Chemistry Form 4 A+ Notes

SEPT 2013

PREPARED BY:

2

CHAPTER 2: STRUCTURE OF THE ATOM

2.1 MATTER

 Any substance or material that occupies space and has mass.  Exists as a solid, liquid or gas (3 states of matter).

 Made up of particles.

 3 kinds of particles – atoms, molecules, ions  Can be divided into elements and compound.

Particles Description

Atoms  Smallest particles of an element that retain the chemical properties of the element.

 Examples : Sodium atom (Na) Zinc atom (Zn) Helium atom (He)

Molecules  Particles composed of two or more atoms.  Can be with the same or different atoms  Examples : Same atoms – Oxygen gas (O2)

Different atoms – Ammonia (NH3) Ions  Charged particles – positive or negative

 Positive charged ion (Cation) – Zinc ion (Zn2+)  Negative charged ion (Anion) – Chloride ion(Cl-)

Matter Descriptions

Elements  Particles made up of the same atoms only.  Can be in the form of atom or molecules.

 Cannot be split into two or more simpler substance by chemical means.

 Examples:

- Metallic  Copper(Cu), Iron(Fe), Gold(Au) - Non-metallic  Oxygen(O2), Sulphur(S8) Compounds  Particles made up of two or more elements.

 Can be molecules or ions.  Examples:

- Molecules  Water (H2O)

Sulphur trioxide (SO3) Tetrachloromethane (CCl4) - Ions  Sodium chloride (Na+

3 Iron(III) oxide (Fe3+, O2-)

Calcium chloride (Ca2+, Cl-)

2. Changes in states of matter  Matter can change its state.  Reversible changes.

 Exists in 3 states, solid, liquid and gas.

 During the changes, the following do not change: - Mass of particles

- Size of particles - Type of particles

 Velocity of the particle increases when - Temperature increases

- Kinetic energy increases

Melting point Boiling point

4  Sublimation can only happen to :

- Ammonium chloride (NH4Cl)

- Solid carbon dioxide / Dry ice (CO2) - Iodine (I2)

 Differences between solid, liquid and gas (Kinetic Theory Of Matter): (Essay)

States of matter Solid Liquid Gas

Arrangement of

particles compact, orderly

manner

Loosely packed, disorderly manner

very far apart, random motion

Particles motion

Vibrate, rotate in a fixed position

Move freely Move freely and randomly

Particles Kinetic energy

Very low Moderate High

Shape Fixed Not fixed (follow

the shape of container)

Not fixed (follow the shape of

container)

Volume Fixed Fixed Not fixed

Rate of diffusion Low Average High

Attractive forces between particles

5 3. Experiment (PeKa)

a. Heating curve of naphthalene/acetamide  Diagram:

 Graph:

AB: Solid

DE: Liquid + Gas

BC: Solid + Liquid

EF: Gas

6  Explanation:

AB: When the solid is heated, heat energy is absorbed. This causes the particles to gain kinetic energy and vibrate faster.

BC: The temperature remains constant because the heat energy absorbed by the particles is used to overcome the forces between particles so that the solid can turn into a liquid. At this temperature, both solid and liquid are present.

CD: The particles in liquid naphthalene absorb heat energy and move faster.  During the heating of naphthalene:

- Water bath is used (ensure uniform heating, naphthalene is flammable) - Naphthalene is stirred continuously (ensure an even heating)

 Water bath: For heating a substance which is less than 100°C.  Oil bath: For heating a substance which is more than 100°C.

 Latent heat of fusion: heat required to convert solid to liquid without a change in temperature.

b. Cooling curve of naphthalene/acetamide  Diagram:

7  Graph:

 Explanation:

RS: When the liquid is cooled, the particles in the liquid lose their kinetic energy. They move slower as the temperature decreases.

ST: The temperature of naphthalene remains constant because the

heat loss to the surroundings is balanced by the heat energy given off

during freezing.

TU: The particles in solid naphthalene release heat energy and vibrate

slower.

 During the cooling of naphthalene:

 Boiling tube containing naphthalene is placed in a conical flask. (to minimize heat loss which may affect the accuracy of freezing point – air trapped in conical flask is poor conductor of heat)

 Stirred by using thermometer (to ensure even cooling) PQ: Gas ST: Solid + Liquid

QR: Liquid + Gas TU: Solid RS: Liquid

8  Super cooling

i. Condition in which the temperature of a cooling liquid drops below the normal freezing point.

ii.

.

2.2 ATOMIC STRUCTURES

1. Historical development of the structure of atom a) John Dalton

- All elements made up of small indivisible particles called atoms. - Atoms made up of tiny particles which cannot be created or destroyed. - Atoms of same element – same mass

- Atoms of different elements – different mass

- Atoms join together to form larger molecules or compounds (in simple ratio)

- Weakness:

 Atoms are not the simplest particles – bigger than proton, neutrons and electron

 Atoms can be destroyed or breakdown – radioisotopes  Atoms of same element have different mass – isotopes

9 b) J.J. Thomson

- Plum pudding model.

- Electron embedded in a sphere of positive charge.

- Electron spreads randomly throughout the positive charge. c) Ernest Rutherford

- All positive charge of an atom is concentrated in the nucleus – contain protons.

- Mass of atom is located in a small area (nucleus). - Number of protons = number of electron

d) Neils Bohr

- Electrons of atom are arranged and move around the nucleus in orbital called electron shells.

- Nucleus contains protons.

10 e) Sir James Chadwick

- Discovered neutrons which are located in the nucleus. - The neutral particle has the same mass as protons. 2. Atomic Structure

 Made up of subatomic particles; protons, electrons and neutron.  Nucleus – situated at the centre of atom.

– has positive charge, protons. Neutrons may also present.  Electrically neutral. (Number of proton = Number of electrons)

 Have electrons which move around the nucleus in its shells.  Mass of proton = mass of neutron

 Nucleus contributes a lot of mass in an atom.

Subatomic particles

Symbol Relative atomic

mass (RAM) Charge Proton p 1 + Neutron n 1 neutral Electron e˗ ˗ 3. Electron Configuration

 Maximum number for each shell:  First shell : 2 electrons

 Second shell : 8 electrons  Third shell : 8 electrons  Forth shell : 2 electrons

11 4. Atomic number & Mass number

 Atomic number = proton number

 Nucleon number = proton number + number of neutrons  Mass number = Nucleon number

2.3 KINETIC THEORY OF MATTER

 Matter consists of tiny and discrete particles.  Particles always move randomly.

 There are forces of attraction between the particles.

