CHEMISTRY SCIENCE UNIT-2 - NIMS Dubai – N.I. · PDF fileCHEMISTRY SCIENCE UNIT-2 Learning outcomes - Core At the end of this unit, students should be able to - Understand that every

Unit - 2

Metals

Chemistry

SCIENCE UNIT-2 CHEMISTRY37

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Syllabus Coverage

Unit 2 - Metals

Chemistry

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General physical and chemical properties of metals .

Reactivity of K, Na, Ca, Mg, Fe, Cu based on their reactions with water or steam, dil HCl

Apparent un-reactivity of Al

Displacement reactions of metals.

Action of heat on the carbonates a n d n i t r i t e s o f m e t a l s -d e c o m p o s i t i o n r e a c t i o n s formation of rust

Extraction of metals from their ores and extraction of iron from haematite ore.

Rusting an oxidation process.

factors which increase the rate of rusting.

Prevention of rust

Metallurgy- steps for extraction on the basis of reactivity of metals.

Reduction of metal oxides with carbon.

Use of metals as alloys

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Reactivity series as the tendency of a metal to form its positive ions

N e w t e c h n i q u e s u s e d f o r extraction of metals.

Composition and uses of different types of iron.

Uses of metals and alloys-Al,Fe,Zn,Cu and their alloys

Composition of alloys.

Diagrammatically represent metals and alloys.

EXTENSION

SCIENCE UNIT-2CHEMISTRY

Learning outcomes - Core

At the end of this unit, students should be able to -

Understand that every element is classified as a metal, nonmetal or semimetal (metalloid) based on its individual properties.

Classify an element as a metal, nonmetal based on experimental observations of physical and chemical properties.

Understand and perform reactions of metals with oxygen, water and acids.

Arrange metals in order of their reactivity by reference to their chemical reactions with water/steam or dilute acid.

Write simple equations to describe oxidation and reduction using both words and symbols.

Identify the steps involved in the extraction of some metals based on their reactivity.

Describe extraction of iron from haemetite ore.

Describe chemical reactions which are not useful-corrosion.

State the conditions necessary for the formation of rust.

Acknowledge rusting as an oxidation process.

Identify the factors which increase the rate of rusting.

Figure out ways to prevent rust.

Understand that an alloy is a metal with some other element mixed in with it

Understand that steel, brass, bronze, solder and amalgam are alloys.

Understand that the properties of an alloy are different from those of the parent metal.

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Learning outcomes - Extension

Cross curricular links

At the end of this unit, students should be able to -

Appreciate that the reactivity of metals is related to the ease with which they form ions .

Understand various techniques used for extraction of metals.

Understand the composition and uses of different types of iron.

Understand that the introduction of 'smart' alloys has increased the number of applications of alloys.

Describe the use of metals- in high strength structures, as alloys, for anti corrosion applications, conductors.

Recognise other examples of corrosion.

English- Discussions in the class on various topics like rusting. History- The sequence of copper age followed by bronze-age and then iron-age can be related to the reactivity of copper and iron and their extraction from respective ores.Architecture- Metals used in the construction.

Economics- Loss of revenue due to rusting.

Physics- Strength of materials, conductivity, resistance, use of metals and their alloys.

Biology- Comparing the concept alloying with hibridisation.

Music- A rap song on properties of metals.

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Teacher's Activity Student's ActivitySteps to be

followed

Teacher may start the unit by

showing the video on metals. The

beginning of the video is to be

shown to build the back ground

knowledge.

Video developed at CIET-NCERT

attached.

Students discuss about what

they know about metals,

they may discuss their

experiences of visit to

m u s e u m s o r f a m i l y

vacations.

Warm up

Students complete the -I

know and want to know

sections of the KWL sheet in

groups of two and then share

in the class.

As they have learnt few

reactions of metals in the unit

'Acids, Bases and Salts' they

should be able to recall

studied properties.

Pre content Teacher may divide the students

into groups of two and give them

the task of completing the K,W

sections of KWL sheet after they

brainstorm what they already

know about metals. They may

include their understandings from

the video shown as warm up. 10 to

20 minutes of the period may be

given to this activity.

Students make a note of the

position of metals in the

periodic table. They come to

the board to locate the metal

on the periodic table.

Content -1

Placement of

metals in the

periodic table

LESSON TEMPLATE : METALS

Students have studied about

periodic table in grade 9 teacher

may either bring the table to the

class or show the metals pps in the

class.

Students part ic ipate in

debates and discussions

initiated by the teacher on the

given topics in worksheet 1.1.

After performing the activity 1

S t u d e n t s c o m p l e t e t h e

worksheets.

Worksheet 1.1 and 1.2

Content-2

PHYSICAL

PROPERTIES OF

METALS AND

NON METALS

The teacher leads the students to

e x p l o r e v a r i o u s p h y s i c a l

p r o p e r t i e s o f m e t a l s a n d

distinguish them from non-metals

on the basis of the properties.

The teacher facilitates students to

u n d e r s t a n d t h a t p h y s i c a l

properties are not enough to

classify elements as metals, non-

metals and metalloids.

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SCIENCE UNIT-2 CHEMISTRY

As the students realize in the

previous content that physical

properties are not enough to

classify the elements as metals

teacher may lead the students to

the concept attainment that

chemical properties of metals are

different and these may be used to

distinguish elements as metals.

The observation expected from the

students is that a combination of

physical and chemical properties

put together is the best suited for

this classification. During the

development of the previous

content students must have

realized that all metals do not react

with air, water and acid with same

strength so we should arrange the

metals in the order of their

reactivity.

The metals may be arranged in the

increasing order of their reactivity

by observing the displacement

reactions between aqueous

solution of a metal salt and another

metal. A more reactive metal

would displace a less reactive one

from its salt solution.

One of the videos attached may be

selected to show how to perform

the activity.

Content -3

CHEMICAL

PROPERTIES OF

METALS AND

REACTIVITY

SERIES OF

METALS

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Students perform activities

for concept attainment under

activity 2 and follow it up

with the worksheets.

SCIENCE UNIT-2CHEMISTRY42

Content-4

Extraction of

Metals

Content-5

Corrosion of iron-

causes and

prevention

Students build links between

the reactivity of metals and

the ease of their extraction.

They should be able to out

line the method used for

extraction of a metal

according to its position in

the reactivity series.

Students make a hypothesis

a b o u t t h e c o n d i t i o n s

necessary for rusting and test

the hypothesis through

experimentation.

Students a lso make a

hypothesis about effect of

presence of salts in water on

rate of rusting and test the

h y p o t h e s i s t h r o u g h

experimentation.

Extraction of metals must be

linked to the reactivity series of

metals. Bringing to the notice of

the students that during the

development of the civilizations

bronze age was followed by the

iron age. The reason is bronze is

an alloy of copper and copper

being less reactive was easy to

extract from its ores as compared

to iron. Iron forms stable oxides as

iron is a reactive metal so it was

difficult for the early man to

recover iron from its ore.

ppt on extraction of iron from its

ore haemetite must be shown

along the video.

Iron age saw the development of

agricultural practices due to

strength of the metal iron. But the

major problem with iron is that

due to its high reactivity it reacts

with oxygen of the air to form a

powdery oxide-rust.

Students should be led to make a

hypothesis about the conditions

necessary for rusting.

Students should be led to make a

hypothesis about effect of

presence of salts in water on rate

of rusting. Rusting of ships in sea

water should be brought to the

notice of the students.


Students depict the structure

of alloys using the plasticine

modellling.

Students may complete the

group projects assigned by

the teacher.

Content-6

Alloys

Post Content

The loss of revenue due to rusting

of iron should be brought to the

notice of the students. A small

discuss ion to bui ld cross

c u r r i c u l a r l i n k a g e w i t h

Economics and English subject

should be initiated in the class.

Methods of prevention of rust to

be discussed in the class with

involvement of students, taking

on from the previous knowledge

of the students.

The concept of alloying should be

brought as the method of

enhancing the properties of

metals.

Revision worksheets 1and 2


ACTIVITY : 1

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Learning objectives:

CLASSIFICATION OF ELEMENTS

The students will be able to-

Understand that every element is classified as a metal, nonmetal or semimetal (metalloid) based on its individual properties.

Classify an element as a metal, nonmetal or semimetal based on experimental observations of physical and chemical properties.

There are 118 chemical elements known at present. On the basis of their properties, all these elements can be broadly divided into two main groups : Metals and Non-Metals. Metals are the elements (except hydrogen) which form positive ions by losing electrons whereas non-metals are the elements which form negative ions by gaining electrons. The elements which show the properties of both metals and non metals are called metalloids.

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A metal is a lattice of positive metal 'ions' in a 'sea' of delocalised electrons.

Metallic bonding refers to the interaction between the delocalised electrons and the metal nuclei.

The physical properties of metals are the result of the delocalisation of the electrons involved in metallic bonding.

Conductivity

Solid and liquid metals conduct heat and electricity.

The delocalised electrons are free to move in the solid lattice. These mobile electrons can act as charge carriers in the conduction of electricity or as energy conductors in the conduction of heat.

PHYSICAL PROPERTIES OF METALS AND NON METALS

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Melting Points

Malleable and Ductile

Optical Properties

In general, metals have high melting and boiling points because of the strength of the metallic bond.

The strength of the metallic bond depends on the

number of electrons in the delocalised 'sea' of electrons.

(More delocalised electrons results in a stronger bond and a higher melting point.)

packing arrangement of the metal atoms.

(The more closely packed the atoms are the stronger the bond is and the higher the melting point.)

Group I metals have relatively low melting points compared to other metals because they: only have 1 electron to contribute to the delocalised 'sea' of electrons are not forming as many metallic bonds as other metals because Group I atoms are inefficiently packed have large atomic radii so the delocalised electrons are further away from the nucleus resulting in a weaker metallic bond

Metals are malleable and ductile.

The delocalised electrons in the 'sea' of electrons in the metallic bond, enable the metal atoms to roll over each other when a stress is applied.

Metals typically have a shiny, metallic lustre.

Photons of light do not penetrate very far into the surface of a metal and are typically reflected, or bounced off, the metallic surface.

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Properties Metals Non-Metals

Hardness 1. Metals are usually very hard except alkali metals like lithium, sodium and potassium which are so soft that they can be cut by a knife.

Non-Metals are normally soft except diamond and boron

Lustre Most of the metals have metallic luster (shine) and they can be polished.

Non metals are non lustrous except iodine crystals

Malleability Metals are malleable i.e. can be beaten into very thin sheets without breaking. Gold and silver are most malleable metals.

Non-Metals are not malleable. They are brittle.

Ductility Metals are ductile i.e. can be drawn into thin wires. Gold and silver are the most ductile metals.

Non-Metals are not ductile. They are brittle.

Thermal Conductivity

Metals are good conductors of heat and have high melting points (Gallium and caesium have very low melting points). The best conductor of heat is silver followed by copper. Lead is poorest conductor of heat..

Non-Metals are poor conductors of heat and usually have low m e l t i n g a n d b o i l i n g points.(Boron, silicon and carbon have high melting points.

Electrical Conductivity

Metals are good conductors of electricity. Silver is the best conductor of electricity. Copper and aluminium are the next best conductors of electricity.

Non-Metals are poor conductors of electricity except graphite

State of existence

Metals are sol id at room temperature except mercury and gallium which is liquid at room temperature.

Non metals are mostly solids or gases except bromine which is liquid at room temperature.

Many non-metals exist in allotropic forms.

Metals normally do not exist in allotropic forms.

Allotropy


Note : When a metal is heated, its atom gain energy and start vibrating vigorously. This energy is transferred to the electrons, which can move through the metal. When the energetic electrons move through the metal, they transfer energy to the other electrons and atoms of the metal. In this way heat is conducted from one end of the metal to its other end.

But these properties are not sufficient for classifying matter as metals or non metals as there are many exceptions to these properties. Let us see what the exceptions to these properties are:

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Properties Exceptions in metals Non-Metals

Sonorous Sodium and Potassium are very soft, Mercury exists in liquid state and hence is not sonorous.

Exceptions in non metals Iodine is a non metal with metallic lusture. Graphite also has some lusture.

S. No.

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Lusturous

Sodium and Potassium are very soft, Mercury exists in liquid state and hence is not malleable

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H a r d b u t not brittle

Sodium and Potassium are very soft; these can be even cut with a knife. Mercury exists in liquid state and hence is not hard. Gallium and Caesium have very low melting point and melt above

030 C. Zinc is brittle

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Malleable3

Ductile and Sodium and Potassium are very soft, Mercury exists in liquid state and hence is not ductile

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State Mercury exists in liquid state. Gallium and Caesium have very low melting point and

0melt above 30 C. Rest all metals exist in solid state.

Hydrogen, Oxygen, nitrogen, H e l i u m , N e o n , A r g o n , Fluorine and Chlorine exist as gases. Bromine exists in liquid s t a t e . C a r b o n , s u l p u r , phosphorous exist in a solid state.

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Thermal Conductivity

Good conductors Diamond is a good conductor of heat.

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Electrical Conductivity

graphite is an excellent conductor of electricity

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Thus we see there are exceptions in physical properties of metals and non metals. We need to study some more properties to classify substances as metals and non metals. As physical properties are not sufficient for this sorting let us study the chemical properties of metals and non metals.

METALS AND NON-METALS DIFFER IN PHYSICAL PROPERTIES.

Elements Symbol Type of Hardness Malleability Ductility Heat Electricity Sonority

Surface conduction conduction

Iron Fe Lustrous Hard High High Good Good Good

Copper Cu Lustrous Hard High High Good Good Good

Aluminium AI Lustrous Hard High High Good Good Good

Magnesium Mg Lustrous Hard High High Good Good Good

Sodium Na Lustrous Soft Non-melleable Non-ductile Good Good Poor

Pottasium K Lustrous Soft Non-melleable Non-ductile Good Good Poor

Carbon C Lustrous Hard Non-melleable Non-ductile Poor Good Poor

(graphic) (diamond (graphite)

is the

hardest

natural

substance)

Sulpher S Non-Lustrous Soft Non-melleable Non-ductile Poor Poor Poor

Iodine I Lustrous Soft Non-melleable Non-ductile Poor Poor Poor


Comments :

Aim :

Materials required:

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Procedure:

We cannot group elements as metals or non metals according to their

physical properties , as there are many exceptions and there is no set pattern of physical properties. Carbon, sulphur and iodine are non metals and rest of the elements in the table above are metals.

