csec chem metals chemistry of gardening etc

5/28/2018 CSEC Chem Metals Chemistry of Gardening Etc

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Metals

Metals always form positive ions

Extremes in Metals Metallic properties

Lightest: LithiumHeaviest: OsmiumMost brittle: Manganese and chromium

Lowest melting point: MercuryHighest melting point: TungstenMost expensive: PlatinumRarest: RhodiumMost abundant: Aluminium

Metals have high:- density: high mass per unit volume- tensile strength: high strength of the metal under

stress- durability: resistant to corrosion- malleability: ability to be made into sheets- ductility: ability to be made into wires- thermal conductivity: ability to conduct heat- electrical conductivity: ability to conduct electricity- sonority: ability to produce sound when struck

Alloys

a mixture of metallic elements or metallic with non-metallic. Pure metals are weak as the layers of atoms slide over each other easily. In alloy of 2 metals, they

have different sizes of atoms so this disrupts the orderly layer of atoms making it difficult for atoms toslide over.

Eg of alloys Steel: iron and carbon bronze: copper and tin brass: copper and zinc duralumin: aluminium, copper, magnesium

Uses of duralumin: it is light but strong and durable so used for aircraft parts,greenhouse frames, overhead cables, curtain walling in high-rise buildings etc.

pewter: tin and lead Uses of solder: mixture of tin and lead, has a much lower melting point than either of its components

so more easily fusible --- suitable for welding electrical wire together Uses of stainless steel: is an alloy of iron containing chromium or nickel. Is the most expensive way

applications for: cutleries medical instruments kitchen sinks steel objects in chemical factories and oil refineries

The Reactivity Series of Metals

is an arrangement of the metals (andcarbon and hydrogen) in order of their

reactivity. The reactivity of a metal is determined by

its ability to form a positive ion.

Reactions of Metals:The reactivity series of metals was deduced byperforming several experiments in the labwhich enabled scientists to arrange metalsaccording to their reactivity with dilute acid,oxygen (air), and water.

Reactive metals tend to form positiveions easily, by losing electrons and

forming compounds unreactive metals prefer to remain in

uncombined form, as the element itself the order of reactivity is worked out

from the metal's reaction (if any) withwater or steam and acids

if there is a reaction, the metaldisplaces hydrogen

Metal + hydrogen ion metal ion + hydrogengas

Elements Symbols

Potassium KSodium Na

Calcium CaMagnesium Mg

Aluminium AlCarbonnot a metal C

Zinc Zn

Iron FeTin Sn

Lead PbHydrogen - not ametal

H

Copper Cu

Mercury Hg

Silver Ag

Platinum Pt

Gold Au

N.B: Some text books list Gold before PlatinumThere are mo re metals in the reactiv i ty seriesbut they arenot studied at this level.

Heres a jingle to help you remember the order of the series:Please Send Chief Minister ACute Zebra In The Largest Heaviest Case Marked Striped Perishable Goods

Two non-metals, carbon andhydrogen, are included in the table for comparison, and are importantchemical reference points concerning the method of metal extraction and reactivity towards acids

ReactivityDecrea

ses

Most reactive

Least reactive

CSEC Chemistry 11Sc/TV1


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o Metals above carbon cannot usually be extracted by carbon or carbon monoxide reduction andare usually extracted by electrolysis.In sense this means metals above carbon in the reactivity seriescannot be 'displaced' from their compounds by carbon.

o Metals below carbon in the series can be extracted by heating the oxide with carbon or carbonmonoxide.

o Metals below hydrogen will not usually displace hydrogen from acidsand can be extracted byheating the oxide in hydrogen, though is rarely done e.g. for cost (not as cheap as coke/carbon) andsafety reasons (hydrogen very explosive in air). Again, you can think of metals above hydrogen in thereactivity series as being reactive enough to displace hydrogen from acids in aqueous solution.

Metals were arranged in order of reactivity starting with the most reactive metal which had the highest rate of

effervescence of hydrogen gas. The rate of effervescence is also the rate of this reaction is measured by measuring the

volume of hydrogen produced per unit time.

Metals Reactivity with Dilute HCl

Potassium, Sodium & Calcium React extremely violently with rapid effervescence and splashing

Magnesium & Aluminum React violently with rapid effervescence

Zinc, Iron & Lead React slowly with bubbles

Copper, Silver, Gold & Platinum Do not react

Metal Metal with water/steam Metal with acid

potassiumsodiumcalcium

react with cold waterM(s) + 2H2O(l) --> MOH(aq) + H2(g)Metal + Water --> Metal Hydroxide + Hydrogen

violent reaction with dilute acidsM(s) + 2HCl(aq) --> MCl2(aq) + H2(g)Metal + Acid --> Metal Chloride + Hydrogen

magnesiumaluminiumzinciron

react with steamM(s) + 2H2O(g) --> MO(s) + H2(g)Metal + Water --> Metal Oxide + Hydrogen

react with dilute acids with decreasing easeM(s) + 2HCl(aq) --> MCl2(aq) + H2(g)Metal + Acid --> Metal Chloride + Hydrogen

leadhydrogen

do not react with water or steam react with dilute acids with decreasing ease

coppermercurysilverplatinum

do not react with water or steam react only with concentrated acids

Reactions with Oxygen in Air:

Most metals react with oxygen from air forming a metal oxide. You have previously studied that metal oxides

are basic oxides and that some of them are insoluble in water and some of them are soluble in water forming

an alkaline solution. The most reactive metals like potassium, sodium, calcium and magnesium react with

oxygen with a very bright flame and producing white ashes and their oxides are soluble. Moderately reactive

metals like aluminum and zinc react with oxygen forming white powdered ashes but their oxides are insoluble.

Iron and copper react very slowly with oxygen. The result of iron oxygen reactions is rust which is reddish

brown iron oxide. When a copper lump reacts with oxygen, a white layer of black copper oxide forms on it.

When the lump gets covered by this layer; the reaction stops. Oxides of iron and copper are insoluble. Metals

that are less reactive than copper like silver, gold and platinum do not react with oxygen.

Note: When aluminum reacts with oxygen, a layer of aluminum oxide adheres and covers the aluminum. At this

point no further reaction can take place.

Metal Observation Inference Equation

Magnesium

(Mg)

Burns vigorously with a very brilliant white

flame. The residue is white when hot and

cold.

The reactivity of Mg

towards O2is very high.

Magnesium oxide is

formed.

2Mg(s) + O2(g)>

2MgO(s)

Zinc (Zn) Burns quicklywith a bright flame. The

residue is yellow when hot and white when

cold.

The reactivity of Zn

towards O2is high. Zinc

oxide is formed.

2Zn(s) + O2(g)>

2ZnO(s)

Iron (Fe) Glowsvery brightly. The residue is

reddish-brown when hot and cold.

The reactivity of Fe

towards O2is medium.

Iron (III) oxide is

formed.

2Fe(s) + O2(g)>

2Fe2O3(s)

Lead (Pb) Glows brightly. The residue is brown when

hot and yellow when cold.

The reactivity of Pb

towards O2is low. Lead

(II) oxide is formed.

2Pb(s) + O2(g)>

2PbO(s)


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Copper

(Cu)

Glows faintly. The residue is black when

hot and cold.

The reactivity of Cu

towards O2is very low.

Copper (II) oxide is

formed.

2Cu(s) + O2(g)>

2CuO(s)

Question: Predict the reactivity of the other metals in the reactivity series with oxygen.

Metal: Reaction:Order ofreactivity:

Products:

PotassiumViolent reaction with cold water. Floats andcatches fire.

1st- mostreactive.

Potassium hydroxide, KOH andhydrogen gas.

SodiumVery vigorous reaction with cold water. Floats.Can be lit with a lighted splint.

2nd.Sodium hydroxide, NaOH andhydrogen gas.

Calcium Less vigorous with cold water. 3rd.Calcium hydroxide, Ca(OH)2andhydrogen gas.

