Transcript metals - whitburnscience

NATIONAL 5 CHEMISTRY
UNIT 3
CHEMISTRY IN SOCIETY
CONTENTS
• Metals
• Properties of plastics
• Fertilisers
• Nuclear Chemistry
• Chemical Analysis
METALS
•Metallic bonding and properties (they conduct electricity)
•Reactivity of metals with oxygen, acid and water (balanced
ionic equations can be written).
•Metal ores and percentage (%) composition
•Extracting metals by heat alone, heating with carbon and
electrolysis. [depends on reactivity of metal and all metals
are reduced in the process]. (balanced ionic equations can
be written and reduction equations for the metals can be
written.)
•Electrochemical cells and REDOX equations
•Fuel Cells and Rechargeable batteries
Metallic bonding
• Positive metal ions surrounded by de-localised
electrons (electrons that are free to move). This
is why metals have the properties they have
Properties of Metals
• Density – this is the mass of a substance in a given
volume.
A high density material is much heavier than the same volume of a
low density material e.g. aluminium (low density) – used to build
aircraft. Lead (high density) – is used as weights for fishing
nets/lines.
• Thermal Conductivity - metals all conduct heat well
• because of the close contact of the atoms.
– E.g. pots/pans.
• Electrical Conductivity - metals all conduct
electricity when solid and when molten because
electrons can travel easily through the structure.
– E.g. cables
• Malleability - metals can be beaten into
different shapes.
– E.g. jewellery.
• Strength - most metals are strong
because of the metallic bond which
holds the atoms together.
– E.g. bridges, cars, buildings etc.
• Recycling Metals - Metals need to be
recycled because they will not last forever
(they are finite resources).
Alloys
• The properties of metals can be
extended or altered by mixing them
with other metals or with non-metals.
• Iron can be changed into stainless steel
by mixing it with small amounts of
chromium. This stops the metal rusting.
Alloy
Main
Metal
Other
Elements
present
Uses
Reason
Stainless
steel
Iron
Chromium,
Nickel
Sinks, Cutlery
Non-rusting,
strong
Mild
steel
Iron
Carbon
Girders, Car
bodies
Strong, rust
resistant
Gold
Gold
Copper
Rings, Electrical
contacts
Good conductor,
unreactive
Solder
Lead
(50%)
Tin (50%)
Joining metals,
electrical
contacts
Low melting
point, good
conductor
Brass
Copper
Zinc
Machine bearings,
ornaments
Hard wearing,
attractive
Reactions of Metals
METAL REACTIVITY
The reactions of metals that we will cover are;
•
reaction with oxygen
metal + oxygen  metal oxide
• reaction with water
metal + water  metal hydroxide + hydrogen
•
reaction with dilute acid
Metal + acid  salt + hydrogen
Reactivity Series
• Metals have similar chemical properties.
However, some metals are more reactive
than others.
• Based on their reactivity, chemists produced
a ‘league table’ of metals as shown below.
Name
Symbol
Potassium
K
Sodium
Na
Lithium
Li
Calcium
Ca
Magnesium
Mg
Aluminium
Al
Zinc
Zn
Iron
Fe
Tin
Sn
Lead
Pb
Copper
Cu
Mercury
Hg
Silver
Au
Gold
Ag
most reactive
least reactive
Metals Reacting with Oxygen
• All metals above silver in the reactivity
series react with oxygen when heated to
form a metal oxide.
• The higher the metal in the reactivity
series the more vigorous the reaction
with oxygen.
METAL
+ OXYGEN
e.g.
magnesium + oxygen 
2Mg(s) + O2(g) 
METAL OXIDE
magnesium oxide
2MgO(s)
• Potassium, sodium and lithium are so
reactive they are stored under oil to
prevent them from reacting with the
oxygen and water in the air.
• Oxygen can be made by heating potassium
permanganate in a test tube and allowing the gas
to pass through the preheated metal.
E.g.
Metal + Oxygen  Metal oxide
Magnesium + Oxygen  Magnesium oxide
Mg
+
O2