 Particles gain kinetic energy and move faster when heated.  Particles lose kinetic energy and move slower when cooled.

 Can be proven by using 2 experiments: Diffusion and Brownian movement.

2. Diffusion

 Occurs when particles of a substance move in between the particles of another substance.

 Random movement of particles from a high concentration region to a lower concentration region.

 Happens in three states of matter; solid, liquid and gas.  Occurs most rapidly in gases, followed by liquid and solid.  Particles diffuse from one medium to another.

 Rate of diffusion increases with the temperature.

 Rate of diffusion decreases when the mass of matter increases.  Diffusion in gases:

12  Diffusion of liquid:

(Blue)  Diffusion of solid:

3. Brownian movement

 Random movement that is shown when colliding with other particle.  Can only be observed under a light microscope.

 Supports the Kinetic Theory Of Matter.

13

2.4 ISOTOPES

1. Atoms of same element with the same number of protons but different number of neutrons.

Isotopes

Uses

Carbon-14  determination of age of carbon-containing artifacts  as a biological tracer, for example, in studies of

photosynthesis

Oxygen-18  biological tracer, for example, in studies of photosynthesis Sodium-24  Detect location of leaks in water pipes,

 studies of body electrolytes

Magnesium-27

 location of leaks in water pipes

Cobalt-60  cancer treatment as tumour cells tend to be more susceptible to radiation than other cells

Krypton-81  lung ventilation studies

Technetium-99

 Medical tracer used to locate brain tumours and problems with the lungs, thyroid, liver, spleen, kidney, gall bladder, skeleton, blood pool, bone marrow, salivary

 to detect infection

Iodine-131  Medical tracer

 treat the thyroid gland &amp

 used in the diagnosis of adrenal medulla

 for imaging suspected neural crest and other endocrine tumours

Iodine-123  used in imaging to monitor thyroid function  detect adrenal dysfunction

14

Americium-241

 Domestic smoke alarms

Phosphorus-32

15

CHAPTER 3: CHEMICAL FORMULAE AND EQUATIONS

3.1 Formula and Chemical Equations

1. Reasons of comparing relative atomic mass(R.A.M, Ar) with one carbon-12 atom:

 Solid and easily handled.  Most abundant carbon isotope.  Easily available.

 Used as a reference standard in spectrometer. 2. Formulae:

3.2 The Mole and the Volume of Gas

 Number of particles in one mole of substance.  6.02 × 1023

2. Standard Temperature and Pressure (S. T. P.)  Temperature = 0°C

MASS OF SUBSTANCE, g

NO. OF MOLES, mol VOLUME OF SOLUTION, cm3 NO. OF PARTICLES, atoms × M.V. ÷ M.V. ÷ NA × NA × M.M. ÷ M.M.

16  Pressure = 1 atmosphere / atm

 Molar volume of 1 mole of gas = 22.4dm3 or 22400cm3 3. Room condition (R. T. P.)

 Room temperature = 25°C  Pressure = 1 atmosphere / atm

 Molar volume of 1 mole of gas = 24dm3 or 24000cm3

3.3 Molecular Formula and Empirical Formula

1. Molecular formula: actual number of atoms in each element that present in a molecule of the compound.

2. Empirical formula: simplest whole number ratio of atoms of each element in the compound.

3. Empirical formula =

4. Example : Glucose – M.F.: C6H12O6 – E.F.: CH2O

5. Determining empirical formula by using table form:

Element - Mass/Percentage x No. of mole _y( Ratio Empirical formula

4. Experiment for empirical formula:

For higher reactivity of metal (Mg, Zn, Ca, Al)

Metal tape Crucible with lid

17 Precaution:

- Lift the lid at intervals to allow oxygen gas to enter for combustion of metal.

- Lid is closed immediately after it is lifted to prevent white fume from escaping to the surroundings.

- Stop heating the metal when it is started to glow.

Reactive metal: both reactant and products are solid and thus, the individual mass of metal and oxygen cannot be determined.

For lower reactivity of metal (Cu, Sn, Pb, Ag)

Chemical used to dry hydrogen gas: Anhydrous cobalt chloride / anhydrous calcium chloride.

Hydrogen gas is flowed through the apparatus throughout the experiment to prevent the air for entering it.

18

CHAPTER 4: PERIODIC TABLE OF ELEMENTS

4.1 INTRODUCTION TO PERIODIC TABLE

1. Classifications of elements with the same chemical properties are placed in the same group.

2. Elements in:

 Group 1 – Alkali metals  Group 2 – Alkali earth metals  Group 3 – 12 – transition elements  Group 17 – Halogens

 Group 18 – Noble gases

 Group 1, 2, transition elements and 13 – metals  Group 15, 16 and 17 – non metals

 Same group – same chemical properties and valence electrons No. of

valence electrons

1 2 3 4 5 6 7 8 / 2

Group 1 2 13 14 15 16 17 18

3. Historical Development of the Periodic Table i. Antoine Lavoisier

19  Classify elements into 4 groups which are gases, metals, non-metals and

metal oxide.

 Not accurate – heat and light are included as gases. ii. Johann W. Dobereiner

 Classify elements with the same chemical properties into groups of three (triads).

 Discover relationship between R.A.M. in each triad. (Middle R.A.M. = average R.A.M.)

iii. John Newlands

 Arrange elements in order of increasing nucleon number (mass number) in horizontal rows. Each row has 7 elements.

 Law of Octaves – every eighth element have similar chemical properties. – Only accurate for the first 16 elements.

 Discover the existence of periodic pattern. iv. Lothar Meyer

 Volume of an atom =

of an element

 Graph of volume of atoms against their R.A.M.  Show the properties of elements recur periodically.

v. Dmitri Mendeleev

 Arrange element in order of increasing atomic mass.  Left gaps for elements yet to be discovered.

vi. Henry G. J. Moseley

i. Different element with high energy electrons & measured the frequency of the X-ray emitted by inert gases elements.

ii. Graph of square root of frequency against proton number. a) Group 18 elements – Noble Gases

 Made up of Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn).

 Exist in monoatomic form.

 Has stable electron arrangement (outermost shell filled with the maximum number of electrons).

 Chemically unreactive (do not share, donate or accept electrons).  Duplet – Helium, Octet – other noble gases.

 Physical properties:

20  Low boiling and melting point (weak Van der Waals forces /

intermolecular forces of attraction.  Do not conduct electricity.

 Low density (atoms are far apart).  Going down the group,

Melting & Boiling point Atomic size

Forces of attraction between atoms Heat energy

Density Atomic mass  Uses:

Helium  Fill airship, bicycle tyres of Olympic cyclist & meteorological balloons.

