Olevels Chemistry Notes - Combined Chemistry

2014 Chemistry Notes For Secondary School Combined Chemistry

O-Levels By Marcus Ng

©2014 Marcus Ng Chemistry Notes for Secondary School O-Levels Combined Chemistry

Chapter 1 Experimental Chemistry 1.1 Measurements Physical Quantity

SI Unit

Time

Second (s)

Temperature

Kelvin (K)

Mass

Kilogram (kg)

Length

Meter (m)

Volume

Cubic Meter (m3)

Apparatus

Accuracy

Digital Stopwatch Analog Stopwatch Mercury Thermometer Alcohol Thermometer Electronic Balance Beam Balance Ruler

0.01 s 0.1s 0.01 K 0.01 K

Vernier Calipers Micrometer Beaker Measuring Cylinder Pipette* Burette

0.01 cm (0.1mm) 0.001 cm (0.01mm)

0.1 cm (1mm)

1 cm3 (1 ml) 0.1 cm3 (0.1 ml) 0.1 cm3 (0.1 ml)

How to read a Vernier Caliper reading & A Micrometer reading

Important Points to remember: 1. When measuring Time:

Digital Stopwatch is more accurate than Analog Stopwatch

2. When measuring Length:

Micrometer is more accurate than Vernier Calipers, which are both more accurate than a Ruler

3. When measuring Volumes:

Pipettes are only used for specific volumes (10 cm3, 25 cm3 or 50 cm3)

4. When measuring Volumes:

Burettes are the most accurate, followed by a Measuring Cylinder and lastly a Beaker

5. When measuring Temperature: Maximum upper limit for Alcohol Thermometer is 351.15K (78 0C) 6. When measuring Temperature: Mercury Thermometers are more expensive and toxic then Alcohol Thermometers.


1.2 Separation Techniques Technique

Purpose

Filtration

Used to separate aninsoluble solid from a liquid

Crystallization

Used to separate asoluble solid from a liquid

Distillation

Used to separate aliquid from a soluble solid

Fractional Distillation

Used to separate aliquid from a mixture of Miscible Liquids*

Separating Funnel

Used to separate aliquid from a mixture of immiscible liquids*

Sublimation

Used to separate asublimable solid from a mixture of solids

Magnetic Attraction

Used to separate aSolid that can be magnetized

*Miscible Liquids refer to liquids that can be completely mixed

Filtration

*Note: When given a soluble salt andinsoluble salt, add water to dissolve the soluble salt. When given a soluble salt and organic compound, add water to dissolve the soluble salt.

Crystallization


Distillation

Fractional Distillation


Separating Funnel

Sublimation


1.3 Collection of Gases

Properties of Gases Gas Ammonia Argon Carbon Dioxide Carbon Monoxide Chlorine Helium Hydrogen Hydrogen Bromide Hydrogen Chloride Methane Oxygen Neon Nitrogen Nitrogen Dioxide Sulphur Dioxide

Solubility Soluble in Water Insoluble in Water Soluble in Water Insoluble in Water Soluble in Water Insoluble in Water Insoluble in Water Soluble in Water Soluble in Water Insoluble in Water Insoluble in Water Insoluble in Water Insoluble in Water Soluble in Water Soluble in Water

Density Less dense than air More dense than air More dense than air Less dense than air More dense than air Less dense than air Less dense than air More dense than air More dense than air Less dense than air Similar density to air Less dense than air Similar density to air More dense than air More dense than air


1.4 Purity of a Substance Important Points/Concepts to remember: 1. A Pure Substance melts and boils at a fixed and constant temperature 2. Impurities decreases the melting point of a substance 3. Impurities increases the boiling point of a substance Methods to check the purity of a substance 1. Melting Point Determination

2. Boiling Point Determination

3. Chromatography


Chromatography

1. Chromatography can be used to determine the purity of a substance 2. Chromatography can be used to identify the substance 3. Chromatography can be used to separate components of asubstance with different solubilities in the same solvent and identify them. 4. There are 2 types of Paper Chromatography: Ascending and Descending 5. There can be 3 types of results (chromatograms that canbe developed) a. Only one spot is seen - A Pure Substance (one solute in sample) b. More than one spot is seen - A Mixture (more than one solute in sample) c. No spots are seen - No soluble* solute in mixture (*in the solventused) Note: Some compounds are colourless and thus a locating agent need to be used. An example of a locating agent is Ninhydrin, used for locating amino acids. 

solvent line

start line


1.5 Tests for Cations Test Cation

Symbol

Add dilute sodium hydroxide solution to a solution of the substance.