MagnesiumVery slow with cold water, but vigorous withsteam.

4th.Magnesium oxide, MgO andhydrogen gas.

Zinc Quite slow with steam. 5th.Zinc oxide, ZnO and hydrogengas.

Iron Slow with steam. 6th

.Iron oxide, Fe203and hydrogengas.

Copper No reaction with steam.7th- leastreactive.

Competition Reactions in Solid State:

Previously youve studied displacement reactions which are pre-formed in aqueous states. A very similar

reaction takes place in the solid state, it is called thermite reaction. This reaction is used to repair damaged

railway lines. In this reaction, aluminum and iron (III) oxide are the reactants. In the reaction, aluminum

removes the oxygen ion from iron and bonds with it. This happens because aluminum is more reactive than

iron. The products are aluminum oxide and iron in molten form. In the fixing procedure, the reactants are put in

the cut in the railway line and the reaction is triggered by heating using a magnesium fuse. The reaction leaves

aluminum oxide and molten iron with then condenses in the cut welding it. Like displacement reactions, thisreaction is exothermic.

2Al + Fe2O3Al2O3+2Fe

Competition Reactions in Aqueous State:

These are ordinary displacement reactions in which the two positive ions compete for the negative ion. The ion

of the more reactive metal wins. Zinc is higher than copper in the reactivity series. If zinc is added to a solution

of copper nitrate, a displacement reaction will take place in which the zinc will displace the copper ion from the

solution in its salt. The products of this reaction are zinc nitrate and copper. Copper salt solutions have a blue

color which fades away as the reaction proceeds because the concentration of the copper salt decreases. This

type of reaction also helped in confirming reactivity of metals since the more reactive metal displaces the less

reactive one.

Zn + Cu(NO3)2 Zn(NO3)2+ Cu

Stability of metal compounds(Action of Heat)

Compounds of metals high up in the reactivity series are stable and not easily decomposed by heating.

Compounds of metals low down in the series are unstable, and are often decomposed by heating, or are easilyreduced.

The oxides of metals above zinc in the series can only be reduced to the metal by using electrolysis.

At cathode, reduction occurs

Al3++ 3e----> Al

The oxides below can be reduced with reducing agents like carbon or hydrogen, except zinc oxide which cannotbe reduced by action of hydrogen

ZnO + C --> Zn + CO

CuO + H2--> Cu + H2O

Metal Oxide Hydroxide Carbonate Nitrate

potassiumsodium

electrolyticreduction

stable to heat stable to heat decompose to nitrite andoxygen

calciummagnesium

aluminium

electrolyticreduction

decompose to metaloxide and steam on

heating

decompose to metaloxide and carbon dioxide

gas on heating

decompose to metaloxide, nitrogen dioxide and

oxygen on heatingzinc

ironleadcopper

reduced byheating withcarbon

decompose to metaloxide and steam onheating

decompose to metaloxide and carbon dioxidegas on heating

decompose to metaloxide, nitrogen dioxide andoxygen on heating


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mercurysilverplatinum

reduced byheating alone

unstable, do not exist unstable, do not exist decompose to metal,oxygen and nitrogendioxide gas on heating

Down the series, reduction becomes easier because the metals prefer to exist as atoms, as opposed toions

For metal oxides like mercury(II) oxide, no reducing agent is needed - just heating alone 2HgO --> 2Hg + O2

Hydroxides of the metals calcium and below decompose to their corresponding oxide and give offsteam, on heating. This can be confirmed by using anhydrous copper(II) sulphate which turns white to

blue with steam Ca(OH)2---> CaO + H2O

Similarly, most carbonates, except sodium and potassium carbonates, undergo thermal decompositionagain to a metal oxide, but this time giving off carbon dioxide gas. This can be confirmed by bubblingthe gas through limewater, which turns milky with carbon dioxide

PbCO3---> PbO + CO2 Nitrates also decompose on heating, but the stable ones at the top of the series only decompose as far

as the nitrite (nitrite(III)), giving off oxygen gas. This can be identified by the gas relighting a glowingsplinter

2KNO3---> 2KNO2+ O2 The majority of nitrates decompose to the metal oxide, giving off brown fumes of nitrogen dioxide aswell as oxygen gas.

2Mg(NO3)2 ---> 2MgO + 4NO2+ O2 The unstable nitrates at the bottom of the reactivity series decompose all the way to the metal itself

2AgNO3---> 2Ag + 2NO2 + O2 Applying heat to a metal compound such as potassium nitrate will cause it to decompose into

potassium nitrite and oxygen. This is a thermal decomposition reaction.

Metal: Anion:

Nitrate (NO3-) Carbonate (CO32-) Hydroxide (OH-)PotassiumSodium

Metal Nitrate Metal nitrite +Oxygen

NO DECOMPOSITION

CalciumMagnesium

AluminumZincIronLeadCopper

Metal Nitrate Metal oxide +Nitrogen dioxide + Oxygen

Metal Carbonate Metaloxide + Carbon dioxide

Metal hydroxide Metaloxide + Hydrogen

SilverGold

Metal Nitrate Metal +Nitrogen dioxide + Oxygen

Metal Carbonate Metal +Carbon dioxide + Oxygen

-

Silver and gold hydroxides do not exist.

Ions of more reactive metals tend to hold on tightly to their anions and do not decompose easily this is why lots of heat

is needed.

Extraction of Metals from their Ores:

Most metals do not exist in nature as pure elements. Instead, they are found as naturally occurring compounds

called ores. Ores are naturally occurring minerals from which a metal can be extracted. Most ores are metals

oxide, carbonate or sulfide mixed with other impurities. The extraction of metal from ores begun long ago when

people started purifying iron from its iron oxide ore by reducing it using charcoal. This was possible because

carbon is more reactive than iron so it can reduce it take the oxygen ion from it. But then other metals were

discovered which were higher than carbon in the reactivity series. Those metals were not possibly extracted

from their ores until in the 19th century when a method of extracting them by electrolysis was invented. The

method extracting a metal depends on its reactivity.

Metals - in decreasing order of reactivity Reactivity

Extract by electrolysis

Extract by reaction with carbon or carbon monoxide

Extracted by various chemical reactions


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Metals from Rocks

Mineralselements/compounds that make up rocks Metal orerock containing metal

Extracting these metals

Metal ores are removed from ground.

The ores contain useful and unwanted materials. Unwanted materials are separated to obtainconcentrated mineral.

Metal is extracted from the mineral.

Occurrence of Metals

Metal ores are compounds, usually as: Metal oxidesmetal + oxygen, eg: Al2O3 Metal sulphidesmetal + sulphur, eg: HgS Metal carbonatesmetal + carbon + oxygen, eg: MgCO3

Least Reactiveeasiest to extract; extracted by physical methods

Less Rectiveharder to extract than least reactive; by blast furnace; usually occur as compounds of oxides orsulphides.

Most Reactivehardest to extractstrong bonds in compounds; by electrolysis decomposing compounds withelectricity.

Uses of Metals

The choice of metals over another depends on 3 factors:

Physical properties (e.g. melting point, strength, density, conductivity)

Chemical properties (e.g. resists corrosion)

Cost

Recycling Metals

There are many iron on the surface but copper and tin are seriously reducing.

High temperatures and pressures and greater depth increases hazards that prevent mining up to the lower part of

crust, although there are more metals further down Ways to conserve metals

Use alternative materials to replace the use of iron (e.g. use of plastic pipes instead of iron, use of glassbottles for soft drinks instead of aluminium)

Recycle unused metals by melting them to produce new blocks of clean metal
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Advantages of recycling metals

Recyling helps conserving metals, especially valuables such as gold and platinum.