 Exist in the gas in diver’s oxygen tank. Neon  Advertising boards / lights.

 Electric discharge through glass tubes produces a red light.

Argon  Electric light bulb.

 Carrier gas in gas-liquid chromatography. Krypton  Laser light

 Flash lamps of a light house Radon  For cancer treatment.

Xenon  For flash lamp. b) Group 1 elements – Alkali metals

 Made up of Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), and Francium (Fr).

 Has 1 valence electron.

 Very reactive to become positive ions (easily to donate valence electron).  Physical properties:

 Soft metal with shiny and silvery surfaces (can be cut by knife).  Good electric and heat conductor.

21  When going down the group,

Melting and boiling point Metallic bond

Forces of attraction Atomic size

Density

Number of occupied shell  Chemical properties:

 Have same chemical properties.

 Electropositivity: measurement of ability of an element to lose an electron and form a positive ion.

 Good reducing agent.  Can be oxidised easily.

 Going down the group, reactivity / electropositivity increases.  Safety precautions when handling Group 1 elements:

 Kept in paraffin oil.  Use forceps to take them.  Wear safety goggles & gloves.  Reactions:

a) Alkali metal + water

 Hydroxide solution produced will turn red litmus paper red.  Products: metal hydroxide + hydrogen gas

Lithium 2Li + 2H2O  2LiOH + H2 Moves slowly with “hiss” sound. Sodium 2Na + 2H2O  2NaOH + H2

Moves quickly and randomly with loud “hiss” sound. Potassium 2K + 2H2O  2KOH + H2

Burns with reddish-purple light, jumps, “hiss” and “pop” sound. Increases

Decreases

Alkali metal

22 b) Alkali metal + oxygen

 Products: metal oxide (white powder).

 When metal oxide dissolves in water, it turns phenolphthalein indicator red (presence of OH- ions – alkaline)

Lithium 4Li + O2  2Li2O

Burns slowly with red light. Sodium 4Na + O2  2Na2O

Burns quickly and brightly with yellow light. Potassium 4K + O2  2K2O

Burns very quickly and brightly with reddish-purple light. c) Alkali metal + halogen gas (Chlorine & Bromine)

 Products: metal halides (metal bromide / chloride – white powder)

23 Lithium 2Li + Cl2  2LiCl / 2Li + Br2  2LiBr

Burns slowly with reddish flame. A white solid is obtained. Sodium 2Na + Cl2  2NaCl / 2Na + Br2  2NaBr

Burns brightly with a yellowish flame. A white solid is obtained. Potassium 2K + Cl2  2KCl / 2K + Br2  2KBr

Burns very brightly with a purplish flame. A white solid is obtained.

4.2 HALOGEN

1. Group 17 elements (Halogens)

 Made up of fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At).

 Exist in diatomic molecules.  Non – metal.

 Physical properties:

 Heat and electrical insulator.

 Low melting and boiling point (weak forces between the molecules).  When going down the group,

 Atomic size

 Van der Waals forces

 Heat energy used to overcome forces  Boiling and melting point

 Colour of halogen darker. 2. Chemical properties:

 Same chemical properties (same valence electrons – 7)  High electronegativity

 When going down the group,  Reactivity/electronegativity  Van der Waals forces

 Tendency to accept electron  Solubility

 Atomic size

 Distance between the nucleus and outermost shell  Reaction:

 Halogen + water  Product: two acids.

24  Halogens act as bleaching agent, except iodine water.

 In general, X2 + H2O  HX + HXO, where X is halogen.  Chlorine water

 Turn blue litmus paper red then decolourise it.

 Prepared from the reaction between potassium manganate (VII) chips with concentrated hydrochloric acid.

16HCl + 2KMnO4  2MnCl + 8H2O + 5Cl2 Chlorine (gas)  Cl2 + H2O  HCl + HClO

 Products: Hydrochloric acid + Hypochlorous acid (bleaching agent).

 Greenish–yellow gas dissolves quickly to form a light yellow solution.

Bromine (liquid)  Br2 + H2O  HBr + HBrO

 Products: Hydrobromic acid + hypobromous acid (bleaching agent).

 Reddish–brown liquid dissolves slowly, forming a brownish–yellow solution.

Iodine (solid)  I2 + H2O  HI + HIO

 Products: Hydroiodic acid + hypoiodous acid (bleaching agent).

 Very little purplish–black solid dissolves, forming a light yellow solution.

 Halogen + Sodium hydroxide solution, NaOH.

 Products: Sodium halide + Sodium halite(I) + Water

 In general, X2 + 2NaOH  NaX + NaOX + H2O, where X is halogen.

Chlorine Cl2 + 2NaOH  NaCl + NaOCl(sodium chlorate) + H2O

Greenish-yellow gas dissolves quickly to form a colourless solution. Bromine Br2 + 2NaOH  NaBr + NaOBr(sodium bromate) + H2O

Reddish-brown liquid dissolve averagely to form a colourless solution.

Iodine I2 + 2NaOH  NaI+ NaOI(sodium Iodate) + H2O

Purplish-black solid dissolves slowly to form a colourless solution.  Halogen + Iron (Fe)

 Product: iron(III) halides (brown solid)  Apparatus set up:

25  Soda lime: mixture of calcium hydroxide and sodium hydroxide

(absorb excess halogen gas)

 Iron wool is heated strongly until red hot.

 Concentrated hydrochloric acid is added to potassium manganate (VII) through a thistle funnel (to produce chlorine gas).

 In general, 3X2 + 2Fe  2FeX3, where X is a halogen. Chlorine 3Cl2 + 2Fe  2FeCl3

Iron wool burns lights up strong and bright. Brown solid is formed. Bromine 3Br2 + 2Fe  2FeBr3

Iron wool glows moderately bright and less vigorously. Brown solid is formed.

Iodine 3I2 + 2Fe  2FeI3

Iron wool glows dimly and slowly. Brown solid is formed. 3. Precaution:

 Halogens are poisonous gas.

 Must be handled in fume chamber.

 When handling halogens. Safety goggles and gloves must be used.  Fluorine is a radioactive substance, astatine is radioactive.

26

4.3 Elements in a Period

1. Period: horizontal row in the Periodic Table. 2. There are 7 periods in the modern periodic table. 3. When it goes across the period from left to right:

 Electronegativity  Proton number  Valence electrons  Non-metallic properties

 Nuclei attraction on valence electrons  Atomic size

 Electropositivity

 Metallic properties (Metallicity)

Element Na2O MgO Al2O3 SiO2 P4O10 SO2 Cl2O7

Characteristics Basic oxides

(Alkali)

Amphoteric oxides

Acidic oxides 4. Amphoteric oxides: react with both acids and alkalis, have base and acidic

properties. ( acid, alkali)

5. Sodium, Magnesium and Aluminium :  Metal

 Strong metallic bonds

 High melting and boiling points  High strength of metallic bond 6. Silicon

 High melting and boiling points

 Has strong covalent bond, forming a 3-dimensional gigantic network.