Type of activity: Content based

To classify an element as a metal, nonmetal or semimetal based on experimental

observations of physical properties.

Eight labels, One dropper bottle, One hammer, Eight pieces of paper each measuring

approximately 3.5 to 5 inches, safety goggles for eye protection.

Seven vials with caps, filled with the following:

Iron filings

Sulphur rolls

Zinc Granules

Graphite (replacement leads for mechanical pencils work well)

Silicon

Tin Granules

Carbon

(If the above materials are not available, some substitutes are: paper clips, nails, fishing weights, charcoal)

A conductivity apparatus such as 9-volt battery, a small appliance light bulb, and three pieces of insulated copper wire to make an open circuit (the circuit will be closed with each of the seven samples)

1. Before beginning this activity, students should already have a basic understanding of the differences in the physical properties of luster, malleability, and conductivity between metals and nonmetals. Students should understand


that metalloids have characteristics of both metals and nonmetals and that acids will react with most metals and not nonmetals. Students should always be reminded to follow proper safety procedures and guidelines when working with any chemicals.

2. Break the class into groups of two or three, depending on class size. Use labels to number a set of vials from 1 through 7. Each group should receive a set of seven vials. Fill and label the complete set of vials as follows:

Vial 1-iron filings

Vial 2-sulphur

Vial 3- zinc granules

Vial 4-graphite

Vial 5-silicon

Vial 6-tin granules

Vial 7-carbon

Teacher's Notes

3. At each group's lab station place a hammer, eight pieces of white paper, a set of vials containing samples 1-7, a conductivity apparatus, a test tube holder, and seven test tubes.

4 Before students begin their experiment, make sure they are all wearing their safety goggles. Instruct a student in each group to take a piece of white paper, fold it in half, open it, and place it on the lab top. Another member of the group should then open vial 1 and shake about a pea-sized portion of the sample onto the white paper. Each student should observe the appearance of the sample and record his or her observations in the "colour" and "lustour" columns of the observation table.

5. Next, have one student in each group place a second piece of paper over the top of the sample and crush the sample with the hammer. The student should then remove the top piece of paper and each student in the group should observe the sample and record his or her observations in the "malleability" column of the observation table.


6. Have one student within each group test the conductivity of the sample with the conductivity apparatus by placing the ends of the wires not attached to the power source or light bulb into the sample vial. DO NOT LET THE WIRES TOUCH EACH OTHER. Each student should observe the light bulb and record his or her observations in the "conductivity" column of the observation table.

7. Each group should then repeat steps 3 through 6 above for the remaining six vial samples.

8. Each student, based on his or her group's experimental observations and the knowledge he or she has about the properties of metals and nonmetals, should then classify each of the samples as a metal, nonmetal, or semimetal. He or she should record the answer in the "classification" column of the data observation table.

9. After submission of the data sheets, discuss with the class their conclusions for each of the seven vials. What conclusions did they draw and why?

Observations:

Sample Colour Lustre Malleability Conductivity Classificationnumber

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2.

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4.

5.

6.

7.

Video on Physical behaviour of metals

http://www.youtube.com/watch?v=5BfqXXUmZdg&feature=BF&playnext=1&list=QL&index=1


WORKSHEET-1.1

1. What accounts for the observed differences between samples? Describe how these samples might be arranged with respect to other elements in the periodic table. What other tests could be performed to help identify materials as metals, nonmetals, and semimetals? What are some useful applications of such tests?

2. Our society is rapidly consuming raw materials that are nonrenewable. Discuss whether we should rely on scientists to develop new materials to replace older ones that are no longer available or consume and discard currently available materials more wisely. Is every material replaceable?

3. Discuss what types of materials are most important to recycle. Debate whether recycling should be made the law nationwide. Is the expense of enforcing such a law justified or is the money better spent on other social or scientific programs?

4. Given the high quality and ready availability of substitutes, debate whether rare materials such as gold or diamonds should be used for jewellery and art or saved solely for technology and industry.

5. Discuss how our global society would change if the technology to change one element into a different element were developed. Is this a technology that, if developed, should be made public?

6. One area of great advancement in materials science is the area of medicine. Artificial joints are available, and scientists are currently developing artificial skin and blood. Discuss whether there should be limits placed on the development of physiological substitutes. Can you think of any reasons why advances in this area should be controlled?

7. Many of the greatest advances in materials science have been made accidentally. In light of this, discuss the value of following the scientific method versus chance. Is there ever room for chance when strictly adhering to the scientific method?

Resource: http://www.discoveryeducation.com/teachers/free-lesson-plans/pursuit-of-the-properties-of-metals-and-nonmetals.cfm


WORKSHEET-1.2

1. Iron is used to make cooking pots.

(a) Give at least three reasons why it is such a good structural material for cooking pots.

2. Copper and silver can both be used in electrical wiring.

(a) Which properties make them good metals to use for electrical wiring?

(b) Why is copper more likely to be used than silver?

3. Give an example of a metal which

(I) is a liquid at room temperature.

(ii) can be easily cut with a knife.

(iii) is the best conductor of heat.

(iv) is a poor conductor of heat.

4. Explain the meanings of malleable and ductile.


ACTIVITY : 2

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CHEMICAL PROPERTIES OF METALS AND REACTIVITY SERIES OF METALS


The students will be able to-

Discuss the reactivity series of metals-place metals (K,Na,Ca,Mg,Fe,Cu) in an order of decreasing reactivity .

Understand the reactions of metals with water or steam,

Understand the reactions of metals with dil HCl,

Understand the reduction of metals oxides with carbon.

Study some typical displacement (redox) reactions

Identify oxidizing and reducing agents and write half reactions.

Develop an activity series for a limited number of elements.


The arrangement of metals in order of decreasing reactivities is called reactivity series or activity series of metals.

1. All metals which are placed above hydrogen in the activity series can lose electrons more readily than hydrogen whereas the metals placed below hydrogen in the activity series lose electron less readily than hydrogen.

2. If a metal loses electron easily to form positive ions, it will react readily with other substances i.e. it will be a reactive metal. On the other hand, if a metal loses electron less rapidly to form positive ion, it will react slowly with the other substances i.e. metal will be less reactive.

The chemical reactivity of elements varies over a wide range. Some elements,like sodium and fluorine, are so reactive that they are never found in the free,uncombined state in nature. Other elements, like xenon and platinum, are nearly inert and can be made to react with other elements only under special conditions.

The reactivity of an element is related to its tendency to lose or gain electrons, which is to be oxidized or reduced.

NOTE: -

Principles:


In principle, it is possible to arrange nearly all the elements into a single series in order of their reactivities. A series of this kind indicates which free elements are capable of replacing other elements from their compounds. Such a list is known as an activity or electromotive series.

To illustrate the preparation of an activity series we will experiment with a small group of selected elements and their compounds. A generalized single replacementreaction is represented in the form:

A + BC → B + AC

Element A is the more active element and replaces element B from the compound BC. But if element B is more active than element A, no reaction will occur.

Let us consider two specific examples, using copper and mercury:

A few drops of mercury metal are added to a solution of copper (II) chloride. In this example no change is observed even after the solution has been standing for a prolonged time, hence, we conclude that there is no reaction.

Hg( l ) + CuCl (aq) → No reaction 2

From this evidence we conclude that mercury will not replace copper in copper compounds, and therefore mercury is a less active metal than copper.

A strip of metallic copper is immersed in a solution of mercury (II) chloride. In this example the copper strip is soon coated with metallic mercury and the solution

2becomes pale bluegreen (the color of Cu + ions in aqueous solution).

Cu (s) + HgCl (aq) → Hg (l) + CuCl (aq)2 2

From this evidence we conclude that copper will replace mercury in mercury compounds. Therefore copper is more reactive than mercury and is before mercury in the activity series.

The single replacement equation given above shows that in terms of oxidation numbers, the chloride ion remains unchanged, but the oxidation number of mercury has changed from 2+ to 0, and the oxidation number of copper has changed from 0 to 2+.

Expressed another way, the actual reaction that occurred was the replacement of a 2 0mercury ion (Hg + ) by a copper atom (Cu ).

Example 1:

Example 2:


This can be expresses more simply in the form of a total ionic equation:

0 2+ - 0 2+ -Cu (s) + Hg (aq) + 2cl (aq) → Hg (l) + Cu (aq) + 2Cl (aq)

Since the chloride ions do not experience chemical change, they are called spectator ions and can be removed from the total ionic equation. This results in a net ionic equation:

0 2+ 0 2+Cu (s) + Hg (aq) → Hg (l) + Cu (aq)

The net ionic equation shows only those reactants that are actually consumed and those products that are actually formed in the reaction. Total ionic equations are an intermediate step in writing net ionic equations, and for this reason they are often omitted. Since single replacement reactions Involve changes in oxidation number, they are also classified as oxidation reduction reactions.

1. Reaction with oxygen- Almost all the metals react with oxygen (or air ) to form metal oxides.

Metal + Oxygen → Metal oxide

These oxides are basic in nature. When these oxides are dissolved in water they give alkaline solutions. For example sodium metal reacts with oxygen of the air and form sodium oxide which reacts with water to form an alkali called sodium hydroxide.

4Na + O → 2Na O2 2

Sodium oxide

Na O + H O → 2NaOH2 2

Sodium hydroxide

But some metal oxides such as aluminium oxide, zinc oxide etc. shows both acidic as well as basic behaviour. Such metal oxide which reacts with both acids as well as bases to produce salts and water are known as amphoteric oxides.Aluminium oxide reacts in the following manner with acids and bases-

Al O + 6HCl → 2AlCl + 3H O2 3 3 2

Aluminium Oxide Aluminium Chloride

Al O + 2NaOH → 2NaAlO + H O2 3 2 2

Aluminium Oxide Sodium aluminate


Different metals combine with oxygen with different speed-

(i) Metals like sodium and potassium (most reactive metals) combine with oxygen of air even at room temperature to form their oxides.

4Na + O → 2Na O2 2

Sodium Oxide

In fact the reaction is so vigorous that these metals catch fire if kept in the open air. That is why they are always kept under kerosene oil.

(ii) Magnesium. reacts with oxygen upon heating and burns brightly to form magnesium oxide which is a white powder.

Heat

2Mg + O → 2MgO2

Magnesium (From air) Magnesium Oxide

MgO dissolves in water to form magnesium hydroxide.

2Mg(s) + O (g) 2MgO(s) Mg(OH) (aq)2 2

(Magnesium) (Magnesium oxide) (Magnesium Hydroxide)

The resultant solution turns the colour of the red litmus to blue indicating that it is basic in nature.

(iii) Iron reacts with oxygen but does not burn to form Fe O .3 4

3Fe + O Fe O2 3 4

Iron Oxygen Iron Oxide

(iv) Copper metal is even less reactive than Iron. It reacts with oxygen on prolonged heating to form a black mass of copper (II) oxide.

2Cu + O 2CuO2

Copper Copper(II) oxide

(v) Silver and gold do not react with oxygen even at very high temperature.

(vi) When sulphur powder is burnt in air, it oxidizes to form sulphur dioxide which dissolves in water to form sulphurous acid.

S(s) + O (g) SO (g) H SO (aq)2 2 2 3

(Sulphur) (Sulphur dioxide) (Sulphurous acid)

H O2

Heat

Heat

H O2


The resultant solution turns the colour of the blue litmus to red indicating that it is acidic in nature.

Most metals give to basic oxide and most non-metals give acidic oxides.

When metal sample are brought over a flame to them, they show different reactivities. Some important observations are -

Table: 3.1

Metal Whether Ease of Flame colour Appearance Reaction Solubility Burning burning while of metal involved of productsobserved burning surface formed

after burning

For silver and gold, burning is not observed even at relatively higher temperatures.

2. Metals reacts with water to form metal oxides or metal hydroxide along with the evolution of hydrogen gas.

(i) Metals like sodium and potassium placed higher in the reactivity series react with water violently even in the cold and immediately catch fire.

2K(s) + 2H O (l) → 2KOH (aq) + H (g) + heat energy2 2

Cold

(ii) Calcium reacts with water less violently while magnesium placed below calcium reacts with water only upon heating.

2Ca(s) + 2H O (l) → 2Ca (OH) (aq) + H (g) + heat energy2 2 2

Calcium Calcium Hydroxide

Comments:

Reaction with water-

Sodium

Potassium

Magnesium

Aluminium

Zinc

Copper

Yes

Yes

Yes

Yes

Yes

No

Very easy

Very easy

Very easy

Difficult

Difficult

Golden

Yellow

Violet

Brilliant

White

Brilliant

White

Bluish white

Turns into ashes

Turns into ashes

Turns into ashes

Dull

Dull

Black

4Na+O → 2Na O2 2

4K+O → 2K O2 2

2Mg+O → 2MgO2

4Al+ O→ 2Al O3 2 2 3

2Zn+O→2ZnO2

2Cu+O → 2CuO2

Soluble

Soluble

Soluble

Insoluble

Insoluble

Insoluble


(iii) Metals like aluminium, zinc and iron react with water only when steam is passed over heated metals. For example,

2Al(s) + 3H O (l) → Al O (s) + 3H (g)2 2 3 2

Aluminium Steam Aluminium Oxide

(iv) Metals like copper, silver, gold , platinum do not react with water under any conditions.

3. Reaction with acids-Metals react with acids to give a salt and hydrogen gas.

Metal + Dilute acid → Salt + Hydrogen

The reactivity of metals with acids is also influenced by their position in the activity series.

Metals placed above hydrogen in the reactivity series react with dilute acids like hydrochloric acid and sulphuric acid to evolve hydrogen gas and forming metal salts.

2Na + 2HCl → 2NaCl+H (g)2

SodiumHydrochloric acidSodium Chloride

Metals which are placed below hydrogen in the reactivity series can not lose electrons to H+ ions of the acid i.e. Hydrogen gas is not liberated.

