Chemistry Form 4 A+ Notes

August 27, 2017 | Author: Febian Henry | Category: Atomic Nucleus, Hydroxide, Acid, Chlorine, Ion
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SEPT 2013

CHAPTER 2: STRUCTURE OF THE ATOM 2.1 MATTER 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 Atoms

 


  


  

Matter Elements

Description Smallest particles of an element that retain the chemical properties of the element. Examples : Sodium atom (Na) Zinc atom (Zn) Helium atom (He) 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) Charged particles – positive or negative Positive charged ion (Cation) – Zinc ion (Zn2+) Negative charged ion (Anion) – Chloride ion(Cl-)

   


  

Descriptions 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) 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+, Cl-) 2

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. SOLID


Boiling point


Melting point

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 3

 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 Arrangement of particles




compact, orderly manner

Loosely packed, disorderly manner

very far apart, random motion

Vibrate, rotate in a fixed position

Move freely

Move freely and randomly

Particles Kinetic energy Shape

Very low




Volume Rate of diffusion

Fixed Low

Not fixed (follow the shape of container) Fixed Average

Not fixed (follow the shape of container) Not fixed High

Attractive forces between particles

Very strong


Very weak

Particles motion


3. Experiment (PeKa) a. Heating curve of naphthalene/acetamide  Diagram:

 Graph:

AB: Solid

DE: Liquid + Gas

BC: Solid + Liquid

EF: Gas

CD: Liquid


 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:


 Graph:

PQ: Gas

ST: Solid + Liquid

QR: Liquid + Gas

TU: Solid

RS: Liquid  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)


 Super cooling



Condition in which the temperature of a cooling liquid drops below the normal freezing point. Occurs when conical flask is not used in the experiment.

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


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. - The orbital has various radius form the nucleus. 9

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 Proton Neutron Electron

Symbol p n e˗

Relative atomic mass (RAM) 1 1

Charge + neutral ˗

3. Electron Configuration  Maximum number for each shell:  First shell : 2 electrons  Second shell : 8 electrons  Third shell : 8 electrons  Forth shell : 2 electrons  Valence electron = electrons found in the outermost shell of an atom.


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 1. According to the 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:


 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.


2.4 ISOTOPES 1. Atoms of same element with the same number of protons but different number of neutrons. Uses Isotopes Carbon-14

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


 biological tracer, for example, in studies of photosynthesis


 Detect location of leaks in water pipes,  studies of body electrolytes


 location of leaks in water pipes


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


 lung ventilation studies


 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




   

Medical tracer treat the thyroid gland & used in the diagnosis of adrenal medulla for imaging suspected neural crest and other endocrine tumours  used in imaging to monitor thyroid function  detect adrenal dysfunction  Enriched as a fuel for most nuclear reactors 13


 Domestic smoke alarms


 Treatment of excess red blood cells


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: MASS OF SUBSTANCE, g

÷ M.M.

× M.M.


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

3.2 The Mole and the Volume of Gas 1.   2. 

Avogadro’s Constant, NA Number of particles in one mole of substance. 6.02 × 1023 Standard Temperature and Pressure (S. T. P.) Temperature = 0°C 15

  3.   

Pressure = 1 atmosphere / atm Molar volume of 1 mole of gas = 22.4dm3 or 22400cm3 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)

Crucible with lid

Metal tape


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.



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 1 2 3 4 valence electrons Group 1 2 13 14 3. Historical Development of the Periodic Table i. Antoine Lavoisier










 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:  Colourless gaseous state at room temperature. 19

 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.  Less dense than water.


 When going down the group, Melting and boiling point Metallic bond Decreases Forces of attraction Atomic size Increases 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 Alkali metal Cold water Lithium Sodium Potassium

2Li + 2H2O  2LiOH + H2 Moves slowly with “hiss” sound. 2Na + 2H2O  2NaOH + H2 Moves quickly and randomly with loud “hiss” sound. 2K + 2H2O  2KOH + H2 Burns with reddish-purple light, jumps, “hiss” and “pop” sound.


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)

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) Lithium



Lithium Sodium Potassium

2Li + Cl2  2LiCl / 2Li + Br2  2LiBr Burns slowly with reddish flame. A white solid is obtained. 2Na + Cl2  2NaCl / 2Na + Br2  2NaBr Burns brightly with a yellowish flame. A white solid is obtained. 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. 23

 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  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 (gas)

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: Chlorine


/ Lime soda

 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. 25

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 Increases  Valence electrons  Non-metallic properties  Nuclei attraction on valence electrons  Atomic size  Electropositivity Decreases  Metallic properties (Metallicity) Na2O MgO Al2O3 SiO2 P4O10 SO2 Cl2O7 Element Acidic oxides Characteristics Basic oxides Amphoteric (Alkali) 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