7. Uses of semi-metals/metalloids(element with properties intermediate between those of metals and non-metals)

 Silicon and Germanium – makes diodes and transistor/switch  Conductivity increases with temperature.

 Important in microelectronic industry

Decreases

27 8. Transition Element

 Elements between Group 3 until Group 12.  Metals

 Show metallic properties:  Shiny surface

 Ductile  Malleable

 Can withstand high tension  High melting and boiling point

 High density (big atomic mass despite small radius)  Electric and heat conductor

 Form coloured compounds or ions

Transition elements Colour

Chromium ion, Cr3+ Green

Iron(II) ion, Fe2+ Green

Iron(III) ion, Fe3+ Brown

Copper(II) ion, Cu2+ Blue

Chromate ion, CrO2-4 Yellow

Manganese ion, Mn2+ Pink

Cobalt ion, Co2+ Pink

Nickel ion, Ni2+ Green

Manganate ion, MnO2-4 Purple

Dichromate ion, Cr2O2-7 Orange

 Act as catalyst to speed up the reaction.

 Iron, Fe – Haber process (producing ammonia, NH3).

 Platinum, Pt – Ostwald process (producing nitric acid, HNO3).  Nickel, Ni – manufacture of margarine.

 Vanadium (V) oxide, V2O5 – Contact process (producing sulphuric acid, H2SO4).

 Form complex ions.

 Polyatomic anion/cation consisting of more than 2 metal ions with other group bonded to it.

 Examples – hexacyanoferrate (II) – [Fe(CN)6] 4-– Tetramine copper (II) 4-– [Cu(NH3)4]2+

28  Have different oxidation number.

 Iron, Fe - +2, +3, +1

 Manganese - +1, +2, +3, +6, +7  Nickel - +2, +3

 Chromium - +2, +3, +6  Give colour to precious stone.

 Presence of ions in a solution can be confirmed by using sodium hydroxide solution, NaOH / ammonia solution, NH3.

 The ions of transition elements will react with hydroxide ion, OH- to form coloured solution / precipitate.

Precious stone Colour Transition elements

Emerald Green Nickel, Iron

Ruby Red Chromium

Sapphire Blue Iron, Titanium

29

CHAPTER 5: CHEMICAL BONDS

5.1 Formation of Chemical Bonds

 Metal element reacts with non-metal element.  Metal element (Group 1, 2, and 13)

 Non-metal element (Group 16 and 17)

 Metal elements donate electrons and produce positive ions.

 Non-metal elements will accept electrons to achieve a stable electron configuration and produce negative ions.

 These ions will attract each other by a strong electrostatic force of attraction (ionic bond).

 Examples: sodium chloride, magnesium oxide, lithium oxide.

Elements which are reacting Formula of ionic compound Metal M Non-metal X Group 1, M+ Group 15, M3- M3X Group 1, M+ Group 16, M2- M2X Group 1, M+ Group 17, M- MX Group 2, M2+ Group 15, M3- M3X2 Group 2, M2+ Group 16, M2- MX Group 2, M2+ Group 17, M- MX2 Group 13, M3+ Group 15, M3- MX Group 13, M3+ Group 16, M2- M2X3 Group 13, M3+ Group 17, M- MX3  Sodium chloride, NaCl

30

4.2 Covalent Bonds

1. Formed by non-metal elements form Group 14, 15, 16, and 17.

2. Atoms of non-metals will combine to donate one, two or three valence electrons to be shared.

3. 3 types of covalent bonds:

 Single – sharing one pair of electrons  Double – sharing two pair of electrons  Triple – sharing three pair of electrons 4. These will form covalent compound. 5. Examples:

Chlorine molecule, Cl2 (Single)

Water molecule, H2O (Single)

Carbon dioxide molecule, CO2 (Double)

31 Non-metal elements which combined Molecular

Formula Element P Element Q Group 14, P4+ Group 17, Q- PQ4 Group 14, P4+ Group 16, Q2- P2Q4 / PQ2 Group 15, P3+ Group 17, Q- PQ3 Group 16, P2+ Group 17, Q- PQ2

6. Physical properties of ionic compounds:  High melting and boiling point  Conducts electricity

 Soluble in water, insoluble in organic solvents  Able to ionize in water.

 Has strong electrostatic force of attraction

 Need a lot of heat energy to overcome the forces  Arranged in lattice structure in solid state

 Contain free-moving ions that carry charges 7. Physical properties of covalent compounds:

 Low melting and boiling points

 Has weak Van der Waals forces – less heat energy is needed.  Insoluble in water, soluble in organic solvent

 Cannot conduct electricity

32 8. Giant molecules covalent compounds:

 Strong covalent bonds combine all atoms in a three-dimensional lattice structure.

 Have high melting and boiling point  Unable to conduct electricity.

 Examples: silicon, graphite, silicon oxide, diamond, protein 9. Covalent compound as organic solvents

 Water

 Dissolves all types of food – sugar and salt  Dissolves food substances in the body  Cleanses or gets rid of dirt

 Organic solvent

 Ethanol – preparation of shellac, lacquer, paint, cosmetic and perfumes

 Petrol / kerosene – cleans greasy and oily dirt stains  Propanone – nail varnish

 Chlorofluorocarbon – cleans circuit board of computer

5.4 Comparisons between ionic and covalent bond

Ionic compound Properties Covalent compound NaCl, MgO, ZnCl2, CuO Examples O2, CO2, H2O, N2, Cl2 High – has strong

electrostatic forces of attraction (A lot of heat energy is needed)

Melting and boiling points

Low – has weak intermolecular forces of attraction (Little heat is needed)

Soluble in water, insoluble in organic solvent.

Solubility Soluble in organic solvent, insoluble in water.

Conduct electricity in both molten and aqueous state – contain free– moving charged ions.

Electrical conductivity Do not conduct electricity.

