NATIONAL 5 CHEMISTRY

UNIT 3 CHEMISTRY IN SOCIETY

CONTENTS

•Metals •Properties of plastics•Fertilisers •Nuclear Chemistry •Chemical Analysis

METALSMETALS•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 alow density material e.g. aluminium (low density) – used to buildaircraft. Lead (high density) – is used as weights for fishingnets/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 Element

s present

Uses Reason

Stainless steel

Iron Chromiu

m, 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 Coppe

r Zinc

Machine bearings,

ornaments

Hard wearing, attractive

Reactions of MetalsReactions of Metals

The reactions of metals that we will cover are;

• reaction with oxygenmetal + oxygen metal oxide

• reaction with water

metal + water metal hydroxide + hydrogen

• reaction with dilute acidMetal + acid salt + hydrogen

METAL REACTIVITY

Reactivity SeriesReactivity 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 OxygenMetals 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.

e.g.magnesium + oxygen magnesium oxide

2Mg(s) + O2(g) 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.

METAL + OXYGEN METAL OXIDE

• Oxygen can be made by heating potassium permanganate in a test tube and allowing the gas to pass through the preheated metal.

Metal + Oxygen Metal oxide E.g.

Magnesium + Oxygen Magnesium oxide Mg + O2 MgO

Metals Reacting with WaterMetals Reacting with Water

• All metals above aluminium in the reactivity series react with water to produce the metal hydroxide and hydrogen gas:

e.g.sodium + water sodium hydroxide + hydrogen

Na(s) + H2O(l) NaOH(aq) + H2(g)

METAL + WATER METAL HYDROXIDE + HYDROGEN

Metals Reacting with AcidsMetals 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

PotassiumSodium CalciumMagnesiumAluminiumZincIronNickelTinLeadCopperMercurySilverGold

Can be broken by heat alone

Separated from ore by heating with

CHARCOAL, thus releasing CARBON

DIOXIDE

Must be electrolysed to release

metal from ore

Batteries and CellsBatteries 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 BatteryDry-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 CellsSimple 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 lithium to silver.

metal B

V

metal A

filter paper soaked in a sodium chloride solution

•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 ReactionsDisplacement 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?Why does this happen?

• The zinc atoms have LOST electrons and turned into zinc ions, which go into solution.

• The copper ions GAIN the electrons lost by the zinc and turns into copper metal atoms.

electrons ions

2 2

atoms

eZnZn aqs

• This is called a DISPLACEMENT REACTION and the overall reaction can be represented by;

s2aqaq

2s Cu Zn Cu Zn

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 ECSHydrogen 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-CellsHalf-Cells

• Cells can also be set up by connecting

two-half 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;

• Zinc metal would turn into zinc ions and the copper(II) ions would decrease until the cell would stop producing electricity.

saq2

aq2

s

Cue2Cu

e2ZnZn

Cells Involving Non-MetalsCells 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.

• The iron (III) ions gain electrons and turn into iron (II) ions.

e2 II2 g2aq

2

aq3aq FeeFe

OxidationOxidationandand

ReductionReduction

OxidationOxidation• Below are the ionic equations for the

reactions of calcium with oxygen, water and dilute acid.

• Calcium atoms have lost electrons to become calcium ions as shown below

• This is an OXIDATION reaction.

22

2

222

2

222

HClCaHCl2Ca

HOHCaOH2Ca

OCa2OCa2

e2CaCa 2

An oxidation reaction is one in which there is a loss of electrons.

ReductionReduction• Reactions in which there is a gain of

electrons are called REDUCTION reactions.

Cae2Ca

He2H2 2

2

A reduction reaction is one in which there is a gain of electrons.

Oxidation

IsLoss of electrons

Reduction

Is

Gain of electrons

REDOX ReactionsREDOX 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

• All displacement reactions are redox reactions

• e.g. the reaction between zinc metal and copper(II) sulphate solution.

REDUCTION Cl2e2Cl

OXIDATION e2Na2Na2

2

OXIDATION

REDUCTION

2

2

Zn Zn 2e

Cu 2e Cu

• The oxidation and reduction equations can be combined to show the overall REDOX reaction.

OVERALL REACTI ON

2 - 2 -(s) (aq) (aq) (s)

2 2

Zn Cu 2e Zn Cu 2e

Zn Cu Zn Cu

• 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.

OXI DATI ON

REDUCTI ON

OVERALL REACTI ON

2

2

22

Mg Mg 2e

2H 2e H

Mg 2H Mg H

• 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 intoelectrical 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 manycells are connected in series to create a useful voltage.

Galvanic cell (battery)Hydrogen fuel cell

Open systemAnode and cathode are gases in contact with a platinum catalyst.Reactants are externally supplied, no recharging required.

Closed systemAnode and cathode are metals.Reactants are internally consumed, need periodic recharging.

Fuel Cell Vs. Battery

Basic operating principles of both are very similar, but there are severalintrinsic differences.

Fuel Cell Vs. Internal Combustion Engine

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.

Differences:

Similarities:

Both use hydrogen-rich fuel. Both use compressed air as the oxidant. Both require cooling.

Some History…

Fuel cell principle first discoveredby William Grove in 1839.

Grove used four large cells, eachcontaining hydrogen and oxygen,to produce electric power which wasthen used to split the water in thesmaller upper cell.

Commercial potential first demonstrated by NASA in the 1960’s with theusage of fuel cells on the Gemini and Apollo space flights. However, thesefuel cells were very expensive.

Fuel cell research and development has been actively taking place since the1970’s, resulting in many commercial applications ranging from low cost portablesystems 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

XCELLSiS fuel cell bus prototypes

Buses are most commerciallyadvanced applications of fuelcells to date.

Are currently being used by many American and Europeancities.

Fuel Cells in Use: Transportation Systems

Many of the major car companies are developing fuel cell car prototypeswhich should come to market during the next decade. The cars use eitherpure hydrogen or methanol with an on board reformer.

Fuel Cells in Use: Hydrogen Fuel Cell System

Fuel Cells in Use: Space Systems

1.5 kW Apollo fuel cellApollo used two of theseunits.

12 kW Space shuttle fuel cellWeight: 120 kgSize: 36x38x114 cmContains 32 cells in series

Fuel Cells in Use: Portable Systems

A laptop using a fuel cell power sourcecan 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 ona hydrogen explosion. However,hydrogen burns invisibly, and noevidence of leaks were ever found(garlic scent added to hydrogen gas).

Using infrared spectrographs, NASAscientists found that the skin of theHindenburg was treated with compoundswhich are found in gunpowder androcket fuel (nitrates and aluminum powder). This, combined with a woodenframe coated with lacquer resulted in a highly flammable ship.

redox

Glossary of Terms Used in Describing Fuel CellTechnology

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.

Ion: An atom that carries a positive or negative charge due to the loss or gain of an electron. Anion is a negative ion, cation is a positive ion.

An electrochemical cell consists of 2 electrodes + 1 electrolyte

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.