E.g. used computer parts processed to extract gold used for electrical contacts of processors andmemory chips

Recycling saves the cost of extracting new metals

Recycling benefits environment, e.g. if there is a car wasteland, it causes eyesore

Disadvantages of recycling metals

Recycling metals can damage the environment by smelting process which sends a lot of fumes into the air

Cost to separate metals from waste is high. E.g. separating metals in alloys is hard

Transport costs for collecting scrap metal is high, e.g. trucks should be used

People are not interested in depositing their used materials in recycling bins

Extraction of Iron

The ore of iron is called haematite. It consists of 60% iron in form of Iron oxide (Fe2O3) with other impurities such as

silicon dioxide (SiO2).

Substances Products and Waste Materials

Iron is extracted from the iron ore haematite, Fe2O3

Iron is extracted from the oxide in a tower called a blast furnace

Several reactions take place before the iron is finally produced.

1. Oxygen in the air reacts with coke to give carbon dioxide:

2. The limestone breaks down to form carbon dioxide:

3. Carbon dioxide produced in 1 + 2 react with more coke to produce carbon monoxide:

4. The carbon monoxide reduces the iron in the ore to give molten iron:

5. The limestone from 2, reacts with the sand to form slag(calcium silicate):

CaO(s)+ SiO2(s) CaSiO3(l)

Both the slagand ironare drained from the bottom of the furnace .


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Steel

Iron made from blast furnace is not good as:

it contains impurities which makes it brittle (can break easily)

it cannot be bent or stretched

Most iron is converted into steel which is an alloyof iron and carbon with small amounts of other elements.

Advantages of steel:

it is strong and tough

it can be bent and stretched without shattering

Making Steel:

Impurities of iron is removed by blowing oxygen into molten iron to change the impurities into oxides. They arethen combined with CaO and removed as slag.

Carbon and other metals are added in certain amount to make steel.

Different Types of Steel:

There are many different forms of steel. Each has different components and properties and is used for different

purposes.

Steel Composition Properties Uses

Mild Steel 99.5% Iron

0.5% Carbon

Easily worked lost brittleness Car bodies

large structures

Machinery

Hard Steel 99% Iron

1% Carbon

Tough and brittle Cutting tools and chisels

Stainless Steel 87% Iron

13%

Manganese

Tough and springy Drill bits and springs and chemical

plants

ManganeseSteel

74% Iron18% Chromium

8% Nickel

Tough and resistant to corrosion Cutlery and surgical tools, kitchensinks

Tungsten Steel 95% Iron

5% Tungsten

Tough and hard even at high

temperatures

Edges of high speed cutting tools

Extraction of Aluminum:

Aluminum exists naturally as aluminum oxide (alumina) in its ore, which is called bauxite. Because

aluminum is a very reactive metal, it holds on very tightly to the anion it bonds with, which is oxide in

this case. This is why the best way to extract and purify aluminum is by electrolysisin a cell like theone below.

1. Mining of Bauxite:

The bauxite (red-brown solid) - aluminium oxide (Al2O3) mixed with impurities - is extractedfrom the earth.

2. Purification of Bauxite:

The extracted aluminium oxide is then treated with alkali (NaOH),to remove the impurities.This results in a white solid called aluminium oxide(or alumina) (Al2O3).

3. Electrolysis of Aluminium oxide:

1 Graphite blocks - anode

2 Graphite lining - cathode

3 alumina is dissolved in molten cryolite

4 Steel container

5 Outlet for molten aluminium


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o The alumina is then transported to huge tanks. The tanks are lined with graphite,thisacts as the cathode.Also blocks of graphite hang in the middle of the tank, and actsas anodes.

o Aluminas melting point (2050oC) is very high, therefore the alumina is dissolved

in molten cryolite- this lowers the melting point to 950 C - saves money!

o When molten, the aluminium ions and oxide ions in the alumina can move.

o Electricity is passed and electrolysis begins. Electrolysis is the decomposition of acompound using electricity.

o At the anode: The oxide ions lose electrons to become oxygen molecules, O2:

o At the cathode: Here the aluminium ions receive electrons to become atoms again:

Overall equation: 4Al3+(l) + 6O2-(l) 4Al(l) + 3O2(g)

Oxygen gas evolves and is collected with waste gases

Aluminum is sucked out of the container at regular intervals

Oxygen gas which evolves reacts with carbon from the cathode forming CO2. The cathodegets worn away. To solve this, the cathode is replaced at regular intervals.

Uses of Aluminium because

aircraft light, strong, resists corrosion

other transport such as ships' superstructures, containervehicle bodies, tube trains (metro trains)

light, strong, resists corrosion

overhead power cables (with a steel core to strengthen them)light, resists corrosion, good conductor ofelectricity

saucepanslight, resists corrosion, good appearance,good conductor of heat

THE MANUFACTURE OF SULPHURIC ACID - THE CONTACT PROCESS

1.Making the sulph ur dioxide

This can either be made by burning sulphur in an excess of air:

. . . ORby heating sulphide ores like pyrite in an excess of air:

2. Making the sulph ur tr ioxide

This is a reversible reaction, and the formation of the sulphur trioxide is exothermic.

Conditions:

o Temperature: 400-450C,

o Pressure: 12 atmospheres,

o Catalyst:vanadium pentoxide (V2O5)

3. Convert ing the sulphur tr ioxide into sulph ur ic acid

This can't be done by simply adding water to the sulphur trioxide - the reaction is so uncontrollablethat it creates a fog of sulphuric acid. Instead, the sulphur trioxide is dissolved in concentratedsulphuric acid to form fuming sulphuric acid (oleum).


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Water is then carefully added to the oleum to produce concentrated sulphuric acid (98%). - twice as

much as you originally used to make the fuming sulphuric acid.

MANUFACTURE OF AMMONIA -THE HABER PROCESS

Nitrogenneeded for this process is obtained from fractional distillation of liquid air.

Hydrogenneeded could be obtained by three ways, either reacting methane with steam,

electrolysis of brine or cracking of alkanes.

N2(g) + 3 H2(g) 2 NH3(g) -a reversible exothermic process

Nitrogen and hydrogenare mixed together in a 1 to 3 ratio with the reactionConditions:

o Temperature: 400 - 450Co Pressure: 200 atmosphere

o Catalyst: finely divided iron

This mixture is cooled to about -40C sothat ammonia condenses and liquifiedbutnitrogen and hydrogen dont. Ammonia iscollected and stored. The unreactednitrogen and hydrogen are then recycledand the ideal conditions are brought back infor it to react. This is repeated until no moreof the reactants remain.

Uses of Ammonia:

Manufacture of fertilisers

Manufacturing of nitricacid

MANUFACTURE of CHLORINE USING A DIAPHRAGM CELL

Background chemistryChlorine is manufactured by electrolysingconcentrated sodium chloride solution(brine). Three useful substances are made inthis process - chlorine, sodium hydroxide andhydrogen.

The chemistry of the electrolysis processSodium chloride solution contains:

o sodium ions,

o chloride ions,

o hydrogen ions (from the water),

o hydroxide ions (from the water).

The diaphragm cell

The diaphragm is made of a porous mixture ofasbestos and polymers. The solution can seep throughit from the anode compartment into the cathode side.

Notice that there is a higher level of liquid on the anodeside. That makes sure that the flow of liquid is alwaysfrom left to right - preventing any of the sodium

hydroxide solution formed finding its way back to wherechlorine is being produced.

Production of the chlorine

Chlorine is produced at the titanium anodeaccordingto the equation:

Production of the hydrogen

The hydrogen is produced at the steel cathode:

Production of the sodium hydroxide


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A dilute solution of sodium hydroxide solution is also produced at the cathode. It is highly contaminated withunchanged sodium chloride solution.

The sodium hydroxide solution leaving the cell is concentrated by evaporation. During this process, most of thesodium chloride crystallises out as solid salt. The salt can be separated, dissolved in water, and passed through thecell again.

Even after concentration, the sodium hydroxide will still contain a small percentage of sodium chloride.