33 Ionic compound Covalent compound Metal

Simple Giant

Examples Cu, Zn, Na,

Ca, Pt, Ni, Mg

M/P & B/P High Low High High Solubility Soluble in water, insoluble in organic solvents Insoluble in organic solvent, soluble in water Insoluble in both Insoluble in both Electrical conductivity Conduct in molten state or aqueous solution

Does not Does not Conduct in solid or liquid

34

CHAPTER 6: ELECTROCHEMISTRY

6.1 Electrochemistry

1. Electrochemistry: study of the interconversion of chemical energy and electrical energy.

2. Electrolyte: chemical substances that can conduct electricity in molten or aqueous form.

Examples:

 Molten potassium iodide, KI  Molten lead(II) chloride, PbCl2  Molten aluminium oxide, Al2O3  Sulphuric acid solution, H2SO4  Copper sulphate solution, CuSO4  Sodium chloride solution, NaCl

3. Non-electrolyte: chemical substances that cannot conduct electricity either in molten or aqueous form as they have no free-moving ions.

Examples:  Sulphur  Wood

 Molten sugar  Naphthalene

 Covalent compounds except ammonia and hydrogen chloride

4. Conductor: substances that can conduct electricity in liquid or solid state (not regarded as electrolyte as they are not decomposed)

 Copper  Iron  Platinum  Silver

5. Electrolysis: process whereby a compound is separated into its constituent elements when an electric current passes through an electrolyte.

35 Set up of apparatus:

Electrolysis of molten compound Electrolysis of aqueous solution

6. 2 types of electrodes: a) Active electrode

- do not react with electrolytes

- do not involve in chemical reactions - Carbon, platinum and graphite electrodes b) Inert electrode

- react with electrolytes

- involves in chemical reactions

- Copper, silver, or mercury electrodes

7. Anode: electrode that connect to the positive terminal of battery. 8. Cathode: electrode that connect to the negative terminal of battery. 9. Anion: negatively charged ions and attracted to anode.

10. Cation: positively charged ions and attracted to cathode. 11. Half equation:

Positive ions (Cations) Negative ions (Anions) K+ + e  K 2F- - 2e  F2 2F-  F2 + 2e Na+ + e  Na Ca2+ + 2e  Ca 2Cl-  Cl2 + 2e Mg2+ + 2e  Mg 2I-  I2 + 2e Al3+ + 3e  Al 4OH-  2H2O + O2 + 4e Zn2+ + 2e  Zn 2O2-  O2 + 4e Fe2+ + 2e  Fe 2Br-  Br2 + 2e Sn2+ + 2e  Sn Pb2+ + 2e  Pb 2H+ + 2e  H2 Cu2+ + 2e  Cu Ag+ + e  Ag

36 12. Electrolysis of molten compounds

Metal Observation

Sodium Shiny grey solid is formed. Lead Shiny grey solid is deposited. Nickel Shiny grey solid is formed. Copper Brown deposit is formed.

Gas Observation

Bromine Brown gas is produced. (pungent smell) Iodine Purple gas is produced.

Chlorine Yellowish-green gas is produced.

Oxygen Colourless gas bubbles are formed. (effervescence) Hydrogen Colourless gas bubbles are formed. (effervescence)

6.2 Electrolysis of Aqueous Solution

 Produced when solute is dissolved in water.

 Electrolyte containing cations, anions, H+ and OH- ions. 2. During electrolysis of aqueous solution:

 2 cations are attracted to cathode (-).  2 anions are attracted to anode (+).

 Only one of the four ions will be chosen to be discharged at anode and cathode.

3. Factors affecting which ions are chosen to be discharged:  Position of ions in the electrochemical series (ECS)  Concentration of ions in the solution

 Type of electrodes used

Test for Oxygen gas, O2

A glowing splinter is placed near the mouth of the test tube containing oxygen gas. It will light up.

Test for Hydrogen gas, H2

A lighted splinter is placed near the mouth of the test tube containing oxygen gas. A “pop” sound is produced.

37 4. Position of ions in the electrochemical series (ECS)

Cations Anions K Na Ca Mg Al Zn Fe Sn Pb H Cu Ag Kalau Nak Cari Minum Air Zappel Free Sila Pergi Hotel Curi Agar F -SO4 2-NO3 -Cl -Br- I -OH -Father Say Nothing Can Buy Indian Oranges

 The lower the position of the ion, the higher the tendency of the ions to be discharged.

 Sulphate ion, SO42- and nitrate ion, NO3- cannot be discharged. 5. Concentration of ions

 The anions in a lower concentration solution will be chosen to be discharged. (diluted)

 The cations in a higher concentration solution will be chosen to be discharged.

 Diluted – 0.0001, 0.001, 0.01 dm-3  Concentrated – 0.1, 1.0, 2.0 dm-3

 K+ and Na+ cannot be discharged even if their concentration of the solution is high.

6. Types of electrodes used

 Inert electrodes: Carbon, graphite and platinum (Both of these electrodes do not react with the electrolytes or products of electrolysis)

 Active electrodes: Silver, copper and nickel (Active anode ionises and concentration of cations in the electrolyte does not change)

6.4 Application of Electrolysis

 Objectives: to prevent corrosion / to improve appearance.

 Plating metals: gold (Au), Platinum (Pt), Chromium (Cr), copper (Cu), Silver (Ag), & Nickel (Ni).

38  Conditions:

 Object to be plated  cathode  Electroplating metal  anode

 Electrolyte used must contain the metal ions.  Surface of electroplating metal must be cleaned.  Set-up apparatus:

2. Extraction of metals

 Reactive metals (Na, Ca, Mg, Al) are extracted from their ores compounds using electrolysis.

 These metals cannot be extracted by reduction using carbon. a) Extraction of aluminium metal from bauxite (aluminium oxide)

 Cryolite is added to bauxite to lower the temperature of bauxite from 2000°C to 950°C.

 Bauxite dissociates. Al2O3  2Al3+ + 3O

39  Half equation at anode : 2O2-  O2 + 4e

 Overall equation : 4Al3+ + 6O2-  4Al + 3O2

 Carbon electrodes react with the oxygen gas to produce carbon dioxide.

 Hall Heroult’s Process.

b) Extraction of sodium metal from sodium chloride  Iron : cathode

 Carbon : anode  Set – up apparatus:

 Calcium chloride is added to lower the melting point of sodium chloride.

 Half equation at cathode : Na+ + e  Na  Half equation at anode : 2Cl-  Cl2 + 2e  Overall equation : 2Na+ + 2Cl-  2Na + Cl2  Downs’ Process

3. Purification of metals

 Impure metal containing impurities can be purified.  Conditions:

 Impure metal : anode  Pure metal : cathode

40  Set-up apparatus:

 Observation:

 Copper anode becomes thinner and the impurities are deposited below it.

 Copper cathode becomes thicker.