MgO
Metals Reacting with Water
• All metals above aluminium in the reactivity
series react with water to produce the metal
hydroxide and hydrogen gas:
METAL + WATER
METAL HYDROXIDE + HYDROGEN
e.g.
sodium + water  sodium hydroxide + hydrogen
Na(s) + H2O(l) 
NaOH(aq) + H2(g)
Metals Reacting with Acids
• All metals above copper in the reactivity
series react with dilute acids such as
hydrochloric and sulphuric acid to produce
a salt and hydrogen gas:
METAL + ACID
SALT +HYDROGEN
e.g.
zinc + hydrochloric acid zinc chloride + hydrogen
Zn(s) +
2HCl(aq)  ZnCl2(aq) +
H2(g)
• When a metal reacts with an acid it
produces bubbles of hydrogen gas.
• Generally, the faster the bubbles are
produced, the more reactive the metal.
• Aluminium is the exception to this. It
reacts very slowly for the first 20 mins,
after which it reacts quickly.
• The reason for this is that the metal is
protected by a thin layer of aluminium
oxide, which must first be removed by the
acid.
Summary
Metal Ores
• Ores are naturally-occuring compounds
of metals from which metals can be
extracted.
• The three main types of ore are metal
carbonates, metal oxide and metal
sulphides.
Common Ores
Common
name
Chemical name
Metal
present
Haematite
Iron oxide
Iron
Bauxite
Aluminium oxide
Aluminium
Galena
Lead sulphide
Lead
Cinnabar
Mercury sulphide
Mercury
Malachite
Copper(II)
carbonate
Copper
Percentage Composition
Extracting Metals
• Metals such as gold and silver occur
uncombined on earth because they are
unreactive and because of this these
elements were among the first to be
discovered.
• Other metals, such as those in the table
are found in compounds and have to be
extracted (which is an example of
reduction).
Extraction of Metals from
Ores
• The method used to extract a metal
depends on the reactivity of the metal.
– The more reactive the metal, the more
difficult it is to extract.
– The less reactive the metal, the easier it is
to extract.
Methods of extraction
a) Heating metal oxides
Silver oxide  Silver + Oxygen
Ag2O