Hydrogen gas is not evolved when a metal reacts with nitric acid (HNO ). The 3

acid is a strong oxidizing agent and oxidizes hydrogen evolved in the reaction to H O and itself gets reduced to an oxide of nitrogen. For example,2

Zn + 4HNO (conc) → Zn(NO ) (aq) + 2H O + 2NO (g)3 3 2 2 2

Zinc Nitric acid Zinc Nitrate

4.

Reactive metals can displace less reactive metals from their compounds in solution or molten form. For example zinc when reacts with copper sulphate solution displaces copper from the solution as it is more reactive than copper.

Zn(s) + CuSO (aq) → ZnSO (aq) + Cu(s)4 4

Zinc Copper sulphate Zinc Sulphate Copper

(Blue colour) (colourless)

Potassium, sodium, lithium and calcium all react violently with dilute sulfuric acid and dilute hydrochloric acid. It is dangerous to put these metals into an acid.

NOTE :

Reaction with solutions of other metal salts-


For example

sodium + hydrochloric acid → sodium chloride + hydrogen.

2Na(s) + 2HCl(aq) → 2NaCl(aq)+H (g) 2

Magnesium, aluminium, zinc, iron, tin and lead react safely with dilute acid. Magnesium is the fastest and lead is the slowest of the six.

magnesium + sulfuric acid magnesium sulfate + hydrogen.

Mg(s)+H2SO4(aq) → MgSO4(aq)+ H2(g)

magnesium + hydrochloric acid → magnesium chloride + hydrogen.

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

aluminium + hydrochloric acid → aluminium chloride + hydrogen.

2Al(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2(g)

zinc + sulphuric acid → zinc sulphate + hydrogen.

Zn(s) + H SO4(aq) → ZnSO4(aq) +H2(g)2

iron + hydrochloric acid iron(II) chloride + hydrogen.

Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)

tin + hydrochloric acid tin(II) chloride + hydrogen.

Sn(s) + 2HCl(aq) → SnCl (aq)+ H (g)2 2

lead + sulphuric acid → lead sulphate + hydrogen.

Pb(s)+ H SO (aq) → PbSO (s) H (g) 2 4 4 + 2

The reaction of zinc with sulfuric acid is often used to make a small amount of hydrogen in the laboratory.

Metals below hydrogen in the reactivity series (copper, silver, gold and platinum) will not react with dilute acid. They cannot displace hydrogen from the non-metal anion.

12. If an iron wire or iron filings are put in copper sulphate solution taking in a beaker, the blue colour of copper sulphate changes into green and reddish brown particles of copper settle at the bottom of the beaker or are deposited on the iron wire or iron filings. This is due to the following reaction taking place:


Fe (s) + CuSO4 → FeSO + Cu(s) 4

(Iron) (Copper(II)Sulphate) (Iron(II) Sulphate) (Copper)

Blue Green Reddish brown

This shows that iron is more reactive than copper. No reaction takes place if copper wire is dipped in iron (II) sulphate solution. This again proves that copper is less reactive than iron.

The displacement reactions can be written as ionic equations. In the example using iron and copper(II) sulfate

iron + copper(II) sulfate → iron sulfate + copper.

Fe(s) + CuSO4(aq) → FeSO (aq) + Cu(s) 4

Copper(II) sulfate and iron sulfate are ionic compounds.

When they are dissolved in water the ions become separated by the water molecules

If we write the equation showing the ions we have

2+ 2 2+ 2Fe(s) + Cu (aq)+ SO -(aq) → Fe (aq) + SO -(aq) + Cu(s) 4 4

2+In going from reactants to products iron metal - Fe(s) has become iron ions - Fe (aq)

2+copper ions - Cu (aq) have become copper metal - Cu(s)

2-sulfate ions - SO (aq) are not changed during the reaction.4

Sulfate ions are the same on the left and the right side of the arrow.

Ions which do not change during the reaction are called spectator ions.

Spectator ions can be left out of the equation, giving

2+ 2+Fe(s) + Cu (aq) → Fe (aq) + Cu(s)

This is the ionic equation

for the reaction between iron and copper(II) sulphate.

Iron is oxidised and copper is reduced.

Similarly, the reaction between tin and lead chloride may be written as

2+ 2+Sn(s) + Pb (aq) → Sn (aq) + Pb(s)

Tin is oxidised and lead is reduced.

Ionic Equations.


Oxidation means losing electrons and reduction means gaining electrons.

In the reaction between iron and copper(II) sulfate

2+ 2+Fe(s) + Cu (aq) → Fe (aq) + Cu(s)

Iron metal (Fe(s)) loses 2 electrons to form iron ions (Fe2+(aq))

- 2+Fe(s) - 2e (aq) → Fe (aq)

This is called oxidation. Iron is said to be oxidised.

Copper ions (Cu2+(aq)) gain 2 electrons to form copper metal (Cu(s))

2+Cu (aq) + 2e(aq) → Cu(s)

This is called reduction. Copper is said to be reduced.

To help you to remember that oxidation means losing electrons and reduction means gaining electrons, remember O I L R I G

(Oxidation Is Loss, Reduction Is Gain - of electrons).

Oxidation and reduction always occur together.

They are called redox reactions (pronounced "reed - ox" for reduction / oxidation).

Oxidation may also be defined as gaining oxygen and reduction defined as losing oxygen.

For example, when copper(II) oxide is heated with carbon

copper(II) oxide + carbon → copper + carbon dioxide.

2CuO(s) + C(s) 2Cu(s) + CO 2(g)

Carbon is oxidised by gaining oxygen to form carbon dioxide.

Copper (in CuO) is reduced by losing oxygen to form copper metal.

2+Note that copper has changed from Cu (in CuO) to Cu(s) and so has gained electrons.

2+Cu (aq) + 2e- (aq) → Cu(s)


ACTIVITY : 2(A)

Type of activity:

Aim:

Procedure:

Content based

To compare the activity of certain given metals.

Materials required:

2 cm Strip of metals like Pb, Cu, Zn, Fe, Al and Mg

Steel wool

24-well plate

Phenolphthalein

Conc HCl

3M HNO3

Beral pipet

0.1 M Pb(N0 ) , 0.1 M Cu(N0 ) , 0.1 M Zn(N0 ) , 0.1 M Pb(N0 ) , Fe(NH ) (S0 ) .3 2 3 2 3 2 3 2 4 2 4 2

(a) Metal and Hot Water - Rate of Disappearance of Metal

1. Pb, Cu, Zn, Fe, Al, and Mg- Thoroughly clean a 2 cm strip of each metal with steel wool to remove any oxide coating. This is especially critical for Al and Mg since they form tough, protective oxide coatings (the dull finish on a metal caused by its reaction with the oxygen in air.) Obtain a 24-well plate, rinse six microcells with boiling, distilled (or deionized) water, and then half-fill with boiling water; add 1 drop of phenolphthalein to each cell.


2. Quickly place each metal into the microcells, Al and Mg first (see Figure below). Look for H (g) evolution, discoloration of the metal surface, a color change of the 2

solution (due to the phenolphthalein), or a disappearance of the metal. Some reactions may not be immediately apparent. Record your observations.

(b) Metal and Nonoxidizing Acid - Rate of Hydrogen Gas Evolution

1. Add 5 drops of conc HCl (Caution: avoid skin contact) to the metals from Part B in which no reaction was observed. Swirl the solutions; allow 10- 15 minutes for evidence of a reaction. Record your observations

(c) Metal and Oxidizing Acid - Rate of Gas Evolution

1. If a metal shows no reaction in Parts A or B, draw off the water/acid solution with a Beral pipet and discard it. Add I mL of 3 M HNO3 (Caution: avoid skin contact). If a reaction is now observed, record your observation.

(d) Metal and More Reactive Metal Cation - A Displacement Reaction

1. Pb, Cu, Zn, and Fe-Place a small amount of a freshly cleaned metal in four consecutive microcells of the 24-well plate-one metal in each microcell. Half-fill each microcell with 0.1 M Pb(N03) . Any tarnishing or dulling of the metal or 2

changing of the color of the solution indicates a reaction. Allow 5-10 minutes for a reaction to be observed.

2. Repeat the procedure, using the following 0.1 M test solutions on each metal: Cu(N03) , Zn(N03) , and Fe(NH4) (S04) . In each case the same metal strip may be 2 2 2 2

reused if it remains unreacted from the previous test, rinsed with distilled water, and cleaned with steel wool. Record.

(e) Establishment of Activity Series

Using the observations from Parts A through E, list the metals in order of decreasing activity.

Dispose of the metals in the "Waste Solids" container and the solutions in the "Waste Salt Solution" container.

www.harpercollege.edu/tm-ps/chm/.../series/3perform.htm

Disposal Information:

Resource:

Observations:


___________________________________________________________________________

___________________________________________________________________________

___________________________________________________________________________

___________________________________________________________________________

___________________________________________________________________________

___________________________________________________________________________

To study the reactions of magnesium, zinc, iron and copper in solutions of their salts, and to work out a reactivity series for them. The most reactive metal will be coated by all the others, and the least reactive metal will not be coated by any of the others.

Goggles

Bench mat

Spotting tile

Teat pipette

Strips of magnesium, zinc, iron, copper

magnesium sulphate

zinc sulphate

iron sulphate

copper(II) sulphate

1. Copy the table shown below into your book. Allow a minimum of three lines for each row.

Conclusion:

Aim:

Materials required:

Procedure:

ACTIVITY-2(B)


2. Put copper(II) sulphate solution into one of the wells in the spotting tile.

3. Hold a strip of magnesium under the surface of the liquid for about a minute.

Note down any changes in the table below. If there is no change, write No change.

4. Repeat step 3 for each of the other three metals.

5. Repeat steps 2, 3 and 4, but replace the copper(II) sulphate solution with each of the other three solutions in turn.

Metal magnesium Zinc iron copper(II) number ofsulphate sulphate sulphate sulphate times coated

Magnesium

Zinc

Iron

Copper

Analysis

1. For each metal, count up the number of times it was coated by the metal from the solution.]

Write the number in the right hand column of your table.

2. Put the metals in order from the most reactive to the least reactive; this is a reactivity series. Explain why you put the metals in this order.

3. Write word equations (and symbol equations if you can) for each of the reactions seen, e.g. for magnesium dipped in copper(II) sulphate solution the word equation is:

magnesium + copper(II) sulphate? magnesium sulphate + copper and the symbol equation is: Mg(s) + CuSO (aq) ??MgSO (aq) + Cu(s)4 4

Consider the limitations of the experiment. How could you improve it, and why?

Observations:

Evaluation:


ACTIVITY : 2C

2

2

2

Aim:

Materials required:

Procedure:

Observation:

To Study whether magnesium oxide is acidic or basic in nature.

3 cm long magnesium ribbon

Water

Blue litmus paper or solution

1. Take a 3cm long piece of magnesium ribbon.

2. Burn it in flame of candle or burner. Magnesium burns in air with a dazzling white flame. (To be performed carefully precaution for the students not to look directly into the flame while magnesium is burning.)

3. Collect the white ash in a dish. The following reaction took place while burning-

Magnesium + Oxygen → Magnesium oxide

(ash)

4. Dissolve the collected ash; magnesium oxide in water.

5. Test the solution separately with red and blue litmus paper or litmus solution.

6. Observe and record the colour change in the table below.

S.No Action on red litmus Action on blue litmus Inference(acid or base)

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ACTIVITY : 2D

2

2

2

Aim:

Materials required:

Procedure:

Observation:

To Study whether rust is acidic or basic in nature.

Rusted iron article

Water

Red and blue litmus paper or solution

1. Collect some rust by scraping it off from a rusted iron article. We have already learnt that oxygen, water and iron react to produce rust. Let us write a word equation for rusting of iron-

Iron + Oxygen + Water → Rust (iron oxide)

2. Add this in water and stir well. The rust will not dissolve in water. Let it sediment and take only the liquid from the top.

3. Test this liquid with red and blue litmus separately and record your observations in the table given.

S.No Action on red litmus Action on blue litmus Inference(acid or base)

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ACTIVITY : 2E

2

2

2

Aim:

Materials required:

Procedure:

Observation:

To Study whether sulphur dioxide is acidic or basic in nature.

Sulphur powder

Spatula/Spoon

Red and blue litmus solution or paper

1. Take some sulphur powder in a spoon/ spatula.

2. Burn it and hold an inverted test tube over it to collect the gas formed. Following reaction takes place, during burning-

Sulphur + Oxygen → Sulphur dioxide

S + O → SO2 2

4. Carefully add water to the test tube in which gas has been collected.

5. Test this solution with red and blue litmus separately and record your observations in the table given in activity 2. Instead of this you can simply bring a wet red and blue litmus paper where the gas is being formed and test whether there is a colour change or not.

6. Similarly perform test with burning a piece of charcoal and checking the action of gas evolved on wet red and blue litmus paper separately.

S.No Element Action on red litmus Action on blue litmus Inference(acid or base)

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Compare the results of activities 2c ,2d and 2E. Do we get results as given in the table below-

S.No Element Action on red litmus Action on blue litmus Inference(acidic or (metal or non metal) basic oxide)

1 Mg Turns blue No change basic oxide)

2 Fe Turns blue No change basic oxide)

3 S No change Turns red acidic oxide)

4 C No change Turns red acidic oxide)

To study the reactions of magnesium, zinc, iron and silver in copper sulphate

solution and to work out a reactivity series for them.

Copper sulphate crystals

Four disposable paper cups

Magnesium ribbon

Zinc granule

Iron nail

Silver ring

Sand Paper

1. Prepare approximately100mL copper sulphate solution by dissolving four tea spoon full (spatula) of copper sulphate in 100mL water.

2. Take four 100mL beakers or disposable paper cups.

3. Take 25mL copper sulphate solution in all the beakers. What is the colour of the solution?

ACTIVITY : 2F

2

2

2

2

2

2

2

Aim:

Materials required:

Procedure:

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4. Label them as A, B, C and D.

5. Put Magnesium ribbon in A, zinc granules in B, Iron nail in C and silver ring in D.

6. Keep the beakers undisturbed and observe carefully the activity going inside the beaker.

7. Record your observations after 10 minutes.

Clean the metal pieces with a sand paper before putting them into solution.