Add dilute ammonia solution to a solution of the substance.

White precipitate

White precipitate

that dissolves in excess sodium hydroxide

that is insoluble in excess ammonia.

3+

Aluminum

Al

Ammonium

NH4

+

Ammonia gas is produced White precipitate

Calcium

Ca

2+

Cu

2+

that dissolves in excess sodium hydroxide Pale green precipitate

Iron(II)

Fe

2+

Iron(III)

Fe

3+

Lead(II)

Pb

No Chemical Reaction

No Chemical Reaction Red Flames are produced

that is insoluble in excess sodium hydroxide. Pale Blue precipitate

Copper (II)

Flame Test

Pale blue precipitate changing to deep blue solution in excess ammonia. Pale green precipitate

that is insoluble in excess sodium hydroxide.


Red-brown precipitate

Red-brown precipitate


that is insoluble in excess ammonia..

White precipitate

White precipitate



2+

White precipitate Magnesium

Mg

Zinc

Zn

2+

Green Flames are produced

Blue Flames are produced

White precipitate



White precipitate

White precipitate


that dissolves in excess ammonia

2+


1.6 Tests for Anions Anion

Symbol

Test

-

Chloride

Cl

Bromide

Br

Iodide

Add aqueous silver nitrate solution to a solution of substance

-

Add acidified lead (II) nitrate solution to a solution of substance

Add dilute hydrochloric acid to the substance.

2-

Carbonate

CO3

Nitrate

NO3

Sulphate

SO4

Add dilute sodium hydroxide solution, followed by Aluminum powder and warm

-

Acidify the solution of the substance (Either HCl or HNO 3)

2-

White precipitate that is soluble in ammonia solution. Cream precipitate, that is slightly soluble in ammonia solution.

Or

I-

Results

Add solution of barium cations (BaCl or BaNO3) to the solution.

Pale yellow precipitate, that is insoluble in ammonia solution.

Carbon dioxide gas is produced

Ammonia gas is produced

White precipitate, does not dissolve in excess dilute acid.

1.7 Tests for Gases Gas

Symbol

Properties

Litmus Test

Splint Test

Hydrogen

H2

Colourless & Odourless

Extinguish a lighted splint with a pop sound

Oxygen

O2


Relights a glowing splint

Carbon Dioxide

CO2


Cl2

Greenish-yellow with a Choking Smell

Chlorine

Turns moist Blue litmus Red

Limewater test

Forms a White precipitate

Turns moist Blue litmus Red and eventually bleaches White

Ammonia

NH3

Colourless with a Pungent Smell

Turns moist Red litmus Blue

Sulphur Dioxide

SO2

Colourless with a Choking Smell

Turns moist Blue litmus Red


Chapter 2 Kinetic Particle Theory 2.1 Three States of Matter Properties

Solid

Liquid

Gas

Volume

Fixed Volume

Fixed Volume

No Fixed Volume

Shape

Fixed Shape

No Fixed Shape

No Fixed Shape

Compressibility

Incompressible

Negligibly compressible

Very compressible

Packing

Particles are very closely packed

Particles are closely packed

Particles are very far apart

Forces of Attraction

Very strong forces of attraction between particles

Strong forces of attraction between particles

Very weak forces of attraction between particles

Motion

Particles vibrate about a fixed position

Particles can slide over each other

Particles are in random motion

Diagrammatic Representation


2.2 Melting & Freezing Important Explanation: Melting

During melting, the particles of a solid gainenergy and vibrate until they overcome the forces of attraction between the particles, movingfaster and further apart. At this point, there is no rise in temperature as all heat energy is used to separate the particles at this point. This is the melting point, at which the temperature remains constant until the whole solid has melted into a liquid. Important points to include: 1. Gaining/Losing Energy 2. Motion of particles 3. Forces of attraction 4. New motion of particles 5. No rise/drop in temperature as all heatenergy is used to separate/combine the particles 6. _____ Point where temperature remains constant Graphical Representation K / e r u t a r e p m e T

d

b

c

a Time/min

1. Between points a and b, the substance is in theSolid state 2. Between points b and c, the substance is in a mixture ofSolid & Liquid states 3. Between points c and d, the substance is in theLiquid state