Ag +
O2
• Few metals can be obtained in this way.
b) Heating Metal Oxides with Carbon
Metal oxide + Carbon  Metal + Carbon dioxide
E.g.
Iron oxide + Carbon  Iron + Carbon dioxide
Fe2O3
+
C
 Fe
+
CO2
• This method is used to extract metals
below aluminium in the reactivity series.
c) Using Electricity
• Electricity can be used to split ionic
compounds into their elements in a
process called electrolysis.
• The method is used to extract reactive
metals above zinc in the reactivity
series.
• A large electric current is passed
through the molten compound, and metal
appears at the negative electrode.
Electrolysis
• http://www.youtube.com/watch?v=i9xS9t-KMpc –
electrolysis explained
Electrochemical Series
Potassium
Sodium
Calcium
Magnesium
Aluminium
Zinc
Iron
Nickel
Tin
Lead
Copper
Mercury
Silver
Gold
Must be electrolysed to release
metal from ore
Separated from ore by heating with
CHARCOAL, thus releasing CARBON
DIOXIDE
Can be broken by heat alone
Batteries and Cells
• We generate electricity from burning
fossil fuels, harnessing the power of
water (hydroelectric), or nuclear energy.
• But, we also need electricity for personal
stereos, mobile phones etc.
• We use batteries.
Chemical reactions in a battery produce electricity
• When electricity is produced in a battery,
electrons flow from the battery, through
the wires, to the device to which it is
connected.
• In most batteries the electrons come from
a layer of zinc metal.
Electricity is a flow of electrons.
Dry-cell Battery
• Zinc cup forms the negative terminal of the
battery and the carbon rod is the positive
terminal.
• Between the two terminals is a paste of
ammonium chloride. This completes the
circuit by allowing ions to flow through it –
acts as an electrolyte.
An electrolyte is a substance that will
conduct electricity when dissolved in
water or melted. This is due to the
movement of ions.
• Some batteries are re-chargable. The
chemicals can be restored by giving the
battery a supply of electrons.
• e.g. a car battery contains lead metal.
When the battery is being used the lead
metal atoms turn into ions. During
recharging, the ions are turned back into
lead atoms.
Simple Cells
•
Electricity can be produced by
connecting different metals together,
with an electrolyte, to form a simple cell.
•
In the cell shown above, electrons flow
from the zinc to the copper. The sodium
chloride solution acts as an electrolyte
and completes the circuit.
• A voltmeter measures the voltage produced
and it is seen that different voltages are
obtained when different metals are used.
• The voltage between different pairs of
metals varies and this leads to the
Electrochemical Series (ECS) (page 7 of
Data Booklet).
• When two different metals are joined
together, electrons flow through the wire
from the metal higher in the ECS series to
the metal lower in the series e.g. from
metal A
lithium to silver.
V
filter paper soaked in a
sodium chloride solution
metal B
•The further apart the metals are in the ECS,
the higher the voltage produced.
•The closer together the metals are in the
ECS, the lower the voltage produced.
Displacement Reactions
•
Displacement reactions occur when a metal is
added to a solution containing ions of a metal
lower in the electrochemical series.
Example
•
If zinc metal is added to a solution of
copper(II) sulphate, the zinc slowly becomes
smaller and a brown solid covers it. At the
same time the blue copper(II) sulphate
solution loses its colour.
Why does this happen?
• The zinc atoms have LOST electrons and
turned into zinc ions, which go into
solution.
Zns   Zn 2 aq   2e 
atoms
ions
electrons
• The copper ions GAIN the electrons lost
by the zinc and turns into copper metal
atoms.
2


Cu
aq   2e

ions electrons
Cus 
atoms
• This is called a DISPLACEMENT
REACTION and the overall reaction can be
represented by;
Zns   Cu
2
aq
 Zn
2
aq
 Cus 
DISPLACEMENT REACTION: Formation
of a metal from a solution containing its
own ions when a metal higher than itself in
the electrochemical series is added to it.
• As a general rule, a metal will displace a
metal lower than itself in the ECS.
• e.g.
• - iron would displace silver ions from a
solution of silver nitrate as iron is above
silver in the ECS.
• - lead would not displace tin ions from a
solution of tin chloride as lead is lower
than tin in the ECS.
Hydrogen in the ECS
• Hydrogen and other non-metals are also in
the ECS.
• Hydrogen can be placed in the ECS by
considering the reactions of metals with
dilute acids.
• Metals down to lead in the ECS react with
dilute acids to produce hydrogen gas, i.e.
they displace hydrogen ions from acids.
• Copper, silver, gold and platinum do not
react with dilute acids.
• So hydrogen can be placed below lead but
above copper in the ECS.
Half-Cells
• Cells can also be set up by connecting twohalf cells together.
• A half-cell consists of a metal in contact
with a solution of its own ions, such as a
strip of copper metal in a beaker of
copper(II) sulphate solution.
• Electricity is produced when two half-cells
containing different metals are connected
as shown:
• The metals are joined by wires (electrons
flow) and the two solutions are connected
using an ion bridge (ions flow). Filter paper
soaked in sodium chloride solution is often
used.
• The ion bridge completes the circuit and
allows ions to move across it. If it is removed,
the circuit will be broken and no electricity
will be produced.
• In the cell shown the zinc atoms lose
electrons and form zinc ions, while
the copper(II) ions gain electrons to
form atoms of copper metal.
• The ion-electron equations to represent
these reactions are;
Zn s   Zn2 aq  2e 
Cu
2
 aq
 2e  Cus 