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 3+ Chromium ion, Cr Green 2+ Iron(II) ion, Fe Green 3+ Iron(III) ion, Fe Brown 2+ Copper(II) ion, Cu Blue 2Chromate ion, CrO 4 Yellow 2+ Manganese ion, Mn Pink 2+ Cobalt ion, Co Pink 2+ Nickel ion, Ni Green 2Manganate ion, MnO 4 Purple 2Dichromate ion, Cr2O 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) – [Cu(NH3)4]2+


 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 Emerald Ruby Sapphire Amethyst

Colour Green Red Blue Purple

Transition elements Nickel, Iron Chromium Iron, Titanium Iron, Manganese


CHAPTER 5: CHEMICAL BONDS 5.1 Formation of Chemical Bonds 1. Ionic bond  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 Metal M Non-metal X + Group 1, M Group 15, M3Group 1, M+ Group 16, M2Group 1, M+ Group 17, MGroup 2, M2+ Group 15, M3Group 2, M2+ Group 16, M2Group 2, M2+ Group 17, MGroup 13, M3+ Group 15, M3Group 13, M3+ Group 16, M2Group 13, M3+ Group 17, M Sodium chloride, NaCl

Formula of ionic compound M3X M2X MX M3X2 MX MX2 MX M2X3 MX3


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)

Nitrogen molecule, N2 (Triple)


Non-metal elements which combined Element P Element Q 4+ Group 14, P Group 17, QGroup 14, P4+ Group 16, Q2Group 15, P3+ Group 17, QGroup 16, P2+ Group 17, Q-

Molecular Formula PQ4 P2Q4 / PQ2 PQ3 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  Do not contain free-moving ions that carry charges


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 NaCl, MgO, ZnCl2, CuO High – has strong electrostatic forces of attraction (A lot of heat energy is needed) Soluble in water, insoluble in organic solvent. Conduct electricity in both molten and aqueous state – contain free– moving charged ions. Non – volatile

Properties Examples Melting and boiling points


Electrical conductivity


Covalent compound O2, CO2, H2O, N2, Cl2 Low – has weak intermolecular forces of attraction (Little heat is needed) Soluble in organic solvent, insoluble in water. Do not conduct electricity.



Ionic compound

Covalent compound Simple Giant


M/P & B/P Solubility

High Soluble in water, insoluble in organic solvents

Low High Insoluble in Insoluble organic both solvent, soluble in water Electrical Conduct in Does not Does not conductivity molten state or aqueous solution Volatility No Yes No

Metal Cu, Zn, Na, Ca, Pt, Ni, Mg High in Insoluble in both

Conduct in solid or liquid



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. Electrical energy  chemical energy


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 35

12.Electrolysis of molten compounds Metal Sodium Lead Nickel Copper Gas Bromine Iodine Chlorine Oxygen Hydrogen

Observation Shiny grey solid is formed. Shiny grey solid is deposited. Shiny grey solid is formed. Brown deposit is formed. Observation Brown gas is produced. (pungent smell) Purple gas is produced. Yellowish-green gas is produced. Colourless gas bubbles are formed. (effervescence) Colourless gas bubbles are formed. (effervescence)

Test for Oxygen gas, O2

Test for Hydrogen gas, H2

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

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

6.2 Electrolysis of Aqueous Solution 1. 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


4. Position of ions in the electrochemical series (ECS) Cations K Na Ca Mg Al Zn Fe Sn Pb H Cu Ag

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Anions -


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 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 1. Electroplating of metals  Objectives: to prevent corrosion / to improve appearance.  Plating metals: gold (Au), Platinum (Pt), Chromium (Cr), copper (Cu), Silver (Ag), & Nickel (Ni). 37

 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+ + 3O2 Half equation at cathode : Al3+ +3e  Al 38

 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  Electrolyte used must contain the metal ions.


 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.


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)


3. Examples of voltaic cells:


4. Advantages and disadvantages of voltaic cells:

5. Comparison between electrolytic cell and voltaic cell:


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 Copper(II) salts Metal (Cu2+) Copper ~ Lead Yes Iron Yes Zinc Yes Magnesium Yes

Lead(II) salts (Pb2+) No ~ Yes Yes Yes

Iron(II) salts (Fe2+) No No ~ Yes Yes

Zinc salts (Zn2+) No No No ~ yes

Magnesium salts (Mg2+) No No No No ~


Chapter 7: Acids and Bases 7.1 Acids and Bases 1. Acid  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) 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. b. c. d. e. f. g.

Organic acid (weak acid) Methanoic acid, HCOOH Ethanoic acid, CH3COOH Lactic acid (sour milk) Citric acid (citrus fruit) Ascorbic acid (vit. C) Ethanediodic acid, H2C2O4 Formic acid (insect bites)


   

Strong 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

   

Weak acid Dissociate partially into hydrogen ions in water. Degree of dissociation is 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.  There is no mobile ion in organic solvents.


7.2 The Strength of Acids and Alkalis    

Strong alkali 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

   

Weak acid Dissociate partially into hydrogen ions in water. Degree of dissociation is
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