2.3 Boiling & Condensation

During boiling, the particles of a liquid gainenergy, sliding over each other until they overcome the forces of attraction between the particles, movingfaster, randomly and very far apart. At this point, there is no rise in temperature as all heat energy is used to separate the particles at this point. This is the boiling point, at which the temperature remains constant until the whole liquid has boiled into a gas. Difference between Boiling & Evaporation Boiling

Evaporation

Occurs at boiling point Occurs throughout the liquid Bubbles observed Occurs quickly

Occurs at any temperature below boiling point Occurs only at the surface of the liquid No bubbles observed Occurs slowly

2.4 Sublimation Example of Substances that sublimes are Carbon Dioxide, Naphthalene (Mothballs) and Iodine


Chapter 3 Atomic Structure and Chemical Bonding 3.1 Atomic Structure (& Symbols to represent Atomic Structure)

Atoms are the basic building blocks of all matter.

Particle

Relative Charge

Relative Mass

Location

Proton Neutron Electron

+1 0 -1

1 1 0.0005

Nucleus Nucleus Electronic Shells outside the Nucleus

1. Atomic Number (or Proton Number) → The number of Protons in an atom The number of Protons in an atom = The number of electrons 2. Mass number (or Nucleon Number) → The number of Protons + Neutrons 3. Atoms with same number of electrons andprotons but different number of neutrons 

= isotopes

 ()    ()   

 

Common Isotopes Element

Hydrogen

Carbon Chlorine

Protium Deuterium Tritium Carbon-12 Carbon-13 Carbon-14

Proton Number 1 1 1 6 6 6

Number of Electrons 1 1 1 6 6 6

Nucleon Number 1 2 3 12 13 14

Number of Neutrons 0 1 2 6 7 8

Chlorine-35 Chlorine-37

17 17

17 17

35 37

18 20

Isotope


3.2 Electron Arrangement

1. 2. 3. 4. 5.

Electrons in an atom are arranged in energy shells The arrangement of the electrons is call its electronic configuration The first shell can hold a maximum of2 electrons The second and third shell can hold a maximum of 8 electrons The outermost shell is called the valence shell. the electrons in this shell is called valence electrons. 6. The shells & the no. of electrons can be represented by a dot-and-cross diagram. Electronic Configuration & Dot-and-cross Diagram Examples Element

Electronic Configuration

Hydrogen

1

Carbon

2.4

Oxygen

2.6

Silicon

2.8.4

Dot-and-cross Diagram


Easy Reference Table (Please try to not rely on this. learn to derive these information from a periodic table) Symbol

Element

Proton Number

Electronic Configuration

                                   or       

Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine

1 2 3 4 5 6 7 8 9

1 2 2.1 2.2 2.3 2.4 2.5 2.6 2.7

Neon Sodium Magnesium Aluminum Silicon Phosphorus Sulphur Chlorine Argon Potassium Calcium

10 11 12 13 14 15 16 17 18 19 20

2.8 2.8.1 2.8.2 2.8.3 2.8.4 2.8.5 2.8.6 2.8.7 2.8.8 2.8.8.1 2.8.8.2

3.3 Formation of Ions

1. elements) Atoms are generally naturally unstable. (With the exception of Group 0/Group 8 2. This is due to the lack of a stable octet (or duplet) structure, with fully filled shells. 3. Atoms can form ions by gaining or losing valence electrons, in their attemptto obtain a stable octet (or duplet) structure, with fully filled shells. 4. Metals usually lose electrons, forming Positive Ions, also known as Cations. 5. Non-Metals usually gain electrons, forming Negative Ions, also known as Anions. Example Lithium can lose a valence electron to form Li +, a positive ion (cation) with a fully filled valence shell of 2 electron, with an electronic configuration of 2. Magnesium can lose two valence electrons to form Mg2+, a positive ion (cation) with a fully filled valence shell of 8 electron, with an electronic configuration of 2.8 Oxygen can gain two valence electrons to form O 2-, a negative ion (anion) with a fully filled valence shell of 8 electron, with an electronic configuration of 2.8 






3.4 Ionic Bonding

1. Ionic Bonds are formed when metallic atoms give their valence electrons to nonmetallic atoms. This enables both the metallic and the non-metallic atoms to achieve a stable octet (or duplet) structure, with fully filled valence shells. 2. These ions formed are oppositely charged, and attract each other through strong electrostatic forces of attraction, thus forming the ionic bond. Examples

Na + Cl → NaCl Sodium (Na) can lose a valence electron to Chlorine (Cl), forming NaCl, with the positively charged Na+ ion, and the negatively charged Cl- ion.