Reactivity Series Involving Non-Metals

Metals react by losing electrons - they are reducing agents. Non-metalsreact by gaining

electrons - they are oxidising agents. In the same way that metals can be ordered in

terms of reducing strength, the non-metals can be ordered in terms of their oxidising

strength. The halogens are a typical example of a non-metal reactivity series.

Reactivity of the halogens

Fluorine most reactive

Chlorine

Bromine

Iodine least reactive

Fluorine is so reactive that we cannot isolate it in the

laboratory very easily, as it reacts with both water and glass.

As a result we don't usually deal with fluorine at pre-university

level but compare only the other three (astatine is very rare

and radioactive)

Do not confuse this order of reactivity with that of the metals

- these are non-metals, their reactivity is in terms of

oxidising power - i.e. chlorine is the best oxidising agent

out of chlorine, bromine and iodine.

1. Chlorine will displace bromine from solutions containing bromide ions

Cl2+ 2Br- Br2+ 2Cl-

In this reaction the chlorine is oxidising the bromide ions by removing an electron from them. Bromine isliberated from the solution and may be detected by its orange/red colour

2. Bromine will displace iodine from solutions containing iodide ions

Br2+ 2I- I2+ 2Br-

In this reaction the bromine is oxidising the iodide ions by removing an electron from them. Iodine is liberated

from the solution and may be detected by its orange/brown colour which turns blue/black in the presence ofstarch indicator.

It is predictable, then, that chlorine will also displace iodine from a solution containing iodide ions

Qualitative Analysis

Test for Gases

Gas Test and Test Results

Ammonia (NH3) turns damp red litmus paper blueCarbon dioxide (CO2) gives white ppt with limewater, ppt dissolves with excess CO2

Chlorine (Cl2) bleaches damp litmus paper

Hydrogen (H2) produces "pop" sound with lighted splint

Oxygen (O2) relights a glowing splint

Sulfur dioxide (SO2) turns aqueous acidified potassium dichromate (VI) from orange to green

Test for Anions

Anion Test Test result


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Carbonate (CO32-) Add dilute acid Effervescence, carbon dioxide produced

Chloride (Cl-)(in solution)

Acidify with dilute nitric acid, then add aqueous silvernitrate

White ppt

Iodide (I-)(in solution)

Acidify with dilute nitric acid, then add aqueous lead(II)nitrate

Yellow ppt

nitrate (NO3-)(in solution)

Add aqueous sodium hydroxide, then aluminium foil,warm carefully

Ammonia produced

Sulfate (SO42-

)(in solution) Acidify with dilute nitric acid, then add aqueous bariumnitrate White ppt

Colours of some metal hydroxides

Metal hydroxide Colour

calcium hydroxide white

copper(II) hydroxide light blue

iron(II) hydroxide green

iron(III) hydroxide red-brown

lead(II) hydroxide white

zinc hydroxide white

Testing for cations

Cations can be identified by their reactions with aqueous sodium hydroxideand aqueous ammonia

A precipitateis an insoluble solid.

When testing for cations, these precipitates only form when a metal ion in solution joins with hydroxide ions insolution to form an insoluble metal hydroxide

eg Cu2++ 2OH---> Cu(OH)2 (blue copper(II) hydroxide precipitate)

Cation Add dilute NaOH

(5 drops)

Add excess dilute

NaOH

Add dilute aqueous NH3

(5 drops)

Add excess

dilute aqueousNH3

Al3+ white ppt of aluminium hydroxide ppt dissolvescolourless solution

white ppt of aluminiumhydroxide

ppt insoluble

Ca2+ white ppt of calcium hydroxide ppt insoluble no reaction no reaction

Cu2+ blue ppt of copper(II) hydroxide ppt insoluble blue ppt of copper(II)hydroxide

ppt dissolvesdeep bluesolution

Fe2+ dirty green ppt of iron(II) hydroxide ppt insoluble dirty green ppt of iron(II)hydroxide

ppt insoluble

Fe3+ red-brown ppt of iron(III) hydroxide ppt insoluble red-brown ppt of iron(III)hydroxide

ppt insoluble

Pb2+ white ppt of lead(II) hydroxide ppt dissolvescolourless solution

white ppt of lead(II)hydroxide

ppt insoluble

Zn2+ white ppt of zinc hydroxide ppt dissolvescolourless solution

white ppt of zinchydroxide

ppt dissolvescolourlesssolution

NH4+ammonium

ammonia gas is produced on warmingwith dilute NaOH. This gas has apungent smell and turns moist redlitmus paper blue

- no reaction -

[Lead(II) ions can be distinguished from aluminium ions by the insolubility of lead(II) chloride.]

Laboratory Preparation of Gases


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Laboratory Preparation of Oxygen

Oxygen has similar density to air hence is collected

by water displacement

Laboratory Preparation of Ammonia Laboratory Preparation of Carbon Dioxide

A method for preparing and collecting a gas lessdense (lighter) than air by heating solid reactants.The less dense gas rises into, and displaces, the moredense air downwards. This method of gas preparation iscalled upward delivery.

e.g. Heating a mixture of ammonium chloride andcalcium hydroxide (slaked lime) solidsgives ammoniawhich has a very pungent odour! andturns red litmus blue.

2NH4Cl(s)+ Ca(OH)2(s)==> CaCl2(s)+ 2H2O(l)+ 2NH3(g)

To make dry ammoniayou need a tube packed withgranules of calcium oxide (quick lime)between thehorizontal pyrex tube and the vertical inverted collectiontest tube.

This method of gas preparation is called downwarddelivery.

(i)Calcium carbonate (limestone/marble chips) withhydrochloric acid makes carbon dioxide. carbon dioxideis moderately soluble in water

CaCO3(s)+ 2HCl(aq)==>CaCl2(aq)+ H2O(l)+ CO2(g)

Laboratory Preparation of HydrogenA method for preparing a sample of hydrogen gas

Method for preparing and collecting a gas less dense(lighter) than air by reacting a liquid and a solid . Theless dense gas rises into, and displaces, the more denseair downwards. This method of gas preparation iscalled upward delivery.

e.g. A mixture of zinc and hydrochloric acidmakes hydrogen. Hydrogen gives a squeaky pop! with alit splint.

Zn(s)+ 2HCl(aq)==>ZnCl2(aq)+ H2(g)

Drying agents: Gas used to dryo concentrated sulphuric acid

o anhydrous calcium chloride

Oxygen, hydrogen, chlorine

o silica gel Carbon dioxide

o calcium oxide ammonia


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MCQ Questions

1. Which statement about salts is not correct?a, salts are made by neutralising alkalis with acids

b. salts contain anions and cationsc. salts are made by dissolving metal oxides in acidsd. salts always contain water of crystallisation

2. An example of a salt which can be prepared byprecipitation is

a. lead(II) nitrateb. sodium carbonatec. silver chlorided. magnesium sulphate

3. A way of distinguishing dilute hydrochloric acidfrom dilute sulphuric acid is toa. add universal indicatorb. add aqueous barium nitratec. add a metal carbonated. add magnesium ribbon

4. Which of these statements about solubility is true?a. all sulphates are soluble in water except calcium andlead sulphateb. all nitrates are insoluble in water except sodium andpotassium nitratec. most metal oxides are soluble in water except those ofGroup I and IId. most metal carbonates are soluble in water

5. Barium sulphate is insoluble in water. It is used

in a 'barium meal' to allow X-ray studies of theintestines. It can be prepared by a precipitationreaction between two aqueous solutions. Whichtwo substances would be suitable for preparingbarium sulphate for use in X-ray radiography?