 Intensity of blue solution remains the same. Rate of formation of copper(II) ions of anode = rate of discharge of copper(II) ions of cathode. Concentration remains the same.

 Half equation at anode : Cu  Cu2+ +2e  Half equation at cathode : Cu2+ + 2e  Cu

6.5 Voltaic Cell

1. Simple voltaic cell

 Uses two metal plates being immersed in an electrolyte (must contain one of the metal ions).

 Two different metals used must have different positions in the electrochemical series.

 Voltage can be measured by using voltmeter.

 The further the distance between those two metals in electrochemical series, the higher the voltage produced.

 Higher position of metal will donate electrons more easily to form positive ion and become a negative terminal (anode).

 Lower position of metal will accept electrons from the electrolyte to form metal and become a positive terminal (cathode).

 This results in the thinning and thickening of the plates.  Unstable and will decrease rapidly.

41 2. Daniell cell

 Produces more stable cell voltage.

 Cell built with two pieces of different metal immersed in a salt solution of their respective metals.

 Porous pot: to complete the circuit by allowing the transition of ions and separate both solutions.

 Porous pot can be replaced by salt bridge.

 Salt bridge: consists of filter paper soaked with a concentrated salt solution such as sodium chloride, potassium chloride, potassium nitrate, ammonium chloride and dilute sulphuric acid.

 Weaknesses:

 Electrolyte can spill out easily.  Difficult to carry around.

 Voltage produced decreases quickly due to the polarity of the cell(formation of gas bubbles around the electrodes)

42 3. Examples of voltaic cells:

43 4. Advantages and disadvantages of voltaic cells:

44

6.6 Construction of Electrochemical Series through Cell Potential

Difference

1. Procedure:

 30cm3 of 1 moldm-3 copper(II) sulphate solution is added into a beaker.  A piece of magnesium tape and copper metal are cleansed with sand

paper and immersed into copper(II) sulphate solution.

 Both pieces of metals are connected to a voltmeter using wires as shown in the diagram.

 The voltmeter reading is recorded. The positive and negative terminals are determined.

 The procedure is repeated by using zinc, iron, lead, aluminium and copper metal.

2. More electropositive metal : negative terminal

3. The further apart two metals are in the ECS, the higher the voltage of the cell.

6.7 Construction of Electrochemical Series through Displacement

Reaction

1. Metal which is more electropositive (placed higher) in the ECS will displace other metals less electropositive (below it) from its salt solutions.

2. Summary: Solution Metal Copper(II) salts (Cu2+) Lead(II) salts (Pb2+) Iron(II) salts (Fe2+) Zinc salts (Zn2+) Magnesium salts (Mg2+) Copper ~ No No No No Lead Yes ~ No No No

Iron Yes Yes ~ No No

Zinc Yes Yes Yes ~ No

45

Chapter 7: Acids and Bases

7.1 Acids and Bases

 Chemical substance that dissociate in water to produce hydrogen ions, H+ or hydroxonium ions, H3O+.

 Depicted as proton donors (H+).

 Strength of acid depends on the degree of dissociation/ionization.  3 types of acids:

 Monoprotic acid (HCl, HNO3)  Diprotic acid (H2SO4)

 Triprotic acid (H3PO4)  Physical properties:

 Sour in taste

 pH value: less than 7

 Turns blue litmus paper red.

 Conducts electricity (has free-moving ions).  Chemical properties:

 Acid + metal  salt + hydrogen gas

 Hydrogen gas can be tested by using a glowing splinter.

 Less reactive metals (Pb and Cu) are not suitable for the reaction.  Acid + carbonate salt  salt + water + CO2 gas

 CO2 gas turns lime water chalky/milky/cloudy.  Acid + alkali (base)  salt + water

 Neutralisation reaction.

Non-organic/mineral acid (strong acid) Organic acid (weak acid) a. Sulphuric acid, H2SO4

b. Hydrochloric acid, HCl c. Nitric acid, HNO3

d. Carbonic acid, H2CO3 e. Phosphoric acid, H3PO4 f. Sulphurous acid, H2So3

a. Methanoic acid, HCOOH b. Ethanoic acid, CH3COOH c. Lactic acid (sour milk) d. Citric acid (citrus fruit) e. Ascorbic acid (vit. C) f. Ethanediodic acid, H2C2O4 g. Formic acid (insect bites)

46

Strong acid Weak acid

 Dissociate completely into hydrogen ions in water.

 Degree of dissociation is 100%.  Produces higher concentration of

hydrogen ions and lower pH value.

 Eg:

 Hydrochloric acid  Sulphuric acid  Nitric acid

 Dissociate partially into hydrogen ions in water.

 Degree of dissociation is <100%.  Produces lower concentration of

hydrogen ions and higher pH value.

 Eg:

 Ethanoic acid  Methanoic acid  Citric acid

2. Water and acidic properties

 Acid shows its acidic properties in the presence of water.

 Acid will ionise to form hydrogen ions that are responsible for the acidic properties.

 Acid cannot be ionized in an organic solvent.

Acid Condition Observation Inference Hydrogen

chloride

Dissolved in methylbenzene

(non-acid)

No changes in the colour of litmus paper. Bulb does not light up.

Does not show acidic properties. Does not conduct electricity as there are no free moving ions.

Aqueous solution (acid)

Blue litmus paper turns red. Bulb lights up.

Show acidic properties.

Conduct electricity as there are free moving ions.

3. Bases

 Chemical substances that can neutralise an acid to produce salt and water (Neutralisation process).

 Most are not soluble in water. Soluble bases are called alkali.  Eg:

 zinc oxide, ZnO

 copper(II) oxide, CuO

47 4. Alkali

 Chemical substance that dissociate in water to produce hydroxide ion, OH-).

 Have alkaline properties as the formation of freely moving hydroxide ions in water.

 Eg:

 Sodium hydroxide, NaOH  Aqueous ammonia, NH3  Potassium hydroxide, KOH  Calcium hydroxide, Ca(OH)2  Physical properties:

 Feel soapy when in touch  Bitter in taste

 Turns red litmus paper blue  Has a pH >7

 Conducts electricity  Chemical properties:

 Acid + Alkali  Salt + Water (neutralization)

 Alkali + ammonium salt  salt + water + ammonia gas 5. Water and alkaline properties

 Alkaline properties only can be shown in the presence of H2O (presence of free-moving ions).

 Ionic compound – NaOH, KOH, Ca(OH)2

 Cannot show their properties in organic solvent.  Ionisation of alkali produces hydroxide ions in water.  Covalent compound – NH3

 Can dissolve in both water and organic solvent (trichloromethane).  Only show its properties in water.