Put all the metal pieces at the same time in different beakers.

Beaker Metal Added Observations

Colour of solution Appearance of the metal piece

A Magnesium ribbon

(Mg)

B Zinc granules (Zn)

C Iron nail (Fe)

D Silver piece/ring

(Ag)

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

Precaution:

Observations:

Conclusion:

2

2

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ACTIVITY : 2G

2

2

2

2

2

2

Aim:

Materials required:

To describe the reactivity of metals and arrange some of the given metals in an

increasing/decreasing order of reactivity.

Solution of Zinc sulphate,copper sulphate and iron sulphate

Six Beakers

Iron filings

Copper turnings

Zinc granules

Silver piece

Procedure:

1. Prepare solutions of zinc sulphate, copper sulphate and iron sulphate.

2. Take each solution in two different beakers, and label the beakers as

A, B, C, D, E and F.

3. Put different metal pieces in these beakers as specified in the table below-

Beaker Solution Metal

A ZnSO Fe (iron)4

B ZnSO Cu (copper)4

C FeSO Zn (zinc)4

D FeSO Ag (silver)4

E AgNO Fe (iron)3

F AgNO Zn (zinc)3

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Observations:

Conclusion:

Beaker Solution Metal Colour of the Observation

solution after 10 (reaction or no

minutes reaction)

A ZnSO Fe (iron)4

B ZnSO Cu (copper)4

C FeSO Zn (zinc)4

D FeSO Ag (silver)4

E AgNO Fe (iron)3

F AgNO Zn (zinc)3

1. Which is the most reactive metal?

________________________________________________________________________

2. Which is the least reactive metal?

________________________________________________________________________

3. Arrange the metals in the order of decreasing reactivity.

________________________________________________________________________

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WORKSHEET-3.1

1. An experiment is carried out to find the relative reactivities of four metals: copper, magnesium, iron and zinc. Strips of three of the metals are placed in dilute solutions of different sulphates, as shown below, and left for the same length of time.

(a) Use the information in the table to place the four metals in order of reactivity, starting with the most reactive.

(b) Use the appropriate descriptions given in the table to help you complete the three Missing parts of the table.

(i) Zn and FeSO4: appearance of metal at end(Answer: dark grey)

(ii) Mg and CuSO4: appearance of metal at end(Answer: brown)

(iii) Mg and CuSO4: colour of solution at end(Answer: colourless)

(iv) The concentrations of the solutions are the same. Suggest which of the four experiments gives out the most energy.(Answer: Mg and CuSO4)

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2. Aluminium and tin-plated steel are used to make cans for food and soft drinks. The table below shows the pH values of some soft drinks and cooked foods.

drinks and foods pH value

Cola 2.0

Lemonade 3.0

Rhubarb 3.0

Beef 7.0

(a) Cans were first used about 150 years ago to store food for soldiers. The cans were made from unplated steel. The soldiers found that beef kept in steel cans was still good to eat after many months. However they found that steel cans of rhubarb bulged, and when the cans of rhubarb were opened a gas escaped.

(i) Why were the steel cans not suitable for storing rhubarb?

(ii) Name the gas that formed in the cans of rhubarb.

Part of the reactivity series is given below:

magnesium

aluminium

zinc

iron (steel)

tin

copper

silver

(b) In modern 'tin cans' the steel is covered with a thin layer of tin.

(i) Use the reactivity series to explain why 'tin cans' are better than steel cans for storing food.

(ii) When 'tin cans' are dented, the layer of tin often cracks. What reaction might happen when the layer of tin is cracked?

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(c) Many drink cans are now made of aluminium. Given the information in the reactivity series, why is this surprising?

3. Aisha placed small samples off four different metals on a spotting tile. She added drops of copper sulphate solution to each metal.

Aisha repeated the experiment with fresh samples of the four metals and solutions of different salts. She recorded some of her results in a table:

(a tick shows that a reaction took place, a cross shows that no reaction took place)

(a) The four metals have different reactivities. Use the information in the table to put the four metals in a reactivity series, starting with the most reactive.

(b) Use the reactivity series to work out how to fill in the empty boxes in the table. Put 'tick' or 'cross' in the boxes below for each combination of reactants.

(i) zinc and copper sulphate

(ii) iron and magnesium sulphate

(iii) zinc and magnesium sulphate

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(c) Copper reacts with silver nitrate solution.

(i) Complete the word equation for the reaction:

copper + silver nitrate ? ________________________________________________

(ii) Platinum does not react with silver nitrate. Put the metals platinum, copper and silver in the correct order according to their reactivity, starting with the most reactive.

(d) (i) In many houses the hot water pipes are made from copper and the boiler is made from iron. Which of these metals will corrode first?

(ii) Explain your answer.

4. A group of pupils placed pieces of metal wire in different salt solutions. They recorded their observations in the table below.

(a) From these observations, work out the order of reactivity of the four metals, copper, lead, silver and zinc, starting with the most reactive.

79


(b) The pupils then dipped a new piece of each of the metal wires into dilute hydrochloric acid. Only one of the metals reacted. Which metal was this?

(c) One pupil predicted that there would be no reaction when he put a piece of zinc wire into a solution of silver nitrate.

(i) Was his prediction correct?

(d) In nature, gold is never found combined with other elements.

(i) Where should gold be placed in the reactivity series in part (a)?

(ii) Explain your answer.

5. Rebecca tested four metals to see which would react with solutions of different metal nitrates. The table shows her results. If the metal reacted she put a tick (?). If nothing happened she put a cross (x).

(a) Use the information in the table to work out the reactivity series for the four metals. (Start with the most reactive.)

(b) None of the metals shown in the table reacts with a solution of magnesium nitrate.

(i) Give the name of a metal which will react with a solution of magnesium nitrate.

(ii) Give the name of a metal nitrate solution, apart from magnesium nitrate, which will not react with magnesium.

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6. The table shows some of the chemical reactions of four elements.

A None None burns to give an acidic oxide

B vigorous, hydrogen violent, gas given offburnsgiven off to give a basic oxide

C None None reacts slowly to give a basicoxide

D None violent, gas given off burns to give a basic oxide

(a) Elements B, C and D are all metals. Write the letters down in order of the reactivity of the metals starting with the most reactive.

(b) (i) Name one gas, present in the air, with which all four elements will react.

(ii) Name the gas which is produced when element D reacts with dilute hydrochloric acid.

(c) What evidence is there to show that element A is the only non-metal?

(d) Give the name of an element which could be element B.

7. Railway lines can be joined together by pouring molten iron into the gap between them.

(a) The molten iron is produced by the reaction between powdered aluminium and iron oxide. Complete the word equation for the reaction.

aluminium + iron oxide → iron + ______________________

(b) Iron can be produced from a mixture of aluminium and iron oxide but not from a mixture of copper and iron oxide. What are the names of the three metals, in the order of their reactivity, starting with the most reactive?

element reaction with water reaction wtih dilute reaction when heated in airhydrochloric acid

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(c) The list shows the names and symbols of five metals in order of their reactivity.

Sodium Na

Calcium Ca

Magnesium Mg

Zinc Zn

Silver Ag

(i) What, if anything, would be the result of heating zinc powder with calcium oxide?

(ii) Write down the name of a metal in the list that will not react with a solution of magnesium sulphate.

(d) The powdered metal with the symbol Zn burns in air. Write the word equation for the reaction.

8. 1. The table gives information about the colours of four metals.

Copper Brown

Iron dark grey

Magnesium Silver

Zinc light grey

A reactivity series of the metals is:

magnesium (most reactive)

zinc

iron

copper (least reactive)

Use this information to help you answer the questions below.

Name symbol

Metal Colour of metal

82


(a) A piece of zinc was placed in a solution of copper sulphate.

(i) Complete the following word equation:

zinc + copper sulphate ? ______________________________________________

(ii) Complete the statement about the appearance of this piece of zinc:

The light grey colour would change to a ___________________ colour

(b) Excess magnesium powder was put into a test tube containing a blue solution of copper sulphate and stirred. The solution soon turned colourless. A powder settled on the bottom of the test tube.

(i) What is the word equation for the reaction?

(ii) The powder was filtered off. Two different coloured solids could be seen. Give the two colours.

(c) (i) A piece of iron was placed in a solution of magnesium sulphate. What reaction, if any, would occur?

9. The metal chromium can be extracted industrially by three different chemical methods. The equations for these chemical reactions are shown below.

Cr O + 2Al → 2Cr + Al O2 3 2 3

2Cr O + 3Si 4Cr + 3SiO2 3 ? 2

2Cr O + 3C → 4Cr + 3CO2 3 2

(a) What name is given to the extraction of a metal from its oxide in this type of chemical reaction?

(b) Use the equations to compare the reactivity of chromium with the reactivities of aluminium, silicon and carbon.

(i) Is aluminium more or less reactive than chromium?

(ii) Is silicon more or less reactive than chromium?

(iii) Is carbon more or less reactive than chromium?

Resource: http://www.yenka.com/activities/Reactivity_of_Metals_-_Activity/

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ACTIVITY : 3

2

2

2


Where does carbon come in the reactivity series?

Teaching notes on Experiment 1 'Heating carbon with metal oxides’

The student will be able to:

Understand the position of carbon in the reactivity series

Understand the importance of carbon as a reducing agent during extraction of metals.

Predict which all metals can be reduced using carbon as reducing agent.

The position of carbon in the reactivity series, is determined by heating carbon with metal oxides and looking for evidence of a reaction. Magnesium is then burned in carbon dioxide producing carbon and magnesium oxide.

Students should see a glow in the carbon/copper oxide tube with the formation of red-brown copper. In the carbon/magnesium oxide tube, no glow is visible and the mixture looks the same (black and white particles) at the end. (Note that some references recommend testing for carbon dioxide, but heating carbon powder on its own under these conditions produces this gas.)

Carbon is above copper but below magnesium in the reactivity series.

The reaction is:

Carbon + copper oxide → copper + carbon dioxide

Copper oxide is reduced to copper by the carbon. (Reduction is removal of oxygen, at this level.)

Lead oxides are also reduced but care needs to be exercised because of the toxicity of lead.

Zinc oxide can be used as another unreactive oxide but the fact that it turns yellow on heating (but then back to white on cooling) may confuse students.

The reaction between carbon and iron oxide is a bit more subtle. There is no change in the contents of the tube but some magnetic particles are often detected. This can be presumed to be iron, so some reduction has occurred.

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Carbon + iron oxide → iron + carbon dioxide

Thus carbon is above iron in the reactivity series (but, for the relative lack of reaction) only just above.

This experiment can lead into a study of the blast furnaces. The ability of carbon to reduce metal oxides changes as the temperature rises. Thus, at a temperature of

0approximately 180 degrees Celsius in a blast furnace, carbon is more easily able to reduce iron oxide.

This can be done as a class practical (Experiment 1) and a demonstration (Experiment 2). Altogether, they should take about one hour.

Content based

To understand the position of carbon in the reactivity series.

(The position of carbon in the reactivity series, is determined by heating carbon with metal oxides and looking for evidence of a reaction. Magnesium is then burned in carbon dioxide producing carbon and magnesium oxide.)

Experiment 1 Heating carbon with metal oxides

Each working group requires:

Hard glass test-tubes, 3 (see note 1)

Bunsen burner

Heat resistant mat

Test-tube holder

Magnet

Carbon (dry powdered wood charcoal)

Magnesium oxide

Copper(II) oxide

Iron(III) oxide

Lesson organisation

Type of activity:

Aim:

Materials required:

85


Experiment 2: The reaction of magnesium with carbon dioxide

For one demonstration:

Magnesium ribbon -10 cm

Carbon dioxide cylinder (or CO gas generator) 2

Gas jar and lid

Tongs

Bunsen burner

Scissors

Experiment 1 Heating carbon with metal oxides

(a) Light a Bunsen burner.

(b) Mix together one small spatula measure of carbon powder and an equal measure of copper(II) oxide in a test-tube. Move the tube from side to side to mix the solids.

(c) Hold the tube in a test-tube holder. Heat the tube strongly with a roaring Bunsen flame. Look for any glow that persists well after the tube has been taken out of the flame. Also look for any colour change in the tube.

(d) Repeat the experiment using a mixture of carbon powder and magnesium oxide.

(e) Prepare a mixture of iron oxide and carbon as in b above.

(f) Hold the test-tube horizontally and run a magnet under the glass. See whether any part of the mixture is magnetic.

(g) Heat the iron oxide and carbon mixture strongly in the test-tube, and watch for signs of any change.

(h) When you have heated for 5 minutes, allow the tube to cool. Test for the presence of any magnetic particles, as in part f.

(i) For each experiment, record the following.

Procedure:

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Appearance of the mixture at the start (including, for the iron oxide/carbon mixture, 'is it magnetic?').

Appearance of the mixture during heating.

Appearance of the mixture after heating (including, for the iron oxide/carbon mixture, 'is it magnetic?').

Experiment 2 The reaction of magnesium with carbon dioxide

Demonstration

(a) Using a gas cylinder, or a gas generator, fill a gas jar with carbon dioxide and cover with a greased lid.

(b) Using scissors, cut a 10 cm piece of magnesium ribbon.

(c) Light a Bunsen burner.]

(d) Hold the magnesium ribbon in tongs, and place one end in a Bunsen burner flame. As soon as it ignites, remove the lid from the gas jar and quickly plunge the ribbon into the carbon dioxide. The magnesium continues to burn in the carbon dioxide, forming some black specks of carbon and white magnesium oxide.

Here are some questions to ask your students about Experiment 'Heating carbon with metal oxides'.

1. In which tube(s) does a reaction occur?

2. What signs of reaction are there?

3. What can you conclude about the positions of magnesium, iron, copper and carbon in the reactivity series?

4. Write word equations for any reactions that occurred.

5. Which substances are being reduced in these reactions?

2

2

2

WORKSHEET : 3

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ACTIVITY : 4

HOW METAL EXIST-MINERALS AND ORES

The metals placed at the bottom of the reactivity series i.e. silver, gold, and platinum generally exist in free state. The metals at the top of the reactivity series (K, Na, Ca, Mg, Al, Fe etc.) are not expected in free state due to their reactive nature. They exist in combination of other elements as oxides, carbonates, sulphates, sulphides etc.The elementary state or compounds in the form of which metals occur in nature in the earths crust are known as minerals. A particular metal can have more than one mineral.