6. Which of these salts is best prepared by reaction

with an acid and a base?a. barium sulphateb. copper(II) carbonatec. magnesium sulphated. silver chloride


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a. barium carbonate and sulphuric acidb. barium chloride and sodium sulphatec. barium oxide and potassium sulphated. barium nitrate and calcium sulphate7. A solution of substance X gave a whiteprecipitate when aqueous NaOH was added.However, when lead(II) nitrate solution was addedto an acidified solution of X, a yellow precipitateformed. What is the correct identity of X?

a. calcium chlorideb. magnesium sulphatec. sodium bromided. zinc iodide

8. Iron(III) hydroxide is precipitated out of solutionwhen aqueous sodium hydroxide solution is added toiron(II) chloride solution.Fe3+(aq) + 3OH-(aq) --> Fe(OH)3(s)

What is the minimum volume of 2 mol/dm3

aqueousNaOH required to precipitate the maximum amount ofiron(III) hydroxide from 20cm3of 1 mol/dm3iron(III)chloride solution?a. 10 cm3b. 20 cm3c. 30 cm3d. 60 cm3

9. After acidification with dilute nitric acid, acolourless solution X reacts with aqueous silvernitrate to give a yellow precipitate. What could Xbe?a. calcium iodideb. copper(II) chloridec. iron(II) iodided. sodium chloride

10. An element reacts with steam but not with coldwater. Its oxide can be reduced by heating it withcarbon. When it is placed in a solution containingiron(II) ions, a grey deposit is formed. The element ismost likely to bea. leadb. magnesiumc. zincd. copper

MCQ Answers

1. d2. c3. b4. b5. b

6. c7. d8. c9. a10. c

Structured Questions Worked Solutions1. A student was given an aqueous solution analyse. Itcontains copper(II) chloride and aluminium nitrate.

a. Describe how he could detect the presence ofchloride ions in the aove solution.b. Name the precipitate(s) formed when excessaqueous ammonia is added to the above solution.

Solution: 1a.

Add dilute nitric acid followed by aqueous silver

nitrate.

A white precipitate will be seen in the presence of

chloride ions

OR

Add dilute nitric acid followed by aqueous lead(II)

nitrate.

A yellow precipitate will be seen in the presence of

chloride ions

2. W is an alkali and X is a salt. When the two solutions were mixed together, a reddish-brown precipitate Y was

obtained. When a salt Z was added to solution W and heated, a pungent gas which turned moist red litmus blue wasevolved. Suggest what could W, X, Y, and Z be.

Solution

W: Sodium hydroxideX: Iron(III) chloride/sulphate/etcY: Iron(III) hydroxideZ: Ammonium chloride/nitrate/etc

3. Give the name and formula of the ions present in each of the solutions X, Y and Z below -a. Solution X gives a white precipitate when dilute hydrochloric acid and aqueous barium chloride are added to it.b. An alkaline gas is given off when sodium hydroxide solution is added to the colourless solution Z and the mixtureheated.

Solution

3a. sulphate ion (SO42-)3b. Ammonium ion (NH4+)

4. A similar reagent is added to zinc carbonate and sample S to initiate both reactions A and B.

ai. Name the reagent(s) required for Reaction A.aii. Write down the chemical equation (with state symbols) for Reaction A.bi. Give a possible identity of Sample S.
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bii. Based on your answer in bi, write a chemical equation with state symbols for Reaction B.

Solution

ai. Hydrochloric acidaii. ZnCO3(s) + 2HCl (aq) ---> ZnCl2(aq) + CO2(g) + H2O (l)bi. Zincbii. Zn (s) + 2HCl (aq) ---> ZnCl2 (aq) + H2(g)

5.

a. Identify A to F.b. Write an ionic equation (with state symbols) for the formation of the yellow precipitate.c. Describe a test for the pungent gas R.d. Pungent gas R dissolves in water to form a solution. Describe the observations when a few drops of copper(II)sulphate solution is added to the solution.

Solution

a.A: lead(II) carbonateB: dilute nitric acidC: ammonia gasD: carbon dioxideE: lead(II) nitrateF: lead(II) hydroxide

b. Pb2+(aq) + 2I-(aq) ---> PbI2(s)

c. Place two pieces of damp red and blue litmus papers at the mouth of the test tube. If the gas is ammonia, the dampred litmus paper will turn blue. A pungent gas will also be detected.

d. Blue precipitate is formed which dissolves in excess aqueous ammonia to form a dark blue solution.

6. In the experiment shown below, the gas X produced by the action of dilute sulphuric acid on the zinc granules waspassed over two heated metallic oxides. A colourless liquid W was collected and the excess gas X was burnt off at Y.

a. What is gas X? Write the ionic equation for the formation of the gas.b. State what is observed of:i. zinc oxideii. copper(II) oxideWrite equation(s) for any change observed.c. Explain your observation made in bi and bii.d. Give a chemical test to identify liquid W.e. Suggest a suitable drying agent to be placed inside the drying bulb.

f. Why was the excess gas X burnt off at Y?g. What precautions should be taken before the excess gas was lit?

Solution
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a.hydrogen gasZn + 2H+---> Zn2++ H2

bi. Zinc oxide turns from white to yellow.bii. Copper(II) oxide turns from black to pinkCuO + H2---> Cu + H2O

c. Zinc is above hydrogen in the reactivity series so zinc oxide will not be reduced by hydrogen to zinc. Upon beingheated, zinc oxide will change its colour from white to yellow.

Copper is below hydrogen in the reactivity series so copper(II) oxide will be reduced to form pink copper.

d. W is tested with anhydrous cobalt(II) chloride paper. If W is water, the cobalt(II) chloride paper will turn from blue topink.

e. fused calcium chloride

f. Because hydrogen gas is flammable and a mixture of hydrogen and air is very explosive.

g. Ensure that there is no leakage in the apparatus.

Students should be able to explain the importance of metals and their compounds to living systems and the

environment.

Metals in Living SystemsSome metals are an integral part of different structures in living systems. It is well known that calcium is necessary

for the formation of bones, teeth and shells. Ironis present in the haemoglobinof red blood, and magnesium ispresent in chlorophyll, the compound found in the green parts of plants. Potassium is an essential constituent of

protoplasm. Generally the metals function by being part of the mineral structure of the plants and animals. Thus

metals in minerals are important to general body function. They are particularly important for regulation of someactivities like fluid balance, acid-base balance and some forms of metabolism.

Harmful Metals

Some metals may also be very harmful to living systems. Lead is used in certain paints, soldering, car batteriesand in ceramics. Leaded gasolene is used in some internal combustion engines because lead prevents 'knocking'.

Careless disposal of industrial wastes into rivers and streams has caused death to fish, possibly other forms of

aquatic life and human beings. Disposal of household refuse containing fluorescent tubes and electrical switcheswhich contain mercury needs to be done carefully. If the fluorescent tubes are left around and get broken and the

switches get burnt, a lot of mercury vapour will get into the atmosphere and eventually settle on the ground

causing possible contamination of soil and ground water.

Another point to note is that levels which may be good or at least harmless may, when increased, become harmful.

Evidence of this is the effect that excess sodium ions has on blood pressure.

Radioactive IsotopesWhen the number of neutrons in the nucleus is much greater than the number of protons, the nucleus is unstable.

An isotope with an unstable nucleus is called a radio-isotope. Spontaneous changes occur in the nuclei of radio-isotopes. For example, in carbon-14, a neutron in a nucleus changes to give a proton and an electron. The atom

produced by the nuclear change is no longer an atom of carbon but one of nitrogen, since it now has seven protons.

Nuclear changes produce different elements. The electrons given off from the nucleus during nuclear reactions

are called beta particles. The rate at which spontaneous nuclear changes occur is fixed for any one radio-isotopebut varies greatly for different radio-isotopes.

Nuclear changes also produce emissions of alpha particles, gamma rays and x-rays. Radiation can damage thecells of organisms. Some metals have radioactive isotopes. Medical diagnosis and treatment make use of radiation.

Undue exposure to radiation should be avoided.Nuclear EnergyNuclear reactions produce much more energy than combustion of fossil fuels. The atom bomb demonstrated the

vast differences in energy between nuclear change and chemical change (involving the extra-nuclear electrons).