 Conduct electricity only in water.

48

7.2 The Strength of Acids and Alkalis

Strong alkali Weak acid  Dissociate completely into

hydroxide ions in water.

 Degree of dissociation is 100%.  Produces higher concentration of

hydrogen ions and higher pH value (pH 14).

 Eg:

 Sodium hydroxide, NaOH  Potassium hydroxide, KOH

 Dissociate partially into hydrogen ions in water.

 Degree of dissociation is <100%.  Produces lower concentration of

hydroxide ions and low pH value.  Eg:

 Magnesium hydroxide, Mg(OH)2

 Aqueous ammonia, NH3  Calcium hydroxide, Ca(OH)2

1. pH concept

 To measure acidity and alkalinity of a solution.  From pH 0  14

 pH 7 : the concentration of H+ ions = the concentration of OH- ions

2. Acid – base indicator

Indicator Acid Alkali Neutral Litmus paper /

litmus solution

Red Blue Purple

Phenolphthalein Colourless Pink Colourless Methyl orange Red Yellow Orange

Universal indicator

Red/orange/yellow Green Blue/purple/violet Bromothymol

blue

Yellow Blue Green

49 3. Concentration / molarity



/



/

4. Standard solution: solution which has a known concentration.  Prepared by using volumetric flask.

 Dilution method:

7.4 Neutralisation

Titration method

 Reaction between an acid and a base to produce salt and water.  Acid + Base  Salt + Water

 H+ ions from acid will react with OH- ions from the alkali to produce water molecules.

H+ + OH-  H2O (ionic equation)  Neutral solution produced pH 7.

 Titration: method used to determine the molarity of a solution by using another solution with a known molarity.

50  End point: all the ions dissociated from acid and alkalis have reacted

completely to form water molecules. (neutral, pH 7)

 The water molecules dissociated into ions and thus do not conduct electricity.

 Formula:

=

(ratio of acid and alkali)

51

CHAPTER 8: SALTS

8.1 Salts

 Ionic compound that is formed when H+ ions in an acid is replaced by a metal ion or ammonium ion.

 Neutral [ pH 7 – phenolphthalein (colourless)]  Neutral in term of electrical charges.

 Can be produced through neutralisation process.  Examples:

Acids

Hydrochloric acid HCl X chloride Nitric acid HNO3 X nitrate Sulphuric acid H2SO4 X sulphate

Carbonic acid HCO3 X carbonate Phosphoric acid H2PO4 X phosphate

Ethanoic acid CH3COOH X ethanoate 2. Solubility of salts

Type of salts Solubility in water Sodium, potassium & ammonium

salts

All are soluble except oxide, hydroxide and carbonate

Nitrate, ethanoate salts All are soluble

Chloride salts All are soluble except PbCl2, AgCl,

HgCl2

Sulphate salts All are soluble except PbSO4, BaSO4,

CaSO4

Carbonate, oxide & hydroxide salts All are insoluble except sodium,

potassium & ammonium

Lead(II) salts All are insoluble except Pb(NO3)2 and Pb(CH3COO)2

***Lead hallides such as lead(II) chloride, lead(II) bromide and lead(II) iodide are insoluble in cold water but soluble in hot water.

***Lead(II) nitrate is soluble in both cold and hot water.

3. Preparation of soluble salts except soluble salts of sodium, ammonium and potassium

 Acid + alkali  salt + water  Acid + metal  salt + hydrogen  Acid + base  salt + water

52 4. Preparation of soluble salts of sodium, ammonium and potassium

i. Titration of an acid and alkali (Neutralisation) ii. Crystallization (Heating)

iii. Recrystallization (Filtration) 5. Crystals

 Formed when a saturated salt solution is cooled down.  Physical characteristics:

 Fixed geometrical shape

 Flat surfaces, straight edges and sharp corners  Fixed angle between two adjacent surfaces  Hard and brittle

 Colour of crystal depends on the ions in the crystals. Salt / metal oxide Colour

Solid Aqueous solution

Copper(II) salts

Copper(II) carbonate

Copper(II) sulphate, copper(II) nitrate, copper(II) chloride Copper(II) oxide Green Blue Black Insoluble Blue Insoluble Iron(II) salts

Iron(II) sulphate, Iron(II) nitrate, Iron(II) chloride

Light green Light green

53 Iron(III) sulphate, Iron(III)

nitrate, Iron(III) chloride

Brown Brown / Green – brownish Zinc oxide Yellow (hot)

White (cold)

Insoluble Lead(II) oxide Brown (hot)

Yellow (cold)

Insoluble Magnesium oxide, aluminium

oxide

White Insoluble Potassium oxide, sodium oxide,

calcium oxide

White Colourless 6.

54 1. Test for gases:

Gases Colour Smell Effect on damp litmus paper

Confirmation test

Oxygen, O2 Colourless - - Light up glowing splinter Hydrogen, H2 Colourless - - Lighted splinter

is placed near the mouth of the test tube. A “pop” sound is produced.

Carbon dioxide, CO2

Colourless - Damp blue litmus paper turns red

Bubbled through lime water. It will turn milky. Ammonia,

NH3

Colourless Pungent Damp red litmus paper turns blue

Forms dense white fumes with hydrogen chloride gas. Chlorine, Cl2 Greenish – yellow

Pungent Damp blue litmus paper turns red, then decolourises / bleaches it. Tested by using litmus paper. Hydrogen chloride, HCl

Colourless Pungent Damp blue litmus paper turns red

Forms dense white fumes with ammonia. Sulphur

dioxide, SO2

Colourless Pungent Damp blue litmus paper turns red

Decolourise purple colour of potassium manganate(VII) solution / changes orange potassium dichromate(VI) solution Nitrogen dioxide, NO2

Brown Pungent Damp blue litmus paper turns red

Tested by using litmus paper.

2. Production of gases:

55 Oxygen, O2 Heating a chlorate(V) or nitrate salt

Hydrogen, H2 Acid-metal reaction

Carbon dioxide, CO2 Heating a metal carbonate or acid-carbonate reaction Ammonia, NH3 Heating a mixture of ammonium salt and alkali

Chlorine, Cl2 Heating a mixture of manganese(IV) oxide and concentrated hydrochloric acid

Hydrogen chloride, HCl Heating a common salt and concentrated sulphuric acid

Sulphur dioxide, SO2 An acid-sulphite reaction Nitrogen dioxide, NO2 Heating a nitrate salt 3. Action of heat on salts

56  Nitrate salts

 Sulphate salts

i. Group 1 and 2 sulphate salts do not decompose when heated. ii. The sulphates of heavy metals decompose into metal oxides and

sulphur trioxide when heated except iron(II) sulphate which release sulphur dioxide gas.

iii. Ammonium sulphate sublimates at first and decompose into ammonia and hydrogen sulphate when further heating.