• Zinc metal would turn into zinc ions and
the copper(II) ions would decrease until
the cell would stop producing electricity.
Cells Involving Non-Metals
• In the cell shown, the two half cells are a
solution of iodide ions and a solution of iron
(III) ions.
2I-  I2 + 2e-
Fe3+ + e-  Fe2+
• Electrons flow from the iodide ions through
the meter to the iron (III) ions. As this
happens, the iodide ions turn into iodine
molecules.
2I

aq
 I2 g   2e

• The iron (III) ions gain electrons and turn
into iron (II) ions.
Fe
3
aq
 e  Fe

2
aq
Oxidation
and
Reduction
Oxidation
• Below are the ionic equations for the
reactions of calcium with oxygen, water and
dilute acid.
2
2
2Ca  O2  2Ca O
Ca  2H2O  Ca2 OH  2  H2
Ca  2HCl  Ca2 Cl 2  H2
• Calcium atoms have lost electrons to
become calcium ions as shown below
Ca  Ca2  2e 
• This is an OXIDATION reaction.
An oxidation reaction is one in which
there is a loss of electrons.
Reduction
• Reactions in which there is a gain of
electrons are called REDUCTION
reactions.
2H  2e  H2

Ca
2

 2e  Ca

A reduction reaction is one in which
there is a gain of electrons.
Oxidation
Is
Loss of
electrons
Reduction
Is
Gain of
electrons
REDOX Reactions
• Oxidation and reduction reactions take
place at the same time.
• In a redox reaction, electrons lost by one
substance during oxidation are gained by
another substance during reduction.
• The formation of a compound by a metal is
a redox reaction.
e.g. when sodium joins with chlorine to form
sodium chloride
2Na  2Na   2e 
OXIDATION
Cl2  2e   2Cl
REDUCTION
• All displacement reactions are redox
reactions
• e.g. the reaction between zinc metal and
copper(II) sulphate solution.
Zn  Zn2  2e 
OXIDATION
Cu2  2e   Cu
REDUCTION
• The oxidation and reduction equations can
be combined to show the overall REDOX
reaction.
Zn(s)  Cu2(aq)  2e -  Zn2(aq)  Cu(s)  2e Zn  Cu2  Zn2  Cu
OVERALL REACTION
• The reaction between a metal and a dilute acid
can also be considered to be a redox reaction.
e.g. The reaction between magnesium and
hydrochloric acid.
• The magnesium is oxidised and forms
magnesium ions, whilst the hydrogen ions in the
acid gain electrons and form hydrogen gas.
Mg  Mg2  2e 
OXIDATION
2H  2e   H2
REDUCTION
Mg  2H  Mg2  H2
OVERALL REACTION
• During electrolysis, oxidation occurs at
the positive electrode and reduction
occurs at the negative electrode.
Hydrogen fuel cell
How Fuel Cells work
• http://www.youtube.com/watch?v=Tk_iIzOUjTU
•
• http://www.youtube.com/watch?v=c3PkgUcI4Z8
•
• http://www.youtube.com/watch?v=5dSYQf8TUhA
• Advantages and disadvantages of hydrogen as a fuel.
• http://www.youtube.com/watch?v=mr2_XqRZjn0
What is a Fuel Cell?
Quite simply, a fuel cell is a device that converts chemical energy into
electrical energy, water, and heat through electrochemical reactions.
Fuel and air react when they come
into contact through a porous
membrane (electrolyte) which separates
them.
This reaction results in a transfer of
electrons and ions across the electrolyte
from the anode to the cathode.
 If an external load is attached to this
arrangement, a complete circuit is formed
and a voltage is generated from the flow
of electrical current.
The voltage generated by a single cell is typically rather small (< 1 volt), so man
cells are connected in series to create a useful voltage.
Fuel Cell Vs. Battery
Basic operating principles of both are very similar, but there are several
intrinsic differences.
Hydrogen fuel cell
Galvanic cell (battery)
Open system
Anode and cathode are gases in
contact with a platinum catalyst.
Reactants are externally supplied,
no recharging required.
 Closed system
Anode and cathode are metals.
Reactants are internally consumed,