Mg + O

→

MgO

Magnesium (Mg) can lose two valence electrons to Oxygen (O), forming MgO, with the positively charged Mg2+ ion, and the negatively charged O2- ion.

Mg + 2 Cl

→

MgCl2

Magnesium (Mg) can lose two valence electrons, one to each Chlorine (Cl), forming MgCl2, with the positively charged Mg2+ ion, and 2 negatively charged Cl- ions.


3.5 Structure of Ionic Compounds

1. A solid ionic compound has a giant lattice structure with alternating positivelyand negatively charged ions. 2. The ions are held in fixed positions by strong electrostatic forces ofattraction.

3.6 Physical Properties of Ionic Compounds Physical Properties of Ionic Compounds

Explanation in terms of their structure and bonding

Ionic compounds are usually The ions are arranged in a highly regular fashion, crystalline solids at room with strong electrostatic forces of attraction (ionic temperature bonds) between the ions. Ionic compounds have melting and boiling points

high The electrostatic forces of attraction between the oppositely charged ions is very strong and extends over the entire crystalline structure. Large amounts of energy is required to separate the ions

Ionic compounds cannot conduct In the solid structure, the ions are held in fixed electricity when solid, but do so in positions. When molten or in aqueous solution, the molten or in aqueous form. ions are mobile, so a flow of charge is possible. Most ionic compounds are water Water molecules are polar, and are attracted to the soluble, but insoluble in organic charged ions in the ionic compound. This helps to solvents. pull the crystalline structure as the solid dissolves.


3.7 Covalent Bonding

1. Covalent bonding occurs when the electrons are shared, so as to achieve a stable octet/duplet structure. 2. Each pairs of shared electrons forms one covalent bond. 3. Covalent bonding occurs mainly between non-metals Examples

H+H

→

H2

Two hydrogen atoms can share an electron each to form a covalent H-H bond, giving both atoms a stable duplet structure.

O + O → O2 Two oxygen atoms can share two electrons each to formtwo covalent O-O bonds, giving both atoms a stable octet structure.

H+H+O

→

H2 O

Each hydrogen atom shares one electron with the oxygen atoms, forming 2 O-H covalent bonds.


3.8 Physical Properties of Simple Covalent Compounds Physical Properties of Simple Covalent Compounds

Explanation in terms of their structure and bonding

Simple Covalent Compounds The inter-molecular forces of attraction are very have low boiling and melting weak, hence very little energy is required to break points. the forces apart. Simple Covalent Compounds There are no mobile ions or electrons in simple cannot conduct electricity in any covalent compounds in any states. state. Simple Covalent Compounds are Simple Covalent Compounds have generally nonsoluble in organic solvents but polar molecules, and thus would be unable to not in water. dissolve in a solvent like water with strong hydrogen bonding, but would be soluble in a organic solvent like ethanol, petrol or trichloromethane.


3.9 Elements, Compounds and Mixtures.

1. Elements are made up of only one kindof atoms, and can be found directly in the periodic table. 2. Elements cannot be further broken down by physical or chemical means (at least for O-levels syllabus). 3. Compounds are made of two or more different kinds atoms chemically combined in a fixed proportion. E.g. Hydrochloric acid comprises of hydrogen atoms and chlorine atoms in a 1:1 fixed proportion only. 4. A mixture is made up of two or more elements and/or compounds physically combined. They can be physically separated by physical means and do not have a fixed proportion. E.g. Saltwater can be 50% salt 50% water, or 40% salt 60% water or 30% salt 70% water … 



Composition

Properties

Mixture No fixed composition/ proportion. The percentage of one element/compound to another in a mixture can vary.

Compounds Fixed composition/ proportion. The percentage of one element to another in a particular compound is always the same.

No set of properties of its own. It exhibits a combination of the properties of the constituent components.

It has its own set of properties.