The mineral from which we can isolate or extract a metal conveniently and economically is called an ore.( An ore is any naturally-occurring source of a metal that you can economically extract the metal from)

Thus all ores are minerals but all minerals are not ores.

Aluminium, for example, is the most common metal in the Earth's crust, occurring in all sorts of minerals. However, it isn't economically worthwhile to extract it from most of these minerals. Instead, the usual ore of aluminium is bauxite - which contains from 50 - 70% of aluminium oxide.

Copper is much rarer, but fortunately can be found in high-grade ores (ones containing a high percentage of copper) in particular places. Because copper is a valuable metal, it is also worth extracting it from low-grade ores as well.

Ores are commonly oxides - for example:

Bauxite Al O2 3

haematite Fe O2 3

Rutile TiO2

. . . or sulphides - for example:

Pyrite FeS2

chalcopyrite CuFeS2

. . . and a whole lot of other things as well.

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EXTRACTION OF METALS (METALLURGY)

CONCENTRATION OF THE ORE OR ENRICHMENT OF THE ORE

The various steps involved in the extraction of the metal from its ores followed by refining of the metal is called metallurgy.

The three main steps involved in the extraction of any metal are :

(i) Enrichment of the ore or concentration of the ore.

(ii) Extraction of the metal from the concentrated ore.

(iii) Refining of the impure metal.

This simply means getting rid of as much of the unwanted rocky material as possible before the ore is converted into the metal.

The ore, as obtained after mining from the ground, contains a large amount of earthy and sandy impurities called 'gangue'. Removal of these impurities from the ore is called concentration or enrichment of ores. It can be done in number of ways depending on the nature of impurities.The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore.

Hand picking- When ore and the impurities differ in size and shape of the particles.

Gravity separation/Hydraulic washing- When impurities and ore have different densities. (Only for oxide ores)

Froth floatation process-It is based on the difference in the wetting properties of the gangue and the ore particles. (Only for sulphide ores)

Magnetic separation-This method is used in those cases where either the ore or the impurities are of magnetic nature.

Chemical separation- It is based on the difference in the solubility of impurities and ore particles.

The ore is first crushed and then treated with something which will bind to the particles of the metal compound that you want and make those particles hydrophobic. "Hydrophobic" literally means "water fearing".

Froth flotation

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In concentrating copper ores, for example, pine oil is often used. The pine oil binds to the copper compounds, but not to the unwanted rocky material.

The treated ore is then put in a large bath of water containing a foaming agent (a soap or detergent of some kind), and air is blown through the mixture to make a lot of bubbles.

Because they are water-repellent, the coated particles of the metal compound tend to be picked up by the air bubbles, float to the top of the bath, and are allowed to flow out over the sides.

The rest of the rocky material stays in the bath.

(ii) The method used for the extraction of the metal from the concentrated ore depends upon the nature of the metal. Based on their reactivity, the metals have been grouped into the following three categories :

Metals of low reactivity (low in the activity series)

Metals of medium reactivity (In the middle of the activity series)

Metals of high reactivity (At top of the activity series)

(i)

This simply means getting rid of as much of the unwanted rocky material as possible before the ore is converted into the metal.

The ore, as obtained after mining from the ground, contains a large amount of earthy and sandy impurities called 'gangue'. Removal of these impurities from the ore is called concentration or enrichment of ores. It can be done in number of ways depending on the nature of impurities.The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore.

When ore and the impurities differ in size and shape of the particles.

Gravity separation/Hydraulic washing- When impurities and ore have different densities. (Only for oxide ores)

Froth floatation process-It is based on the difference in the wetting properties of the gangue and the ore particles. (Only for sulphide ores)

EXTRACTION OF THE METAL FROM THE CONCENTRATED ORE :

CONCENTRATION OF THE ORE OR ENRICHMENT OF THE ORE

Hand picking-

2

2

2

2

2

2

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Magnetic separation-This method is used in those cases where either the ore or the impurities are of magnetic nature.

Chemical separation- It is based on the difference in the solubility of impurities and ore particles.

The ore is first crushed and then treated with something which will bind to the particles of the metal compound that you want and make those particles hydrophobic. "Hydrophobic" literally means "water fearing".

In concentrating copper ores, for example, pine oil is often used. The pine oil binds to the copper compounds, but not to the unwanted rocky material.

The treated ore is then put in a large bath of water containing a foaming agent (a soap or detergent of some kind), and air is blown through the mixture to make a lot of bubbles.

Because they are water-repellent, the coated particles of the metal compound tend to be picked up by the air bubbles, float to the top of the bath, and are allowed to flow out over the sides.

The rest of the rocky material stays in the bath.

(ii) : The method used for the extraction of the metal from the concentrated ore depends upon the nature of the metal. Based on their reactivity, the metals have been grouped into the following three categories :

Metals of low reactivity (low in the activity series)

Metals of medium reactivity (In the middle of the activity series)

Metals of high reactivity (At top of the activity series)

Reducing the metal compound to the metal

Why is this reduction?

At its simplest, where you are starting from metal oxides, the ore is being reduced because oxygen is being removed.

Fe O F2 3 e

Al O Al2 3

2

2

Froth flotation

2

2

2

EXTRACTION OF THE METAL FROM THE CONCENTRATED ORE

removal of oxygen = reduction

removal of oxygen = reduction

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It is much more helpful to use the definition of reduction in terms of addition of electrons.

To a reasonable approximation, you can think of these ores as containing positive metal ions. To convert them to the metal, you need to add electrons - reduction.

3 –Fe + 3e Fe

3 –Al + 3e Al

2+ –Cu + 2e Cu

In some compounds the metal may not literally be present as a positive ion.

Instead, it may be part of a covalent bond - but will always be the least electronegative element present, and so will carry some degree of positive charge. That means that its oxidation state will always be positive. Reducing that oxidation state to zero (in the raw element) will always involve adding electrons.

If you aren't sure about oxidation states you could follow this link to find out about them.

If you choose to follow this link, use the BACK button on your browser to return to this page later.

There are various economic factors you need to think about in choosing a method of reduction for a particular ore. These are all covered in detail on other pages in this section under the extractions of particular metals. What follows is a quick summary.

You need to consider:

the cost of the reducing agent;

energy costs;

the desired purity of the metal.

There may be various environmental considerations as well - some of which will have economic costs.

Note:

Choosing a method of reduction

2

2

2

addition of electrons = reduction



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Carbon reduction

Carbon (as coke or charcoal) is cheap. It not only acts as a reducing agent, but it also acts as the fuel to provide heat for the process.

However, in some cases (for example with aluminium) the temperature needed for carbon reduction is too high to be economic - so a different method has to be used.

Carbon may also be left in the metal as an impurity. Sometimes this can be removed afterwards (for example, in the extraction of iron); sometimes it can't (for example in producing titanium), and a different method would have to be used in cases like this.

Reduction using a more reactive metal

Titanium is produced by reducing titanium(IV) chloride using a more reactive metal such as sodium or magnesium. As you will see if you read the page about titanium extraction, this is the only way of producing high purity metal.

TiCl + 4Na Ti + 4NaCl4

The more reactive metal sodium releases electrons easily as it forms its ions:

–4Na 4Na + 4e

These electrons are used to reduce the titanium(IV) chloride:

– –TiCl + 4e Ti + 4Cl4

Extraction of Metals low in the activity series (Cu, Hg, Ag, Au, Pt)

As these metals have low reactivity, their oxides can be reduced to metals by the action of heat alone. For example cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO), which on further heating is reduced to mercury.

2HgS(s) + 3O (g) 2HgO(s) + 2SO (g)2 2

Cinnabar ore Mercuric Oxide

2HgO(s) 2Hg (l) + O2 (g)

Mercury

Heat

Heat

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Similarly copper which is found as Cu2S in nature can be obtained from its ore by just heating in air.

2Cu S(s) + 3O (g) 2Cu O(s) + 2SO (g)2 2 2 2

Copper Pyrite ore Copper Oxide

2Cu O(s) + 2Cu S(s) 6Cu(s) +SO (g)2 2 2

Extracting Metals in the Middle of the Activity Series (Fe, Zn, Pb, Sn, Ni etc.)

These metals are found in nature in the form of their oxide, sulphide or carbonate ores. Furthr as it is easier to reduce oxides than sulphides or carbonates, therefore the sulphide and carbonate ores are first converted into the corresponding metal oxides. Thus the different steps involved for the e4xtraction of the metal from the concentrated ore are as follows:

(a) Conversion of Carbonate or Sulphide ore into metal oxide: This is done by either of the following two methods:

Calcination (For carbonate ores): It is the process of heating the ore strongly in the absence of air. The metals carbonate decomposes to form metal oxide. For example,

ZnCO (s) ZnO (s) + CO (g)3 2

Zinc carbonate (Absence of air)

Roasting (for sulphide Ores)- It is the process of heating the ore strongly in the presence of exces of air. As a result, the sulphide ore is converted into metal oxide. For example,

2ZnS(s) + 3O (g) 2ZnO(s) + 2SO (g)2 2

Zinc sulphide Presence of air Zinc Oxide

(b) Reduction of the metal oxide to metal: As these metals are moderately reactive, their oxides cannot be reduced by heating alone. Hence their oxides are reduced to metals by using a suitable reducing agent such as carbon or some highly reactive metals like sodium, calcium, aluminium etc.

Reduction by heating with carbon: Carbon combines with the oxygen of the metal oxide forming carbon monoxide. As a result, metal oxide is reduced to metal. For example oxides of iron and Zinc are reduced to their respective metals by heating with coke.

2

Heat

Heat

Heat

Heat

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2Fe O (s) + 3C(s) 2Fe(s) + 3CO(g)2 3

Ferric oxide Iron

ZnO (s) + C(s) Zn (s) + CO(g)

Zinc oxide zinc

Carbon monoxide formed also acts as reducing agent and further reduces the metal oxide to metal.

Fe O (s) + 3CO(s) 2Fe(s) + 3CO (g)2 3 2

Ferric oxide Iron

ZnO (s) + CO(s) Zn (s) + CO (g)2

Zinc oxide zinc

The reduction of metal oxides by heating with coke is called smelting.

Reduction by heating with aluminium- Oxides of certain metals e.g. manganese oxide(MnO ) chromium oxide(Cr O ) etc are easily reduced to their corresponding 2 2 3

metals by heating with aluminium powder. As a result, aluminium is converted into aluminium oxide whereas the metal oxide is reduced to the metal. The reaction is highly exothermic. The heat evolved is so high that the metal is obtained in the molten state.

3MnO (s) + 4Al (s) 3Mn (l) + Al O (s)2 2 3

Manganese dioxide Aluminium Manganese Aluminium

(powder) Oxide

Similarly, when iron (III) oxide is heated with aluminium powder ,the heat evolved is so high that iron obtained melts.

Fe O (s) + 2Al (s) 2Fe (l) + Al O (s)2 3 2 3

Iron (III) Oxide Aluminium Iron Aluminium

(powder) Oxide

This reaction is therefore used for welding the broken parts of iron machinery, railway tracks, girders etc. The reaction is known as thermite reaction.

Heat

Heat

Heat

Heat

Heat

Heat

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Extraction of Metal High Up in the Activity Series(K, Na, Ca, Mg and Al) : The highly reactive metals such as potassium, sodium, magnesium, aluminium, etc. cannot be obtained by reduction of their oxides by heating with carbon or aluminium. This is because these highly reactive metals have greater affinity for oxygen than for carbon or aluminium. Hence, these metals are obtained by electrolysis of their molten or fused oxides or chlorides. The method is called "electrolytic reduction". For example-

Electrolysis of molten sodium chloride When sodium chloride melts, it splits into sodium ion(Na+) and chloride ions(Cl-)

NaCl (s) → Na+ (l) + Cl- (l)

When electricity is passed through the molten NaCl ,Na+ ions go to the cathode whereas Cl- ions are liberated at the anode.

–At Cathode: Na + (l) + e Na (s)

Sodium ions Sodium metal

–At Anode: Cl- (g) Cl (g) + e

Chloride ion Chlorine atom

Cl + Cl → Cl2 (g)

Chlorine gas

If electrolysis of aqueous solution of sodium chloride is carried out, the sodium

metal obtained at the cathode reacts with water to form sodium hydroxide and hydrogen gas. Thus, instead of sodium metal, hydrogen gas is liberated at the cathode.

(iii) REFINING OF IMPURE METALS

The process of purifying the impure (crude) metal is called refining of the metal. The most commonly employed method for the purification of metals is Electrolytic refining. The general procedure is as follows:

(i) The impure metal is taken in the form of a thick block and made the anode in the electrolytic bath , by connecting it to the positive terminal of the battery

(ii) A thin plate of pure metal is made the cathode by connecting it to the negative terminal of the battery.

NOTE:

reduction

Oxidation

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(iii) A suitable water soluble salt of the metal is taken as an electrolyte.

When an electric current is passed through the solution, the pure metal from the anode passes into the solution in the form of metal ions and an equivalent amount of metal ions from the solution are deposited as pure metal on the cathode.

n+ –At Anode: M(s) → M (aq)+ ne

Metal atom Metal ions

At Cathode: M n+ (aq) + ne- → M(s)

Metal ions Metal atom

The soluble impurities present in the impure metal pass into solution whereas insoluble impurities fall below the anode as anode mud/anode sludge.

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K, Na, Ca, Mg, Zn, Fe, Pb Cu, Hg Ag, Au

Al

Ore Ore Ore

↓ ↓ ↓ Found in native state only needed to be purified

Crush Crush Crush

↓ ↓ ↓

Concentration Concentration of the ore Concentration of

of the ore the ore

↓ ↓ ↓

Melt Conversion of ore to its Melt

oxide by calcinations or

roasting

↓ ↓ ↓

Electrolysis of Reduction of oxide to Electrolysis of

molten ore metal molten ore

↓ ↓ ↓

Pure metal Purification to get Pure metal

pure metal

Metals of high Metals of medium Metals of low Noble metals

reactivity reactivity reactivity

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1 Carried out for carbonate ores Carried out for sulphate ores

2 Heating concentrated ore strongly in absence of air. Heating concentrated ore strongly in presence of air

3 Removal of moisture ,CO and other volatile impurities process of roasting is 2

performed to remove moisture, CO , 2

impurities of sulphur, arsenic etc.