Nuclear reactions can be so controlled that the energy released is used to generate electricity.

The nuclear fuel, usually a radio-isotope of a heavy metal, is made to split into atoms with about half the relative

mass. The energy released can convert water to steam which is used to drive generators to produce electricity.

There is always the p6ssibility of leaks of radioactive material. There are also serious problems associated with

disposal of radioactive wastes. As spontaneous decay of some radio-isotopes is slow, they can be harmful for a

very long time. Stable elements produced from radioactive decay may also be harmful. Some related accidentshave already occurred.

EVALUATION


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In addition to the project mentioned under 'Teaching Strategies', direct questions could be:

1. Name four metals essential to the health of plants and/or animals.2. Name two metals which are harmful to living things.

3. It has been found that corrosion of a metal which involves a chemical reaction is speeded up if the metal

is in contact with a less reactive metal and an electrolyte. The less reactive metal does not react under

these circumstances.(a) Explain what you would expect to be the long-term consequences of using uncoated iron nails to

secure roofing material made of (i) sheets of iron coated with zinc and (ii) sheets of iron coated

with tin, which is less reactive than iron.(b) Which of these two types of sheets would you select for your roof given that tin is cheaper and

more attractive than zinc and that the sheets will eventually become scratched? Justify your choice.

(c) Explain briefly why anodizing can be used to protect aluminium but not iron articles fromcorrosion.

Students should be able to discuss the harmful effects of non-metal and metal compounds to living systems

and the environment.

Sulphur dioxide is a soluble, acidic gas. Sulphur is present in association with coal and hydrocarbons and sulphur

dioxide is released from refineries and is also released when fuels burn. Sulphur dioxide is also produced duringthe extraction of metals from sulphide ores. Sulphur dioxide can cause a disease of the lungs. Sulphur dioxide

also poisons plants in two main ways. It may enter the stomata of leaves and react with the tissue fluids to form

sulphites; with the on-going water loss, the concentration of sulphites could build up, killing the cells. The secondadverse effect on plants occurs when sulphur dioxide dissolves in rain, thereby increasing soil acidity. Aprolonged acid soil situation may lead to plasmolysis and the death of vegetation in an area.

Carbon monoxide is always formed when carbon or carbon containing compounds is oxidized in insufficient

oxygen. This includes oxidation of decaying organic matter and burning of fossil fuels. Carbon monoxide is also

produced during reduction of metallic oxides for the extraction of metals. Carbon monoxide combines much more

readily than oxygen with haemoglobin. When carbon monoxide is inhaled, it prevents the haemoglobin fromtaking oxygen to the body's cells. Continued breathing of air containing carbon monoxide eventually causes death.

Carbon dioxideThe carbon cycle shows a number of ways in which carbon dioxide is produced. The great increase in the use of

fossil fuels and population pressures contribute to the increase in carbon dioxide concentration in the earth's

atmosphere.

When solar radiation reaches the earth's atmosphere about fifty percent of the visible light is reflected into space.The remainder causes warming on earth. Warmed surfaces on earth re-radiate energy as heat. Carbon dioxide,

water vapour and ozone, referred to as "greenhouse gases" acting as a blanket readily absorb some of the radiated

heat energy and in turn warms the earth's atmosphere. This is called the "greenhouse effect". Scientists have founddirect correlation between carbon dioxide concentration and global warming. An increase in carbon dioxide

concentration resulted in an increase in global temperatures. Global warming is a major potential problem that

the world faces.

Hydrogen sulphide has a distinctive smell of rotten eggs(rotten eggs contain hydrogen sulphide). Hydrogen

sulphide is produced during the putrefaction of once-living material which contained sulphur. It is also present in

volcanic gases and some mineral springs. Hydrogen sulphide reacts readily with many cations, forming insolublesulphides Many cations contribute to good health by contributing to enzyme activity, and to the formation of


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bone, pigment and haemoglobin. Although these cations are required in extremely small concentrations, they are

removed by being precipitated as sulphides, and thus are not available for regular function. Careless dumping ofrefuse contributes to atmospheric pollution by hydrogen sulphide.

Aerosols also contribute to atmospheric pollution. The material we emit into the air diffuses, so that we may not

smell it after a while. However, unless it has reacted with some other substance in the air, whatever we emitremains in the atmosphere and could eventually affect us. Some aerosols contain chlorofluorocarbons (CFCs)as propellants. CFCs are low molecular mass hydrocarbons in which some or all of the hydrogen atoms have been

replaced by chlorine and fluorine atoms. CFCs are odourless, non-toxic, stable at high temperatures, inert non-flammable and can be easily liquified. These properties allow CFCs to be very useful. In addition to aerosols,

they are used in foamed plastic products, fast food containers, as coolants in air conditioners and refrigeration. It

is believed that CFCs might be responsible for destroying the protective ozone layer which prevents harmfulultraviolet (UV) radiation from reaching the earth. The carbon-chlorine bond in a CFC molecule is easily broken

by a UV photon producing an active chlorine atom. This atom combines with an ozone molecule forming a

chlorine oxide radical and oxygen. The radical reacts with more ozone producing the chlorine atom again. This

kind of reaction is repeated thousands of times. This type of reaction is thought to account for about eighty percentof the loss of ozone observed.

Phosphates, used in detergents to increase cleaning power and phosphates and nitrates from fertilizers, contribute

to environmental pollution. High concentration of phosphates and nitrates, nutrients for plant growth, causeexcessive growth of aquatic flora. Excessive growth of algae and other aquatic plants and their putrefaction use

up oxygen from the water thus reducing the amount available to aquatic animals resulting in their death.

Lead is a heavy metal poison that is widely encountered. It may occur in food, beverages, paint, water from lead-

sealed pipes, in air polluted by combustion of fuels containing lead compounds. The average person can excrete

about 2 mg of lead per day. Above this quantity accumulation begins in bone and soft tissue. Soluble lead saltsare toxic but metallic lead can be absorbed through the skin.

Children are especially vulnerable to lead toxicity since they retain a larger fraction of absorbed lead than adults.

They do not store lead as quickly as adults, so the lead circulates in the blood stream exerting its toxic effect onvarious organs and systems. Some of the effects of lead in children's blood are as follows:

rising blood pressure lowered intelligence, irreversible mental retardation

reduced haemoglobin formation resulting in anaemia

reduced calcium and vitamin D metabolism

impaired central nervous system function

convulsions, coma may result in death.

The teacher may need to have his/her own scrapbook of newspaper clippings concerning the dangers that resultfrom pollution. For example, in one territory people died when they were sent to clean out a molasses tank, which

had been closed up for a while. In another territory a motorist died in bumper-to-bumper traffic, which stood still

for quite a long period. Elsewhere, desert conditions have been created in the vicinity of some industries. Theseor other situations could be used to introduce the topic.

Students should be able to list uses of the non-metals: carbon, sulphur, phosphorus, chlorine nitrogen,

silicon, and their compounds

An outline of the uses of the non-metals and their compounds is given. Generally, substances as used as theirparticular properties allow.

Table 6: Uses of Non-Metals and their compounds

Non-metal/Compounds Uses

Carbon(Allotropes,carbon dioxide)

Reducing agent in extraction of some metals, lubricant, electrodes, abrasives, solvents,

'rubbers, plastics, fuels, jewellery, decolourizing agent, for gas absorption in masks,cutting glass, drill bits, dry ice as refrigerant, manufacture sodium hydrogen carbonate

and sodium carbonate for water softening, raising agent, carbonated beverages.

Silicon As an alloy with iron for acid-resistant steels, as a semiconductor in transistorized circuits;

as quartz in electronic equipment and lenses; as silicates in jewellery, cement, glass;

silicones in lubricants, water repellants, defoaming agents, electrical insulators, paints,

polishes, finishes; as silicon carbide (in abrasives; as silica gel, a drying agent.