Melting Point No fixed M.P. or B.P. & Boiling Point Preparation No chemical reaction has to occur. Separation Can be separated into its components by physical means

A fixed M.P. and B.P. A chemical reaction has to occur. Can only be separated into its components by chemical means


Chapter 4 Stoichiometry & Mole Concept 4.1 Chemical Formulae    







Number of Atoms/Ions are denoted by subscript. Charge of ions are denoted by superscript. Metals and/or positive cations are placed first in the chemical formula. Brackets are used for repeated clusters of atoms (like anions) o E.g. Magnesium Nitrate = Mg(NO 3)2 Mono- is a prefix for indicating only 1 of a particular atom/ cluster of atoms Carbon Mono xide = CO o Di- is a prefix indicating 2 of a particular atoms/cluster of atoms o Carbon Dioxide = CO2 Tri- is a prefix indicating 3 of a particular atoms/cluster of atoms Dinitrogen Trioxide = N2O3 Prefixes are only used for covalent molecules. For ionic compounds, the formulae have to be deduced from the valency of the components. For transition metals, the valency is indicated in brackets Iron (II) has a valency of 2, Iron (III) has a valency of 3. o o





Common Ions Name

Formulae

Ammonium Carbonate Chromate (VI) Dichromate (VI) Ethanoate Hydrogencarbonate Hydroxide

+

NH3 CO2 CrO4 2Phosphate Cr2O7 CH3CO2 HCO2 OH

Name

Formulae

Nitrate Nitrite Oxide

NO3 NO2 O

-

PO Sulphate / Sulfate Sulphite / Sulfite Sulphide / Sulfide

34 -

SO4 SO3 2S

4.2 Balancing Equations (With state symbols) 

A chemical equation is used to shows information in a chemical reaction. What chemicals are used (Reactants). o What chemicals are created (Products). o What states they are in: (s), (l), (g) or (aq) The reactants are always on the left, and the products on the right. Ensure the left side of the equation equals the right side. o

 

Example 1:

Fe(s) + 2 C5H6(g) → Fe(C5H5)2(s) + H2(g)   

There is 1 Iron atom on both sides. There are 10 Carbon atoms on both sides. There are 12 Hydrogen atoms on both sides.

Example 2:

CH3CH2OH + CH3CO2H ⇌ CH3CO2CH2CH3 + H2O    

There are 4 Carbon atoms on both sides. There are 10 Hydrogen atoms on both sides. There are 4 Oxygen atoms on both sides. *The catalyst HCl is not included in the equation as it appears in the same form on both sides.


-

4.3 Ionic Equations 

An ionic equation only shows the ions involved in the reaction An ion is involved in the reaction if its charge changes during the reaction. Ions uninvolved are called spectator ions o Note: Insoluble compounds should not be broken up into its component ions o Ensure the total charge on the left side of the equation equals the total charge on the right side. o



Example 1: Chemical Equation: Ionic Equation:  

CuCO3 (s) + 2 HCl(aq)  CuCl2(aq) + CO2(g) + H2O(l) + 2+ CuCO3 (s) + 2 H (aq)  Cu (aq) + CO2(g) + H2O(l)

Charge on the left side 2 x (+1) = +2 Charge on the right side +2 = Charge on the left side

Example 2: Chemical Equation: Ionic Equation:  

NaOH (aq) + HCl(aq)  NaCl(aq) + H2O(l) + OH (aq) + H (aq)  H2O(l)

Charge on the left side - 1 +1 = 0 Charge on the right side 0 = Charge on the left side

4.4 Relative Atomic Mass (Element) & Relative Molecular Mass The relative atomic mass (A r) of an element is the average mass of one atom of an element compared to  of the mass of a carbon-12 atom. 





Relative Atomic Mass (A r) may sometimes have the same values as the mass number, but they are conceptually DIFFERENT from each other. o Mass number refer to the number of protons and neutrons in an atom. They can differ betweens isotopes of the same elements. Atoms of different elements can have the same mass number. o Relative Atomic Mass (A r) refers to the AVERAGE mass of atoms of a particular element in accordance with isotopic composition. Relative Atomic Mass (A r) has no units.

The relative molecular mass (M r) of a substance is the average mass of one molecule of the substance  compared with of the mass of a carbon-12 atom. 

4.5 % by Mass of an Element in a Compound % by Mass of an Element in a Compound =

      ()      ()   

Example 1: % by Mass of Oxygen in Carbon Dioxide (CO 2) = =

     ()   ()         

x 100%

x 100%

= 72.73% (2 decimal points)


x 100%

4.6 Mole Concept 



A mole is the number of particles which contains the same number of atoms in a 12.0g sample of carbon-12. 23 This number is 6.02 x 10 . This number is also known as Avogadro's number. o Mole can be abbreviated as mol.