This part looks at the use of the Blast Furnace in the extraction of iron from iron ore, and the conversion of the raw iron from the furnace into various kinds of steel.

Extracting iron from iron ore using a Blast Furnace

The most commonly used iron ores are haematite (US: hematite), Fe O , and 2 3

magnetite, Fe O .3 4

The process of the extraction of iron is carried out by the following steps:

Concentration of ore

Roasting of ore

Reduction of ore

In this metallurgical operation, the ore is concentrated by removing impurities like soil etc. The process involves the crushing and washing of ore.

The concentrated ore is now heated in the presence of air. The process of roasting is performed to remove moisture, CO , impurities of sulphur, 2

arsenic. Ferrous oxide is also oxidized to ferric oxide.

This is carried out in a blast furnance

S.No. Calcination Roasting

Iron ores

Concentration of ore:

Roasting of ore:

Reduction of ore:

IRON AND STEEL

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The Blast Furnace

The charge:

The heat source

The blast furnace is a cylindrical tower like structure about 25m to 35m high. It has an outer shell of steel. Inside of furnace is lined with fire bricks. The top of the furnace is closed by a cup-cone feeder.

The consists of :

roasted ore

Coke

Limestone

The air blown into the bottom of the furnace is heated using the hot waste gases from the top. Heat energy is valuable, and it is important not to waste any.

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The coke (essentially impure carbon) burns in the blast of hot air to form carbon dioxide - a strongly exothermic reaction. This reaction is the main source of heat in the furnace.

At the high temperature at the bottom of the furnace, carbon dioxide reacts with carbon to produce carbon monoxide.

It is the carbon monoxide which is the main reducing agent in the furnace.

In the hotter parts of the furnace, the carbon itself also acts as a reducing agent. Notice that at these temperatures, the other product of the reaction is carbon monoxide, not carbon dioxide.

The temperature of the furnace is hot enough to melt the iron which trickles down to the bottom where it can be tapped off.

Iron ore isn't pure iron oxide - it also contains an assortment of rocky material. This wouldn't melt at the temperature of the furnace, and would eventually clog it up. The limestone is added to convert this into slag which melts and runs to the bottom.

The heat of the furnace decomposes the limestone to give calcium oxide.

This is an endothermic reaction, absorbing heat from the furnace. It is therefore important not to add too much limestone because it would otherwise cool the furnace.

Calcium oxide is a basic oxide and reacts with acidic oxides such as silicon dioxide present in the rock. Calcium oxide reacts with silicon dioxide to give calcium silicate.

Flux + Gangue Slag

CaO + SiO CaSiO2 3

The reduction of the ore

The function of the limestone

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The calcium silicate melts and runs down through the furnace to form a layer on top of the molten iron. It can be tapped off from time to time as slag.

Slag is used in road making and as "slag cement" - a final ground slag which can be used in cement, often mixed with Portland cement.

The molten iron from the bottom of the furnace can be used as cast iron.

Cast iron is very runny when it is molten and doesn't shrink much when it solidifies. It is therefore ideal for making castings - hence its name. However, it is very impure, containing about 4% of carbon. This carbon makes it very hard, but also very brittle. If you hit it hard, it tends to shatter rather than bend or dent.

Cast iron is used for things like manhole covers, guttering and drainpipes, cylinder blocks in car engines, Aga-type cookers, and very expensive and very heavy cookware.

It is quite difficult to find examples of uses for cast iron, because it is nowadays

often replaced by other materials. For example, although guttering and drainpipes were once made of cast iron, apart from special old buildings, it is now quite hard to find any which aren't made of plastic!

Most of the molten iron from a Blast Furnace is used to make one of a number of types of steel. There isn't just one substance called steel - they are a family of alloys of iron with carbon or various metals. More about this later . . .

Impurities in the iron from the Blast Furnace include carbon, sulphur, phosphorus and silicon. These have to be removed.

Sulphur has to be removed first in a separate process. Magnesium powder is blown through the molten iron and the sulphur reacts with it to form magnesium sulphide. This forms a slag on top of the iron and can be removed.

Cast iron

Note:

Steel

Steel-making: the basic oxygen process

Removal of sulphur

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Removal of carbon etc

Types of iron and steel

Wrought iron

Mild steel

High carbon steel

The still impure molten iron is mixed with scrap iron (from recycling) and oxygen is blown on to the mixture. The oxygen reacts with the remaining impurities to form various oxides.

The carbon forms carbon monoxide. Since this is a gas it removes itself from the iron! This carbon monoxide can be cleaned and used as a fuel gas.

Elements like phosphorus and silicon react with the oxygen to form acidic oxides. These are removed using quicklime (calcium oxide) which is added to the furnace during the oxygen blow. They react to form compounds such as calcium silicate or calcium phosphate which form a slag on top of the iron.

Cast iron has already been mentioned above. This section deals with the types of iron and steel which are produced as a result of the steel-making process.

If all the carbon is removed from the iron to give high purity iron, it is known as wrought iron. Wrought iron is quite soft and easily worked and has little structural strength. It was once used to make decorative gates and railings, but these days mild steel is normally used instead.

Mild steel is iron containing up to about 0.25% of carbon. The presence of the carbon makes the steel stronger and harder than pure iron. The higher the percentage of carbon, the harder the steel becomes.

Mild steel is used for lots of things - nails, wire, car bodies, ship building, girders and bridges amongst others.

High carbon steel contains up to about 1.5% of carbon. The presence of the extra carbon makes it very hard, but it also makes it more brittle. High carbon steel is used for cutting tools and masonry nails (nails designed to be driven into concrete blocks or brickwork without bending). You have to be careful with high carbon steel because it tends to fracture rather than bend if you mistreat it.

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Special steels

Some environmental considerations

Environmental problems in mining and transporting the raw materials

Extracting iron from the ore

These are iron alloyed with other metals. For example:

This section is designed to give you a brief idea of the sort of environmental issues involved with the extraction of iron and its conversion to steel. I wouldn't claim that it covers everything!

Think about:

Loss of landscape due to mining, processing and transporting the iron ore, coke and limestone.

Noise and air pollution (greenhouse effect, acid rain) involved in these operations.

Think about:

Loss of landscape due to the size of the chemical plant needed.

Noise.

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2

stainless steel

titanium steel

manganese steel

iron mixed with

c h r o m i u m a n d nickel

titanium

manganese

special properties

resists corrosion

withstands high temperatures

very hard

uses include

c u t l e r y , c o o k i n g utensils, kitchen sinks, industrial equipment for food and drink processing

gas turbines, spacecraft

rock-breaking machinery, some railway track (e.g. points), military

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Atmospheric pollution from the various stages of extraction. For example: carbon dioxide (greenhouse effect); carbon monoxide (poisonous); sulphur dioxide from the sulphur content of the ores (poisonous, acid rain).

Disposal of slag, some of which is just dumped.

Transport of the finished iron.

Think about:

Saving of raw materials and energy by not having to first extract the iron from the ore.

Avoiding the pollution problems in the extraction of iron from the ore.

Not having to find space to dump the unwanted iron if it wasn't recycled.

(Offsetting these to a minor extent) Energy and pollution costs in collecting and transporting the recycled iron to the steel works.

http://www.chemguide.co.uk/inorganic/extraction/iron.html#top

Video on extraction of metals

http://www.youtube.com/watch?v=kNzc6JP9OrY

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Recycling

Resource:

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WORK SHEET : 4

Note : Answer the following questions in brief:

1. What is an Ore?

2. What is the Difference between an Ore and a Mineral?

3. What are the most common Metal Ores?

4. Are Ores a Finite Resource?

5. Are Ores Renewable? Name a Source of Metals other than Ores?

6. When is Carbon used for Extraction?

7. When is Electrolysis used for Extraction?

8. What does Native Metal mean?

9. What was the first really Useful Metal?

10. What is Haematite?

11. Which three Raw Materials are added through the Top of the Blast Furnace?

12. What is Coke?

13. Write the Balanced Equation for the Combustion of Coke

14. Write the Balanced Equation for the Reduction of Iron in the Blast Furnace

15. Which compound Reduces Iron in the Blast Furnace?

16. Why is Limestone used in the Blast Furnace?

17. What is Thermal Decomposition of Limestone.

18. Write the Balanced Equation for the Thermal Decomposition of Limestone

19. Give one Use of Slag.

20. Where does the Carbon in Cast Iron come from?

21. How can the amount of Carbon in Cast Iron be Reduced?

22. Which Alloy is the majority of Iron made into?

23. Does Cast Iron Rust more easily than Steel? Why?

24. Give one Use of Cast Iron.

25. Give one Use of Wrought iron.

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ACTIVITY : 5


CORROSION(RUSTING)

The students will be able to:

(i) Identify what is iron rust.

(ii) Demonstrate how a piece of iron can be made to rust.

(iii) Understand conditions of rusting

(iv) Show that rusting in iron can be prevented.

Rusting is a chemical process that, can take place in metals exposed to the atmosphere. Not all metals however rust. Rusting is common with the metal, iron. Certain conditions in the air around a metal have to be present for iron to rust. Iron rust which, generally appears brown, is itself a chemical compound quite different from the iron itself. The rusting of iron can be prevented. In towns and large cities where iron sheets are used for roofing houses, rusting can also take place if the iron sheets are not protected from rusting. When an iron sheet rusts, it forms a new coating around the iron itself.

The process of slowly eating up of metals due to their conversion into oxides carbonates, sulphides, sulphates, etc. by the action of atmospheric gases and moisture is called corrosion.

(i) When iron is exposed to moist air for a longtime, it surface gets covered with a coating of a brown, flaky (or non-sticky) substance called rust. Rust is mainly hydrated ferric oxide(Fe O .XH O)2 3 2

(ii) Copper objects lose their lustre or shine after sometime. This is because when copper objects remain exposed to air their surface is attacked by CO2 and water vapour present in the air forming a green coating of basic copper carbonate, CuCO .Cu(OH) .3 2

(iii) The surface of silver metal gets tarnished on exposure to air. This is due to the formation of coating of black silver sulphide (Ag S) on its surface by the action 2

of H S gas present in the air.2

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Conditions Necessary for Rusting if Iron -The two things which are necessary for rusting of irons are:

(i) Presence of air(or Oxygen) and

(ii) Presence of water(or moisture)

Prevention of Rusting : The various methods which are commonly used for preventing the rusting of iron are:

(i) Painting: The most common method used for preventing rusting of iron objects is to coat their surface with paint. By doing so, air and moisture are not allowed to come in contact with the surface of iron objects and hence rusting does not take place.

(ii) Greasing and oiling : When a thin film of grease or oil is applied to the surface of an iron object, the moisture and air cannot come in contact with it and hence rusting is prevented.

(iii) Galvanization : The process of coating iron objects with a thin layer of zinc is called galvanization. Galvanization is usually done by dipping the iron object in molten zinc. But zinc is more reactive than iron. Therefore it reacts with air, moisture and carbon dioxide of the atmosphere to form an invisible thin layer of basic zinc carbonate, ZnCO .Zn(OH) which protect it from further corrosion.3 2

(iv) Coating with tin, chromium and nickel : These metals are resistant to corrosion. Thus when a thin layer of these metals is deposited on the iron and steel objects by electroplating, they are protected from rusting.

(v) Alloying with nickel and chromium: When iron is alloyed with chromium and some carbon (Fe = 74%, Cr = 18%, Ni = 8%), stainless steel is obtained. Stainless steel resists corrosion and hence does not rust.

In this experiment students protect iron nails using a variety of methods including painting, greasing and sacrificial protection. The nails are placed in test-tubes and covered with corrosion indicator solution. This contains gelatine and so sets to a jelly-like consistency. The indicator changes colour from yellow to blue to show where rusting is taking place. By comparing the amount and position of the blue indicator on each nail, the effectiveness of the different types of protection can be assessed.

Preventing rusting

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Lesson Organisation

Type of activity:

This can be set up as either a demonstration or a class practical. Students can be told how to carry it out or left to plan it for themselves. If they are to plan it themselves then it would be a good idea to demonstrate the use of the indicator with unprotected iron nails, before they start to think about their plans.

Even as a demonstration, the majority of the work can be done by individual pupils. They can paint the nails, cover them in plastic, wrap them in another metal and so on. It is not necessary to do all the suggested tests - around six, including the control nail, will give the idea. It is probably best to include some familiar methods of preventing rust, such as painting, as well as at least one example of sacrificial protection, such as wrapping with magnesium.

Although the results are obtained quickly for a corrosion practical, it will still take around half an hour for the indicator to change and so it is advisable to have something else planned for the students to do while they wait for the colour to develop.

Content based

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ACTIVITY : 5A

Aim:

Lesson Organisation:

Materials required:

Procedure:

To put iron nails in various conditions including wet, dry, air-free and salty to find

out what causes iron to rust.

This experiment will need to be set up in one lesson and then left for more than 3 days before being re-examined. It could be left set up for longer if necessary.

Per pair or group of students:

Eye protection

Test-tubes, 4

Cotton wool

Iron nails, 4

Test-tube rack

Rubber bung to fit a test tube

Forceps

Pen or other means of labelling test tubes

Calcium chloride (anhydrous granules) (Irritant) Students with sensitive skin should be offered gloves.

Cooking oil

Deionised water

Boiled deionised water (15 min boil) (see note 1)

Sodium chloride (table salt) (Low hazard)

(a) Label the test tubes 1-4.

(b) About ¼ fill tube 1 with deionised water and add a nail.

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(c) About ¼ fill tube 2 with boiled deionised water and add a nail. Carefully pour a little oil over the surface to prevent air from reaching the water.

(d) Mix some salt with some deionised water to make a solution. About ¼ fill tube 3 with this mixture and add a nail.

(e) Put a nail into tube 4 and add about 2 cm depth of anhydrous calcium chloride granules. These absorb water. Put a bung in this tube to prevent any further water from getting in.