Nitrogen Manufacture of ammonia, nitric acid, fertilizers, dyes, explosives and textiles. As an inert

atmosphere for easily oxidizable substances and electric lamps. As a cooling agent.Phosphorus As phosphor bronze alloys. For fertilizers, detergents and safe matches.

Sulphur Manufacture of sulphuric acid for production of fertilizers, detergents, paints, enamels,

vulcanized rubber, dyes, textiles, pharmaceuticals, anti-fungal ointments and oils.


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Chlorine For manufacture of bleaches, hydrochloric acid, solvents, drug, insecticides, antiseptics,

anaesthetics, refrigerants, plastics, p~ removers and aerosol sprays.


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Chemistry of Gardening

When you have completed this chapter, you will be able to: list the major and trace elements needed by plants

describe the composition of soil

test soil samples describe the importance of nutrients

understand how to improve soil

argue the case for and against herbicides and hydroponics.

Plant nutrients

Plants need the right combination of nutrients to live, grow and reproduce. When plants suffer frommalnutrition, they show symptoms of being unhealthy. Too little or too much of any one nutrient can cause

problems. Plant nutrients fall into 2 categories: macronutrients and micronutrients.

Macronutrientsare those elements that are needed in relatively large amounts. They include nitrogen,

potassium, sulfur, calcium, magnesium and phosphorus.

Micronutrientsare those elements that plants need in small amounts (sometimes trace amounts), like iron,boron, manganese, zinc, copper, chlorine and molybdenum.

Both macro- and micronutrients are naturally obtained by the roots from the soil.All plants require carbon, hydrogen and oxygen for life processes. Carbon dioxide in the air, and water in the

soil supply these elements to plant'.

Elements have special uses in plants, for example:

nitrogen making proteins

sulphur making proteins and oils

calcium cell walls

magnesium making chlorophyll

phosphorus making enzymes, root growth

potassium controlling osmotic pressuretrace elements making enzymes (biological catalysts).

When plants are grown in soil, some of the nutrients are removed from the soil as the plants grow. These must

be replaced or the soil will gradually lose its fertility. The most important elements are nitrogen, phosphorusand potassium (NPK). Organic and inorganic fertilisers can be added to soil to replace missing nutrients.

What is in soil?

Soils vary greatly from place to place, but they have some features in common. All soils contain: rock particles from weathered rocks

organic matter (humus) produced from the remains of plants, animals and soil bacteria.

We can classify soils according to the type and size of particles present.


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Silt is made of fine particles which are smaller than sand particles but larger than those in clay.What are trace elements?

Where do plants obtain most of the elements they need?

Name one element needed to make chlorophyll.

Problems with nutrientsWhen one or more elements are missing from the soil, plant growth suffers.

Nutrient deficient symptom

nitrogen yellowing of leaves

undersized leaves

poor growth rate

phosphorus weakened root system

stunted growthLeaves dull, blue-green or very dark green in colour

low yield of grain or fruit

potassium leaf tips die back

magnesium poor leaf growth and leaves will take on a pale green colour

sulphur yellowing of leaves

MACRONUTRIENTSReplace macronutrients in soils regularly (at least once per growing season) Nutrient Comments Fertilizer Sources

calcium(Ca) Desert soils and water generallyhave plenty of calcium, so

deficiency problems are rare.Excessive calcium can limit the

availability of other nutrients.

Anything with the wordcalcium;

also gypsum.

nitrogen(N) Most plants absorb nitrogen in the form of ammonium or

nitrate.These forms readily dissolve inwater and leach away.

Anything with the words

ammonium,nitrate,orurea.

Also manures.

magnesium

(Mg)

Plants absorb magnesium as anion (charged particle), which

canbe readily leached from soil. Maybe readily leached from

soil ifcalcium is not present.

Anything with the word

magnesium; also Epsomsalts

(magnesium sulfate).

phosphorus

(P)

Plants absorb phosphorus in theform of phosphate. This form

dissolves only slightly in water,but pH strongly affects uptake.


phosphateor bone.Also

greensand.

potassium

(K)

Plants absorb potassium as an ion,which can be readily leached

fromsoil. Desert soils and watergenerally have plenty of

potassium, so deficiency problemsare rare.


potassiumor potash.

sulfur(S) Plants absorb sulfur in the formof sulfate. This readily leaches

from the soil. Sulfur may acidifythe soil (lower the pH).

Anything with the wordsulfate.

Why is humus important?

Humusis the dark, sticky material in soil. It is made up of decaying organic matter and is essential for plantgrowth as it is rich in nutrients.

Benefits of humus: combines with fine clay minerals in the soil. has the ability to absorb water and then release it to plant roots. gives soil an open texture and allows air into it. Mineral salts from soil particles dissolve in the soil

water and become available for use by plants.

If there is li ttle humus, water drains away and the soil is less fert il e.

Improving the soilSome soils are acidic. Few plants grow well in acid soils so gardeners and farmers need to find ways of

reducing the acidity. They can add limeto soil to neutralise the acidity(raise the pH value). Lime is acalcium compound and calcium is needed by plants for good growth.


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The addition of lime

improves soil texture (by causing small clay particles to stick together) releases other nutrients such as potassium.

Problems can arise if a gardener adds both lime andan ammonium fertiliser at the same time. The two

materials react together.

ammonium salt + alkaliammonia gas + waterNH4++ OH-NH3+ H2O

The soil loses nitrogen compounds as ammonia gas escapes into the air.Limecan be added to the soil in threeways:

1. as hydrated lime (calcium hydroxide),2. ground chalk or limestone (calcium carbonate) or3. as burnt lime (calcium oxide).

4. What might be the cause of yellowing leaves on a plant?5. What is humus?

6. How can you neutralise acidic soils?

Herbicides

A chemical which destroys weeds is called a herbicide.Some types of herbicide are selective in their action, for example one type kills broadleaved plants but leaves

grasses and cereal crops untouched. This type is fine for using on lawns and cereal crops, but is no help if the

crop itself is broadleaved, such as soya beans.

Herbicides work in different ways, for example by

disrupting photosynthesisin the plant or by causing the plant to grow rapidlyand then die.

Regular use of herbicidescan lead to long-term problems. Some of them build up in the soil and do not easily biodegrade.

Others may encourage the development of herbicide-resistant weeds. Herbicides can increase crop yields by reducing competition between crop plants and weeds but there is

an environmental cost.

Hydroponics (Growth without soil)

It is possible to grow plants without any soil at all. The trick is to supply the plants with a solution of all thenutrients they require. There are arguments both for and against the use of hydroponics in agriculture.

The advantagesinclude:

Plants can be grown anywhere, even in a desert, provided the facilities are available. There are no problems with weeds. Crops can be grown outside the normal growing season.

alarge crop can be obtained from a small growing area.The disadvantagesare:

The system is very expensive to set up. It requires heavy use of energy resources. It is regarded as 'unnatural' by some consumers. Highly trained staff are required to operate and maintain the system. It is only really suitable for 'luxury' crops rather than basic food crops.

7.What is a herbicide? Give two examples.8. Give one advantage and one disadvantage of using herbicides.9. What is hydroponics and why is this system used for growing certain plants, even though it is expensive?

Biological and chemical control of pestsCrop yields can be severely reduced by attack from various pests. Various methods of pest control are employedto eradicate crop pests.

In biological control, one living organism is used to control the population of another one. For example, ifthere are too many birds eating the seeds on a farm, one solution is to use a trained bird of prey such as a hawk.

The hawk will kill or frighten the other birds away. Some insect populations can be controlled by introducing an

insect predator. For example, ladybirds can be introduced to eat mealybugs. Some plants have the ability to

repel insect pests. If these plants are grown around the crop, less of the crop will be damaged by insects.

In chemical control, various chemicals, called pesticides, are used to destroy insects, rats or other crop pests.