57  Chloride salts

i. All are stable to heat except ammonium chloride.

ii. Ammonia gas emerges first, then followed by hydrogen chloride. 4. Test for anions

58 5. Confirmatory tests

59 6. Tests for cations

60

CHAPTER

9:

MANUFACTURED

SUBSTANCES

IN

INDUSTRY

9.1 Sulphuric Acid

 To manufacture synthetic fibres (boat, wall)  To manufacture paint

 To manufacture metal oxide

 As an electrolyte (lead-acid accumulator) 2. Contact process

Sulphur, S _{dioxide, SO}Sulphur

2 Sulphur trioxide, SO₃ Oleum, H₂S₂O₇ Sulphuric acid, H₂SO₄ Step I Step II Step III

Catalyst: Vanadium(V) oxide Temperature: 450°C - 550°C Pressure: one atmosphere (atm)

61  Step I

i. Production of sulphur dioxide ii. S + O2  SO2

 Step II

i. Production of sulphur trioxide ii. 2SO2 + O2 2SO3

iii. High % of SO2 is converted into SO3.  Step III

i. SO3 + H2SO4(concentrated)  H2S2O7 (oleum) ii. H2S2O7 + H2O  2 H2SO4

9.2 Ammonia

 To manufacture fertilizers

 As a cooling agent in refrigerator

 To produce nitric acid (Ostwald process)  To make explosives

 To prevent coagulation of latex  To produce ammonium chloride 2. Haber process + Nitrogen, N2 Hydrogen, H2 Ammonia, NH3 Excess N2 and H2

62

3. Nitrogen gas is obtained from fractional distillation of liquid air. 4. Source of hydrogen gas:

 Reaction between steam and heated coke (carbon)  Reaction between steam and natural gas (methane) 5. Equation: N2 + 3H2  2NH3

6. Ratio of nitrogen gas to hydrogen gas  1 : 3

9.3 Alloy

1. Advantages of alloying:

 Increase the hardness/strength of metal  Prevent corrosion/rusting

 Improve appearance of metal 2. Physical properties:

 Ductile – can slide over when external force is applied. Catalyst: Iron powder

Promoter: Aluminium oxide Temperature: 450 – 550 Pressure: 200 – 500 atm

63  Malleable – slide into new positions in the empty spaces of alloy.

 High boiling and melting points  High density

 Good conductor of electricity

Alloy Composition Properties Uses Carbon steel 99% iron

1% carbon

Hard For construction, bridges, vehicles, tools, heavy machinery Stainless steel 74% iron

18% chromium 8% nickel

Rust resistant For crockery, kitchenware and machine parts

Bronze 90% copper 10% tin

Hard & shiny For kitchenware, ship propellers, decorative ornaments and art crafts.

Brass 70% copper 30% zinc

Hard & shiny For musical instrument, electrical connecter, decorative ornaments

Magnalium 70% aluminium 30% magnesium

Light & hard Duralumin 95% aluminium

1% magnesium 4% copper

Light & hard

Pewter 97% tin

3% lead and antimony

Hard & shiny For mugs, candlesticks,

decorative ornaments and souvenirs.

Solder 50% tin 50% lead

Hard, shiny and low melting point

For soldering electrical wires and metal pipes Cupro-nickel Copper, nickel Hard, shiny and

corrosion resistant

For coins

9.4 Synthetic Polymers

1. Polymers: large long-chain molecules formed by joining together many identical repeating sub-units called monomer.

2. Polymerisation: chemical process by which the monomers are joined together into chain-like molecule called polymer.

64 4. Advantages of synthetic polymers

 Strong and light  Cheap

 Able to resist corrosion  Inert to chemical reactions

 Easily moulded or shaped and be coloured  Can be made to have special properties

Name(s) Monomer Properties Uses Polyethylene

low density (LDPE)

ethylene CH2=CH2

soft, waxy solid film wrap, plastic bags Polyethylene high density (HDPE) ethylene CH2=CH2 rigid, translucent solid electrical insulation bottles, toys Polypropylene (PP) different grades propylene CH2=CHCH3

atactic: soft, elastic solid isotactic: hard, strong solid similar to LDPE carpet, upholstery Poly(vinyl chloride) (PVC) vinyl chloride CH2=CHCl

strong rigid solid pipes, siding, flooring Poly(vinylidene chloride) (Saran A) vinylidene chloride CH2=CCl2 dense, high-melting solid seat covers, films Polystyrene (PS) styrene CH2=CHC6H5

hard, rigid, clear solid soluble in organic solvents toys, cabinets packaging (foamed) Polyacrylonitrile (PAN, Orlon, Acrilan) acrylonitrile CH2=CHCN high-melting solid soluble in organic solvents rugs, blankets clothing Polytetrafluoroeth ylene (PTFE, Teflon) tetrafluoroethylen e CF2=CF2 resistant, smooth solid non-stick surfaces electrical insulation Poly(methyl methacrylate) (PMMA, Lucite, Plexiglas) methyl methacrylate CH2=C(CH3)CO 2CH3 hard, transparent solid lighting covers, signs skylights Poly(vinyl acetate) (PVAc) vinyl acetate CH2=CHOCOC H3

soft, sticky solid latex paints, adhesives

65 cis-Polyisoprene natural rubber isoprene CH2=CH-C(CH3)=CH2

soft, sticky solid requires vulcanization for practical use Polychloroprene (c is + trans) (Neoprene) chloroprene CH2=CH-CCl=CH2

tough, rubbery solid synthetic rubber oil resistant

9.5 Glass and ceramics

1. Properties of both glass and ceramics:

 Main component: silica or silicon dioxide, SiO2  Hard but brittle

 Inert towards chemicals

 Insulator of heat and electricity

 withstand compression but not stretching  Can be easily cleaned

 Low cost of production

2. Differences between glass and ceramics:  Glass – transparent , ceramic – opaque

 Glass – cannot withstand high temperature, ceramics – can withstand high temperature

66 4. Special purpose glass and ceramics

I. Photochromic glass

II. Conducting glass – produced by adding tin(IV) oxide (conduct electricity) used to make LCD.

III. Super conductor – ceramics used to make light magnets, electrical generators & electric motors

9.6 Composite Materials

1. New material produced from a complex mixture of two or more materials with different physical properties.

67

~ THE END ~

~ febianhenry_96 ~