(f) Leave for at least 3 days and then note any changes in appearance of the nails.

Test tube Effect on nail Conditions in the test tube

1

2

3

4

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

Observations:

Conclusion:

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ACTIVITY : 5B

Aim:

Materials required:

To compare the effectiveness of different types of protection including painting,

greasing and sacrificial protection.

Eye protection

Per demonstration or pair or group of students:

Test-tubes, at least 6 (see note 1)

Test-tube rack

Iron nails, at least 6 (see note 2)

Galvanised (zinc coated) nails, 2 (see note 3)

Stainless steel nail, screw or bolt

For corrosion indicator (see note 4):

Gelatine, 5 g

Potassium hexacyanoferrate(III), 0.2 g

Cleaning solution (see note 5)

-3Dilute hydrochloric acid, 1 mol dm , (Low Hazard at concentration used), about 10 3cm

Paint and small brush or Correcting fluid (see note 6)

Magnesium ribbon (Highly flammable), about 2 cm

Zinc foil (Low Hazard), small piece

Copper foil (Low Hazard), small piece

Vaseline

Oil (see note 7)

Clingfilm or similar plastic film

Marker pen or labels for test-tubes

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Teaching notes:

After about half an hour it will be possible to see the indicator changing from the starting yellow colour to dark blue in patches on the nails. These dark blue patches indicate areas where rusting is starting.

Rusting is a complex reaction between iron, oxygen and water to form hydrated 2+iron(III) oxides. Initially iron goes into solution as Fe ions, losing electrons:

2+ –Fe(s) -> Fe (aq) + 2e

This is oxidation and occurs on the iron where the protective oxide layer is weakest or damaged. Such areas are called anodic. The Fe2+ ions combine with the indicator to form a blue solid.

2+In the absence of the indicator the Fe ions combine with OH- ions produced at cathodic areas by reduction of oxygen:

–O (aq) + 2H O(l) + 4e -> 4OH-(aq) 2 2

(By adding a few drops of phenolphthalein indicator solution when making up the gelatine mixture, so-called 'Ferroxyl indicator' is obtained. This indicator will show the cathodic areas as well, as the hydroxide ions cause the phenolphthalein to turn pink.)

The iron(II) hydroxide formed is oxidised further by oxygen, to form rust, Fe O .xH O. For more detail on the reactions involved in the rusting process, see the 2 3 2

weblink below.

Typically, the magnesium wrapped nail will rust the least. The magnesium donates electrons to the iron, which slows down the rusting process. This is effective even for the parts of the iron which are not in direct contact with the magnesium. The magnesium corrodes instead of the iron, 'sacrificing' itself. This is called sacrificial protection, and is used commercially to protect iron structures in corrosive environments.

The nail wrapped in copper will rust the most. This is due to the opposite process. The more reactive metal, iron, donates electrons to the copper and becomes electron deficient itself. This increases the rate of the rusting.

The other nails will rust in a variable way, depending on how effectively they have been coated. Any chips in the paint, or gaps in the plastic or grease, will leave some of the iron nail exposed to oxygen and water, and these will be the first areas on those

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nails to rust. It is worth making the comparison with nails in contact with magnesium, which are protected even in areas that are not directly touching the magnesium. Alloying is also an effective method of rust prevention and chips and scratches in the surface are generally not important. As a result, the stainless steel nail will generally not rust much, if at all.

Students could be asked to tabulate the results of the experiment. They could also think about where each method of rust prevention is used in real life, why that method is chosen and how effective it is.

(a) Select some nails which do not show any signs of rusting. Clean them thoroughly with the cleaning solution and dry them.

(b) Place one nail as a control into a test-tube.

(c) Treat the other nails as suggested below. Not every test needs to be done. Label each test-tube.

Wrap one nail in thin plastic film, such as 'clingfilm'.

Paint one nail and let it dry.

Coat one nail with Vaseline or other grease, or oil.

Wrap a small piece of magnesium ribbon or zinc foil around a section of a nail.

Wrap a small piece of copper foil around a section of a nail.

Place these nails in separate test-tubes.

Place a stainless steel nail, bolt or screw into a test-tube.

Place two galvanised (zinc coated) nails, one which has been scratched with a file to remove a patch of zinc coating, into a test-tube.

(c) Carefully pour the corrosion indicator into each test tube, completely covering each nail. Leave for at least half an hour.

_________________________________________________________________________

_________________________________________________________________________

Resource: http://www.practicalchemistry.org/print/experiments/preventing-rusting,251,EX.html

Procedure:

Observation:

2

2

2

2

2

2

2

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ACTIVITY : 6

2

2

2

2

Learning Objectives:

ALLOYS

Students will be able to understand -

that an alloy is a metal with some other element mixed in with it

that steel, brass, bronze, solder and amalgam are alloys

that the properties of an alloy are different from those of the parent metal

that the introduction of 'smart' alloys has increased the number of applications of alloys

An alloy may be defined as a homogenous mixture of two or more metals, or a metal and a non metal. The most important metallic component of an alloy (often representing 90 percent or more of the material) is called the main metal, the parent metal, or the base metal. The other components of an alloy (which are called alloying agents) can be either metals or nonmetals and they're present in much smaller quantities (sometimes less than 1 percent of the total). Although an alloy can sometimes be a compound (the elements it's made from are chemically bonded together), it's usually a solid solution (atoms of the elements are simply intermixed, like salt mixed with water).

Preparation of alloys- An alloy is usually prepared by first melting the main metal and then dissolving the other elements in it in a definite proportion. It is then cooled to room temperature.

An alloy containing mercury as one of the constituent metals is known as amalgam. For example sodium amalgam, zinc amalgam etc.

Objectives of alloy making-The main objectives of alloy making are :

(i) To increase hardness

(ii) To increase tensile strength

(iii) To increase resistance to corrosion

(iv) To lower the melting point

(v) To modify chemical reactivity

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(vi) To reduce electrical conductivity

(vii) To produce good casting

(viii) To modify colour.

Alloying of gold-Purity of gold is expressed in carats. Pure gold is 24 carat. It is very soft and is therefore, not suitable for making jewellery. To make it hard it is alloyed either with copper or silver. In India generally 22 carat gold is used for making ornaments. This means 22 parts by weight of pure gold is alloyed with 2 parts by weight of either copper or silver.

If you look at a metal through a powerful electron microscope, you can see the atoms inside arranged in a regular structure called a crystalline lattice. Imagine a small cardboard box full of marbles and that's pretty much what you'd see. In an alloy, apart from the atoms of the main metal, there are also atoms of the alloying agents dotted throughout the structure. (Imagine dropping a few plastic balls into the cardboard box so they arrange themselves randomly among the marbles.)

If the atoms of the alloying agent replace atoms of the main metal, we get what's called a substitution alloy. An alloy like this will form only if the atoms of the base metal and those of the alloying agent are of roughly similar size. In most substitution alloys, the constituent elements are quite near one another in the periodic table. Brass, for example, is a substitution alloy based on copper in which atoms of zinc replace 10-35 percent of the atoms that would normally be in copper. Brass works as an alloy because copper and zinc are close to one another in the periodic table and have atoms of roughly similar size.

Alloys can also form if the alloying agent or agents have atoms that are very much smaller than those of the main metal. In that case, the agent atoms slip in between the main metal atoms (in the gaps or "interstices"), giving what's called an interstitial alloy. Steel is an example of an interstitial alloy in which a relatively small number of carbon atoms slip in the gaps between the huge atoms in a crystalline lattice of iron.

The structure of alloys

Substitution alloys

Interstitial alloys

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How do alloys behave?

People make and use alloys because metals don't have exactly the right properties for a particular job. Iron is a great building material but steel (an alloy made by adding small amounts of nonmetallic carbon to iron) is stronger, harder, and rustproof. Aluminum is a very light metal but it's also very soft in its pure form. Add small amounts of the metals magnesium, manganese, and copper and you make a superb aluminum alloy called duralumin, which is strong enough to make airplanes. Alloys always show improvements over the main metal in one or more of their important physical properties (things like strength, durability, ability to conduct electricity, ability to withstand heat, and so on). Generally, alloys are stronger and harder than their main metals, less malleable (harder to work) and less ductile (harder to pull into wires).

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How are alloys made?

Note on Alloy Structure

Photo: Scientists at NASA Ames have developed a technique called high-pressure gas atomization for simplifying the production of magnesium alloys. Photo by courtesy of US Department of Energy.

You might find the idea of an alloy as a "mixture of metals" quite confusing. How can you mix together two lumps of solid metal? The traditional way of making alloys was to heat and melt the components to make liquids, mix them together, and then allow them to cool into what's called a solid solution (the solid equivalent of a solution like salt in water). An alternative way of making an alloy is to turn the components into powders, mix them together, and then fuse them with a combination of high pressure and high temperature. This technique is called powder metallurgy. A third method of making alloys is to fire beams of ions (atoms with too few or too many electrons) into the surface layer of a piece of metal. Ion implantation, as this is known, is a very precise way of making an alloy. It's probably best known as a way of making the semiconductors used in electronic circuits and computer chips. (Read more about this in our article on molecular beam epitaxy.)

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1. Shows the regular arrangement of the atoms in a metal crystal and the white spaces show where the free electrons are (yellow circles actually positive metal ions).

2. Shows what happens when the metal is stressed by a strong force. The layers of atoms can slide over each other and the bonding is maintained as the mobile electrons keep in contact with atoms, so the metal remains intact BUT a different shape.

3. Shows an alloy mixture. It is NOT a compound but a physical mixing of a metal plus at least one other material (shown by red circle, it can be another metal e.g. Ni, a non-metal e.g. C or a compound of carbon or manganese, and it can be bigger or smaller than iron atoms). Many alloys are produced to give a stronger metal. The presence of the other atoms (smaller or bigger) disrupts the symmetry of the layers and reduces the 'slip ability' of one layer next to another. The result is a stronger harder less malleable metal.

4. The main point about using alloys is that you can make up, and try out, all sorts of different compositions until you find the one that best suits the required purpose.

Some common alloys and what we use them for

There are zillions of different alloys used for zillions of different purposes. We've listed 20 of the more common (or otherwise interesting) ones in the table below. There are lots of different variations on most alloys and the precise mixture can vary widely, so the percentage figures you see quoted in different books will often not agree exactly.

Alloy Components Typical uses

Alnico Iron (50%+), aluminum (8-12%), nickel (15-25%), Magnets in louds

cobalt (5-40%), plus other metals such as copperd peakers and pickups

an titanium. in electric guitars.

Amalgam Mercury (45-55%), plus silver, tin, copper, and zinc. Dental fillings.

Babbitt metal Tin (90%), antimony (7-15%), copper (4-10%). Friction-reducing

("white coating in machine

metal") bearings.

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Brass Copper (65-90%), zinc (10-35%). Door locks and bolts,

brass musical

instruments, central

heating pipes.]

Bronze Copper (78-95%), tin (5-22%), plus manganese, Decorative statues,

phosphorus, aluminum, or silicon. musical instruments.

Cast iron Iron (96-98%), carbon (2-4%), plus silicon. Metal structures such

as bridges and heavy

-duty cookware.

Cupro-nickel Copper (75%), nickel (25%), plus small amounts of

(copper manganese. Coins.

nickel)

Duralumin Aluminum (94%), copper (4.5-5%), magnesium Automobile and

(0.5-1.5%), manganese (0.5-1.5%). aircraft body parts

, military equipment.

Gunmetal Copper (80-90%), tin (3-10%), zinc (2-3%), and Guns, decorative

phosphorus. items.

Magnox Magnesium, aluminum. Nuclear reactors.

Firework ignition

chromium (20%).

Nichrome Nickel (80%), devices, heating

elements in electrical

appliances.

Nitinol Nickel (50-55%), titanium (45-50%). Shape-memory alloy

used in medical items,

spectacle frames that

spring back to shape,

and temperature

switches.

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Solder Varies. Old-fashioned solders contain a mixture

of tin (50-70%), lead (30-50%), copper, antimony,

and other metals. Newer solders dispense with Connecting electrical

lead for health reasons. A typical modern solder components into

has 99.25 percent tin and 0.75 percent copper. circuits.

Steel Iron (80-98%), carbon (0.2-2%), plus other metals Metal structures, car

(general) such as chromium, manganese, and vanadium. and airplane parts, and

many other uses.

Steel Iron (50%+), chromium (10-30%), plus smaller Jewelry, medical tools,

(stainless) amounts of carbon, nickel, manganese, tableware.

molybdenum, and other metals.

Stellite Cobalt (67%), chromium (28%), tungsten (4%), Coating for cutting

nickel (1%). tools such as saw teeth,

lathes, and chainsaws.

Sterling silver Silver (92.5%), copper (7.5%). Cutlery, jewelry,

medical tools, musical instruments.

White gold Gold (75%), palladium (17%), silver (4%), copper Jewelry.

(18 carat) (4%)

Wood's metal Bismuth (50%), lead (26.7%), tin (13.3%), cadmium Solder, melting

(10%). element in fire

sprinkler systems.

Resource: http://www.explainthatstuff.com/alloys.html

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ACTIVITY : 6(A)

Type of activity:

Aim:

Materials required:

Content based

To experience how alloying can be used to change the properties of a metal.

(Modelling alloys with plasticine)

To make one set of models:

Different coloured lumps of plasticine (35 g),

4 Sand (12 g)

Rough paper

Access to a balance

Per pair or group:

Magnifying glass

Lesson organisation

Plasticine is mixed with varying amounts of sand and its *ductility * is measured in a simple test. The plasticine is used to represent the main metal in the sample - eg iron, and the sand an added substance - eg carbon in steel.

The practical is suitable for students of all abilities and can remain at the level of observation for less able students, or act as a springboard to explain the properties of alloys to more able student.

The model alloy is a mixture of plasticine and varying quantities of sand. These can either be made by students or made up in advance. The plasticine can be used and re-used many times, so once the mixtures are made-up for the first time - which can take a while - the actual practical work is very straightforward.