There are problems associated with the use of pesticides. Many pesticides can injure or kill other wildlife. Thepesticides may accumulate in the bodies of animals that feed on the crop and move up the food-chain, causing

damage to other species. Pesticides are often toxic to people, both those who work on farms and those who buy

the food. In time, pest populations develop resistance to the pesticide and it is no longer effective.


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10. What is biological pest control?11. Describe one advantage of using biological pest control over the use of pesticides.

12. How can pesticides harm animals that are not crop pests?

Nitrogen Cycle

How does the Nitrogen Cycle work?

Although nitrogen gas (N2) makes up 78% of the air, plants cannot use it for growth unless it is turned into

nitrate (NO3) in the soil.


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Making nitrogen in the atmosphere into nitrate in the soil is called fixing. Nitrogen fixing occurs naturally in

four ways. Nitrogen fixing also occurs with the Haber process - see below.

1. Nitrogen fixing bacteriain the soil turn nitrogen in the soil into nitrate.

2. Nitrogen fixing bacteria on some plant root nodulesturn nitrogen into nitrate. Plants that have these kinds ofroot nodules include peas, beans and clover.

3. Compounds containing ammonia occur in animal excretion and in dead animals. The ammonia turns intonitrite and there are nitrifying bacteriain the soil that turn nitrite (NO2) into nitrate (NO3). They are called

nitrifying bacteria because they increase the amount of nitrate in the soil.

4. Lightning fixation. Lightning can cause chemical reactions in the atmosphere that make nitrogen react with

oxygen producing nitrous oxides. Nitrous oxides are also made from petrol engine pollution. The nitrous oxidescan react with more oxygen and dissolve in rain water to make dilute nitric acid (HNO3(aq)) (see acid rain).

HNO3contains nitrate and so rain water containing HNO3increases the amount of nitrate in the soil.

Most of the fixing of nitrogen occurs through the natural processes described above. The Haber processformaking fertilisers accounts for about 30% of nitrogen fixing. Nitrate in the soil is essential for plant growth.

5. There are denitrifying bacteriain the soil that turn nitrate (NO3) into nitrogen gas (N2) and the nitrogen gas

goes back into the atmosphere. They are called denitrifying bacteria because they decrease the amount of nitratein the soil. Compare this with the nitrifying bacteria described above at number 3.

The Carbon Cycle

The carbon cycle is simply the process by which Carbon Dioxide is put into and removed from the atmosphere.

This process is very finely balanced to keep the percentage of CO2in the atmosphere at 0.03%. Even so, man

is doing his part to unbalance this cycle by excess burning of fossil fuels and deforestation.

The reactions involved in the carbon cycle are as follows:

1. Combustion- exothermic reactions at a very fast rate

Fuel [Methane] + Oxygen ----> Carbon Dioxide + Water (+ energy)CH4(g) + 2O2(g) ----> CO2(g) + 2H2O (l)

2. Evaporationof Sea Water (containing dissolved CO2). Carbon Dioxide dissolves in sea water.Some of the water evaporates and CO2is released. This is a reversible reaction:

Sea Water + Sunlight Water + Carbon DioxideH2O (l) + CO2(g) H2CO3(aq) [Carbonic Acid]

3. Decayof Organic Matter. Plants and animals decay to leave behind coal, oil, and gas after millionsof years of being squashed by materials above. Some materials just decay and a form of respirationoccurs that releases Carbon Dioxide. CO2is also given off when the raw materials mentionedabove decay.

4. Photosynthesis. Plants make sugar from light using carbon dioxide and water. These constituentsare catalysed by chlorophyll in green leaves to form glucose, which the plant requires. The bi-product is Oxygen.


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5. Respiration is the name of the process which involves animals ingesting food and inhaling air. Theinhaled air dissolves in the blood and is transported around the body. Some of the oxygen is usedwithin cells to oxidise sugars. When this occurs, CO2, H2O, and O2are released.

Glucose + Oxygen ----> Carbon Dioxide + Water + Energy

6. Sedimentation. Under extreme pressure, dead sea creatures decay without air deep underground.Over millions of years, the sea creatures become oil, or natural gas. Also, when shells of sea

creatures etc... build up in layers over millions of years, sedimentary rock can be formed.

Water Cycle

Processes

Many different processes lead to movements and phase changes in water

Precipitation

Condensed water vapor that falls to the Earth's surface. Most precipitation occurs as rain,but also includes

snow, hail,fog drip,andsleet.

Runoff

The variety of ways by which water moves across the land. This includes both surface runoff andchannel

runoff.As it flows, the water may seep into the ground, evaporate into the air, become stored in lakes or

reservoirs, or be extracted for agricultural or other human uses.

Infiltration

The flow of water from the ground surface into the ground. Once infiltrated, the water becomessoil moisture or

groundwater.

Evaporation

The transformation of water from liquid to gas phases as it moves from the ground or bodies of water into the

overlying atmosphere.[6]The source of energy for evaporation is primarilysolar radiation.Evaporation often

implicitly includestranspiration fromplants,though together they are specifically referred to

asevapotranspiration.86% of global evaporation occurs over the ocean.

Condensation

The transformation of water vapor to liquid water droplets in the air, creating clouds and fog.

Transpiration

The release of water vapor from plants and soil into the air. Water vapor is a gas that cannot be seen.

PercolationWater flows horizontally through the soil and rocks under the influence ofgravity.
http://en.wikipedia.org/wiki/Precipitation_(meteorology)http://en.wikipedia.org/wiki/Rainhttp://en.wikipedia.org/wiki/Fog_driphttp://en.wikipedia.org/wiki/Ice_pelletshttp://en.wikipedia.org/wiki/Runoff_(hydrology)http://en.wikipedia.org/wiki/Channel_runoffhttp://en.wikipedia.org/wiki/Channel_runoffhttp://en.wikipedia.org/wiki/Infiltration_(hydrology)http://en.wikipedia.org/wiki/Soil_moisturehttp://en.wikipedia.org/wiki/Evaporationhttp://en.wikipedia.org/wiki/Water_cycle#cite_note-6http://en.wikipedia.org/wiki/Water_cycle#cite_note-6http://en.wikipedia.org/wiki/Water_cycle#cite_note-6http://en.wikipedia.org/wiki/Solar_radiationhttp://en.wikipedia.org/wiki/Transpirationhttp://en.wikipedia.org/wiki/Planthttp://en.wikipedia.org/wiki/Evapotranspirationhttp://en.wikipedia.org/wiki/Condensationhttp://en.wikipedia.org/wiki/Cloudhttp://en.wikipedia.org/wiki/Transpirationhttp://en.wikipedia.org/wiki/Percolationhttp://en.wikipedia.org/wiki/Gravityhttp://en.wikipedia.org/wiki/Gravityhttp://en.wikipedia.org/wiki/Percolationhttp://en.wikipedia.org/wiki/Transpirationhttp://en.wikipedia.org/wiki/Cloudhttp://en.wikipedia.org/wiki/Condensationhttp://en.wikipedia.org/wiki/Evapotranspirationhttp://en.wikipedia.org/wiki/Planthttp://en.wikipedia.org/wiki/Transpirationhttp://en.wikipedia.org/wiki/Solar_radiationhttp://en.wikipedia.org/wiki/Water_cycle#cite_note-6http://en.wikipedia.org/wiki/Evaporationhttp://en.wikipedia.org/wiki/Soil_moisturehttp://en.wikipedia.org/wiki/Infiltration_(hydrology)http://en.wikipedia.org/wiki/Channel_runoffhttp://en.wikipedia.org/wiki/Channel_runoffhttp://en.wikipedia.org/wiki/Runoff_(hydrology)http://en.wikipedia.org/wiki/Ice_pelletshttp://en.wikipedia.org/wiki/Fog_driphttp://en.wikipedia.org/wiki/Rainhttp://en.wikipedia.org/wiki/Precipitation_(meteorology)

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metals reaction

metals andcarbon

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reactions of metals

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