell", then "Energizer", then "Eveready".
This sequence reflects the relative cost of the battery.

33.1.2.2 Battery comparison, common rechargeable batteries

Type (Chemistry)	Cell voltage	Energy density	Cycle life	Charge Time	Discharge Rate	Cost
SLA: 33.1.8.2	2.0	30 Wh / kg	long	8-16	0.2 C	low
RAM: 33.6.13	1.1	75 Wh / kg	short	2-6	0.3 C	low
NiCad: 33.1.8.5	1.25 V	45-80 Wh / kg	long	14-16	>2 C	medium
NimH: 33.1.8.6	1.25 V	60-120 Wh / kg	medium	2-4	>0.5 C	higher
Li-ion: 33.1.8.4	3.6 V	110-160 Wh / kg	long	3-4	<1C	highest
Lead acid, SLA: 33.1.8.2	2.0 v					low

Other types of rechargeable batteries include:
Nickel-iron battery, NiFe cell, Edison cell, nickel (III) oxide-hydroxide +ve, iron -ve, potassium hydroxide electrolyte,
Silver-zinc battery.