need periodic recharging.
Fuel Cell Vs. Internal Combustion Engine
Similarities:
 Both use hydrogen-rich fuel.
 Both use compressed air as the oxidant.
 Both require cooling.
Differences:
Fuel cell:
 Output is electrical work.
 Fuel and oxidant react
electrochemically.
 Little to no pollution produced.
I.C. Engine:
 Output is mechanical work.
 Fuel and oxidant react combustively.
 Use of fossil fuels can produce significant pollution.
Some History…
Fuel cell principle first discovered
by William Grove in 1839.
Grove used four large cells, each
containing hydrogen and oxygen,
to produce electric power which was
then used to split the water in the
smaller upper cell.
Commercial potential first demonstrated by NASA in the 1960’s with the
usage of fuel cells on the Gemini and Apollo space flights. However, these
fuel cells were very expensive.
Fuel cell research and development has been actively taking place since the
1970’s, resulting in many commercial applications ranging from low cost portab
systems for cell phones and laptops to large power systems for buildings.
Fuel Cells in Use: Stationary Systems
Fuel Cells in Use: Stationary Systems
Fuel cell system for submarine
Fuel Cells in Use: Transportation Systems
Buses are most commercially
advanced applications of fuel
cells to date.
Are currently being used by
many American and European
cities.
XCELLSiS fuel cell bus prototypes
Fuel Cells in Use: Transportation Systems
Many of the major car companies are developing fuel cell car prototypes
which should come to market during the next decade. The cars use either
pure hydrogen or methanol with an on board reformer.
Fuel Cells in Use: Hydrogen Fuel Cell System
Fuel Cells in Use: Space Systems
12 kW Space shuttle fuel cell
Weight: 120 kg
Size: 36x38x114 cm
Contains 32 cells in series
1.5 kW Apollo fuel cell
Apollo used two of these
units.
Fuel Cells in Use: Portable Systems
A laptop using a fuel cell power source
can operate for up to 20 hours on a
single charge of fuel (Courtesy: Ballard
Power Systems)
But Isn’t Hydrogen Explosive?
Many have blamed this disaster on
a hydrogen explosion. However,
hydrogen burns invisibly, and no
evidence of leaks were ever found
(garlic scent added to hydrogen gas).
Using infrared spectrographs, NASA
scientists found that the skin of the
Hindenburg was treated with compound
which are found in gunpowder and
rocket fuel (nitrates and aluminum
powder). This, combined with a wooden
frame coated with lacquer resulted in
a highly flammable ship.
Glossary of Terms Used in Describing Fuel Cell
Technology
Electrochemical reaction: A reaction involving the transfer of electrons
from one chemical substance to another.
Electrode: An electrical terminal that conducts an electric current into
or
out of a fuel cell (where the electrochemical reaction occurs).
Anode: Electrode where oxidation reaction happens (electrons are
released).
Cathode: Electrode where reduction reaction occurs (electrons are
acquired).
In a fuel cell, hydrogen is oxidized at the anode and oxygen reduction
occurs
at the cathode.
Electrolyte: A chemical compound that conducts ions from one
electrode to
the other.
Terminology (cont.)
Catalyst: A substance that participates in a reaction, increasing its rate,
but is not consumed in the reaction.
Polymer: A natural or synthetic compound made of giant molecules which
are composed of repeated links of simple molecules (monomers).
Inverter: A device used to convert direct current electricity produced by a
fuel cell to alternating current.
Reformer: A device that extracts pure hydrogen from hydrocarbons.
Stack: Individual fuel cells connected in series within a generating assembly.