The samples can be prepared by a technician, or by the students. The samples can be used several times, so it is worth ensuring that they are colour coded so that they are easy to identify - such as all 2 g samples in blue. For the experiment to give good results, it is very important that the sand is mixed thoroughly and evenly with the plasticine. If the samples are to be stored, it is worth wrapping them in clingfilm or placing them in plastic bags, to prevent the plasticine from drying out.

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Sand from a builders merchants is the best, as the particle sizes tend to be fairly even. Sand from a fire bucket usually has a wide range of particle sizes and is also often dirty, so is not recommended.

Making the model mixture

1. Weigh out 2 g of sand onto a piece of rough paper.

2. Take one of the plasticine lumps and mould it until it is warm and malleable. Work it into a flat shape about 0.5 cm thick.

3. Sprinkle the sand onto the plasticine and roll it up. Working over the rough paper, mould and work the plasticine, adding in any sand that falls out, until the sand is evenly distributed.

4. Repeat with 4 g and 6 g of sand in the plasticine.

1. Mould each of the plasticine samples for about a minute until they are at hand temperature.

2. Shape each of the samples into cylinders of about the same size and shape - 6 cm long and 1.5 cm in diameter.

3. Hold the ends of the cylinder firmly and pull the ends of the piece of plasticine apart, slowly and steadily. If your hands fly apart it is a failed test and you need to re-mould the cylinder and try again.

4. Repeat for each specimen in turn, pulling with about the same force each time.

5. Examine the fracture surface of the plasticine with a magnifying glass.

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

Procedure:

Using the mixture

Observation:

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Extension:

Students should be able to notice that:

If the same plasticine samples are used repeatedly they always snap in the same place. To solve this simply remould the plasticine for a couple of minutes until all the sand is evenly distributed again and the sample is warm.

The plasticine and sand can be investigated further, and other factors such as temperature can be tested to see the effect they have on the samples' properties.

Another idea is to remove the 6 g sample from students at the start of the experiment. Students do the test on the 0 g, 2 g and 4g samples, and then have to predict what they expect for the 6 g sample. Alternatively, give them a 'mystery' sample with an unknown amount of sand in it (either 3 g or 5 g.) They have to determine the quantity of sand present by performing the same test and comparing the results to the other tests they have done.

Students often get the terms brittle, malleable or ductile, and strong muddled up. It is worth ensuring at some stage during the lesson that they are happy with the use of these. A material that can be stretched or drawn into wires is ductile (malleable means that it can be moulded into shape when cold). If it does not stretch but snaps instead then it is brittle. A material can be strong but brittle - and indeed many are. The opposite of ductile is brittle - not weak.

It is worth asking students to describe the results in detail, and to focus particularly on: the size of the fracture surface; what happens to the plasticine before it breaks; comparing the fracture surface with a piece of plasticine which has just been snapped.

the size of the fracture surface increases as the amount of sand increases;

the plasticine thins before it breaks;

the fracture surface has a larger amount of sand in it than a piece which has just been snapped.

When the sand is added to the plasticine, the properties of the plasticine, including its malleability and ductility, are altered. This models the way that alloying a metal will alter its properties. When a metal is alloyed, the added material interferes with the neat packing of the original, pure metal. This prevents the layers from sliding over each other and reduces the malleability and ductility.

Resource: http://www.practicalchemistry.org/print/experiments/modelling-alloys-with-plasticine,135,EX.html

2

2

2

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ACTIVITY : 6(B)

2

2

2

2

AIM:

Materials required:

Procedure:

Observations:

Conclusion:

To investigate the different melting points of solder, lead and tin.

old tin lid for heating

samples of lead, tin and solder

Bunsen burner

tripod, gauze and heatproof mat

1. Put a sample of solder, lead and tin into an old tin lid.

2. Heat strongly with a Bunsen burner. Watch what happens to each sample.

Which sample melts first? ____________________________

Which sample melts second? ____________________________

Which sample melts last? ____________________________

This experiment shows a general rule about the melting point of alloys.

Suggest what this rule is.

_______________________________________________________________________

_______________________________________________________________________

Watch Video on Alloys and its properties:

http://www.youtube.com/watch?v=-KOlIg8WLw8

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WORKSHEET : 6.1

1. Draw a line to match each of the alloys below to its use. One example has been done for you.

alloy use

aluminium alloy is strong and lightweight general construction

brass helps kill some bacteria aeroplanes

normal steel is strong and cheap statues

solder has a low melting point knives and forks

stainless steel is strong and won't rust hospital door handles

bronze is easy to cast into shapes sticking metals together

2. Draw lines to match each of the alloys below to what it is made from, what it is used for and why. One has been done for you.

alloy Contains used for why

bronze Mercury taps harder than copper water resistant

brass copper and tin joining metals soft when made hardens quickly

stainless lead and tin casting statues shiny and decorativesteel doesn't rust

amalgam copper and zinc filling teeth very low melting point

solder Iron cutlery shrinks on cooling

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3. There are 12 words from the last two lessons in this word search. See if you can find them all. The words are listed below to help you.

alloy amalgam anode bronze cathode

copper

electrolysis ore shape memory smart alloy steel tin

Resource: www.swanshurst.org/documents/none/22577_O_C2d_2.doc

s S x c r e p p o c h f

h I g j h s e f y k l p

a S s m a r t a l l o y

p Y g y u l i s d t l v

e L d m a g l a m a f k

m O n s c h s o a s h n

e R n c a p o h y r k I

m T l c t s g e l l o t

o C s b h t g d l p o o

r E u k o e y o l o a n

y L t e d e z n o r b o

x E r r e l n a u e a l

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WORKSHEET : 6.2

REVISION WORKSHEET -1

1. Jewellery and tableware are sometimes made of sterling silver. Sterling silver is 92.5% silver (Ag) and 7.5% copper (Cu). If you wanted to make 30 grams of sterling silver for a ring, how many grams of silver and copper would you need to start with?

2. Some gold used for jewellery is called '18-karat white gold'. This alloy is 75% gold (Au), 12.5% silver (Ag), and 12.5% copper (Cu). If you wanted to make 25 grams of 18-karat white gold for a ring, how many grams of each metal would you need to start with?

3. In jewellery, 14-karat gold is not pure gold. It is actually 58% gold (Au), 4 to 28% silver (Ag) and 14 to 28% copper (Cu). When you buy a 10 gram ring, how many grams of gold are you really getting?

4. Solder is used for electronic connections and for making stained glass projects. Electronic solder is specifically 63% tin (Sn) and 37% lead (Pb). To make 75 grams of solder, how many grams of each metal would you need to start with?

5. Bronze is typically used for statues and castings. If a sample of bronze was created using 85% copper (Cu), 10% zinc (Zn), and 5% tin (Sn), how many grams of each metal are required for a 150 gram statue?

1. Why is limestone added to the blast furnace?

A. to produce carbon monoxide

B. to lower the melting point of the ore

C. to react with impurities and form slag

D. to remove oxygen from the iron ore

2. Iron oxide ore in a blast furnace is mainly reduced by carbon monoxide to form iron.

(a) Fe O + b CO ==> c Fe + d CO2 3 2

Which four quantities a, b, c and d give the ratio to balance the equation?

A. 1, 2, 3, 2

B. 1, 3, 2, 3

2. Iron oxide ore in a blast furnace is mainly reduced by carbon monoxide to form iron.

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C. 2, 4, 6, 4

D. 2, 6, 2, 3

3. Which is oxidised in the blast furnace extraction of iron?

A. coke

B. air

C. limestone

D. haematite

4. Reduction of a substance can be defined as?

A. a molecule not reacting with oxygen molecules

B. an atom or molecule combining with oxygen

C. a molecule losing oxygen atoms

D. an atom gaining oxygen atoms

4. Reduction of a substance can be defined as?

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5. Which reaction produces the molecule that reduces the iron ore to iron?

A. ? iron(III) oxide + carbon monoxide ==> iron + carbon dioxide

B. ? calcium carbonate + silicon dioxide ==> calcium silicate + carbon dioxide

C. ? carbon dioxide + carbon ==> carbon monoxide

D. ? carbon + oxygen ==> carbon dioxide

6. Which chemical change corresponds to a neutralisation?

A ?calcium hydroxide reacts with acid waste gases from an industrial metal extraction process to form a harmless salt and water

B ?aluminium and oxygen are produced when an electric current is passed through molten aluminium oxide

C ?calcium carbonate forms calcium oxide and carbon dioxide in the blast furnace extraction of iron

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D ?carbon monoxide reacts with iron oxide to form carbon dioxide in the blast furnace

7. Which is formed on burning coke in the blast furnace?

A. carbon dioxide

B. calcium oxide

C. slag

D. carbon monoxide

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8. In the reactions in a blast furnace what happens to most of the carbon monoxide?

A. it reacts with silica and limestone to form the 'slag’

B. it reacts with carbon to form carbon dioxide

C. it reacts with the oxygen in the hot air to form carbon

D. it combines with the oxygen in the iron oxide ore

1. (a) Complete the following table.

Elements present Alloy formed Uses

copper and tin

iron and carbon

copper and zinc

Iron, Cobalt,Nickel and

Chromium

Iron and Carbon

(b) With reference to the list of metals below, answer the following questions.

magnesium, potassium, calcium, copper, zinc, lead, silver

(i) Arrange the metals in order of decreasing reactivity with water.

(ii) Which metal reacts violently with cold water with a lot of heat produced?

(iii) Write a chemical equation for the above reaction.

(iv) Which metal reacts very slowly with cold water but vigorously with steam?

(v) Which metal has no reaction with water or steam?

(vi) Which metal reacts very slowly with warm dilute hydrochloric acid?

(vii) Write a chemical equation for the above reaction.

REVISION WORKSHEET-2

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(c) State whether the metals of the following salts will be displaced or not (displaced/no reaction) when potassium, magnesium and silver are added separately into the salt

Solutions

Metal/Salt solution Sodium Nitrate Zinc Nitrate Lead nitrate

Potassium

Magnesium

Silver

2. (a) Draw a particulate model of the arrangement of atoms in a pure metal and an alloy.

(b) Is steel a pure metal or an alloy?

(c) Using the diagrams drawn in (a), explain why steel is stronger and harder than aluminium.

(d) State one reason why aluminium is preferred over steel in the making of soft drink cans.

(e) State one property of aluminium that differentiates it from a non-metal.

3. (a) What type of oxide do most metals form?

(b) Calcium, copper, magnesium and iron are metals.

(i) Place these metals in order of reactivity, most reactive first.

(ii) State how iron and copper each react, if at all, with steam.

(iii) Suggest how calcium reacts with steam.

(c) Aluminium is high in the reactivity series.Why do aluminium saucepans not react with steam?

4. (a) Mild steel should not be used for making knives and forks. Why?

(b) Aluminium, mild steel and nylon are used in making everyday objects. Suggest why

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(i) mild steel is used for making car bodies instead of aluminium,

(ii) both aluminium and mild steel are better than nylon for making cooking pots and frying pans.

5. The extraction of metals from their ores is of great economic importance.

(a) What type of reaction is required to produce iron from its ore?

(b) Name a metal which occurs uncombined in nature. Where does the metal occur in the reactivity series?

(c) Name two industrial methods which are mainly used in extracting metals from their ores.

(d) Of the two methods named in (c), state which one is suitable for extracting the following metals:

(i) Aluminium:

(ii) Zinc:

(iii) Sodium:

6. The diagram below shows a blast furnace.

(a) What substances are added at A other than the iron containing mineral (usually

haematite, which is largely Fe O )?2 3

(b) What is the hot gas that is passed into the bustle pipe at B?

(c) Explain the rise in temperature that occurs within the furnace at C.

(d) Explain briefly how the slag formed at E removed impurities.

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(e) Which gases leave the furnace by way of the gas outlet F ?

7. Forces are applied to the metal and its alloy.

Explain why a force applied in the directions shown will distort the shape of the pure metal but not that of the alloyed metal.

8. (a) An iron ore, A , is mixed with coke and limestone. It is loaded into a blast furnace. Hot B is blasted in through a ring of pipes at the base of the furnace. The coke burns producing gas C which is reduced by reaction with more coke to give gas D. This gas reduces the iron ore to iron. The limestone decomposes to form carbon dioxide and a white solid E which reacts with silica, SiO , to give a molten slag containing calcium silicate.2

(i) Identify A , B , C , D and E .

(ii) Since compound E is basic, what does this tell you about a chemical property of silica?

(iii) Write a balanced equation for the reaction between iron(III) oxide and carbon monoxide.

(iv) Explain why the iron(III) oxide is considered to have been reduced in this reaction.

(b) Zinc is a grey-coloured metal while copper(II) sulphate solution is blue in colour.

(i) Describe what you would see if pieces of zinc metal were placed in a solution of copper(II) sulphate.

(ii) Write the ionic equation for the reaction (include the state symbols).

You may visit the following web-link for more questions and sample answers.

Resource:

http://www.weberian.com/ChemNotes/WS13b%20Metals%20_5116_%20answer.pdf

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Parameter Beginning

(0)

Approaching

(1)

Meeting

(2)

Exceeding

(3)

Classifies elements as metal, nonmetal or semimetal (metalloid) based on its physical properties.

Classifies an element as a metal, nonmetal based on experimental observations of chemical properties.

Understands reactions of metals with oxygen, water and acids.

Arrange metals in order of their reactivity by reference to their chemical reactions with water/steam or dilute acid.

Write simple equations to descr ibe oxidat ion and reduction using both words and symbols.

Identifies the steps involved in the extraction of some metals based on their reactivity.

Identifies steps involved in extraction of iron from haemetite ore.

Describes chemical reactions w h i c h a r e n o t u s e f u l -corrosion

RUBRICS OF ASSESSMENT FOR LEARNING

Unit 2 - Metals

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Parameter Beginning

(0)

Approaching

(1)

Meeting

(2)

Exceeding

(3)

Identifies conditions necessary for the formation of rust

Acknowledges rusting as an oxidation process.

Identifies the factors which increase the rate of rusting.

Understands ways to prevent rust.

Understand that an alloy is a metal with some other element mixed in with it.

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CHEMISTRY SCIENCE UNIT-2 - NIMS Dubai – N.I. · PDF fileCHEMISTRY SCIENCE UNIT-2 Learning outcomes - Core At the end of this unit, students should be able to - Understand that every

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