4.7 Molar Mass and Molar Volume 

The mass (in grams) of 1 mole of a substance, is called its molar mass. The molar mass of a substance is equal to its relative atomic mass or relative molecular mass. E.g. The relative atomic mass (A r) of Helium is 4.0. The molar mass of Helium is 4.0 g. The 23 mass of 1 mole of Helium atoms is 4.0 g. The mass of 6.02 x 10 Helium atoms is 4.0 g. E.g. The relative molecular mass (M r) of Carbon Dioxide is 44.0. The molar mass of Carbon 23 Dioxide is 44.0 g. The mass of 6.02 x 10 Carbon Dioxide molecules is 44.0 g.

o







The volume occupied by 1 mole of a Gas, is called the molar volume. 3 o

3

The molar volume of ALL gases at r.t.p. is 24.0 dm = 24 000.0 cm r.t.p refers to Room Temperature and Pressure o r.t.p: Temperature = 25 C and Pressure = 1 atm 3 3 The molar volume of ALL gases at s.t.p. is 22.4 dm = 22 400.0 cm r.t.p refers to Standard Temperature and Pressure o s.t.p: Temperature = 0 C and Pressure = 1 atm 3 E.g. 1 mole of Chlorine gas at r.t.p. has a volume of 24.0 dm . 1 mole of Bromine gas at r.t.p. also 3 has a volume of 24.0 dm , despite having a larger atom than Chlorine.  

o

 

o

Mole = 

The number of moles present in a sample =

      

    

Example 1: Calculate the number of moles in a 142.0g sample of Chlorine gas. Molar Mass of Chlorine gas (Cl2)

= 2 x 35.5g = 71.0g

Number of moles ofpresent Cl2

   

= =

      

= 2 mols Example 2: Calculate the number of moles in a 100.0g sample of NaCl. Molar Mass of NaCl

Number of moles ofpresent Cl2

= 23.0g + 35.5g = 58.5 g    

= =

      

= 1.71 mols (3.s.f)


4.8 Molar Solutions (Concentration) 

3

The concentration of a solution refers to the amount of solute in 1 dm of solution. 3 3 1 000 cm = 1 dm

o 

Concentration can be presented in 2 ways: Concentration or Molarity

-3

Concentration in (g dm )= -3

Molarity (mol dm or M) =

   ()    ()

     ()    ()

  ( )

-3

Molarity (mol dm or M) =     () 3

Example 1: A 100 cm solution of HCl contains 1g of HCl. Concentration of Solution

= =

   ()    ()   

= 0.1 g dm Number of Moles of HCl

= =

3

   

= 0.0274 mols (3.s.f) Molarity of Solution

= =

     ()    ()    

= 0.00274 mol dm Molarity of Solution

= =

    ()   

= 0.00274 mol dm 

-3

  ( )

-3

The concentration of a solution changes when diluted

M 1V1 = M 2V2 M1 V1 M2 V2

= Original Molarity = Original Volume = New Molarity = New Volume 3

Example 1: a 10 cm sample of a 1M HCl solution is diluted to 50cm

3

M1V1 = M2V2 3 3 (1 M) x (10 cm ) = M2 x (50cm )      M = 2

 

= 0.20 M


4.9 Empirical Formulae 



The empirical formulae shows the simplest integer ratio of the different types of atoms in a compound. The empirical formulae for Benzene (C 6H6) is CH o The empirical formulae for Butane (C 4H10) is C 2H5 o The empirical formulae may be determined using the following table if the mass of each individual constituent elements are given: Element X

Element Y

Mass of each Individual Element Molar Mass Number of Moles Smallest Mole Divide by the smallest Mole Ratio Example 1: A Sample of Iron Sulphide contains 5.373g of iron and 4.627g of sulphur.

Mass of each Individual Element Molar Mass 

Number of Moles



Smallest Mole

Iron

Sulphur

5.373g

4.627g

56.0

32.0 

= 0.0959 mol



0.0959 mol 

Divide by the smallest Mole



Ratio

= 0.145 mol

0.0959 mol 

=1



2

= 1.51 3

Empirical Formulae = Fe2S3 Example 2: A compound contained (by mass) 23.3% Magnesium, 30.7% Sulphur and 46.0% Oxygen. Magnesium

Sulphur

Oxygen

23.3g

30.7 g

46.0 g

Mass of a 100g Sample Molar Mass Moles Smallest Mole Divide Ratio

24.0  

= 0.97 mol

0.96 mol  

=1

1

32.0  

= 0.96 mol

0.96 mol  

=1

16.0  

= 2.88 mol

0.96 mol 

1

Empirical Formulae = MgSO 3




=3

3

4.10 Molecular Formulae 



The molecular formulae shows the actual number of atoms of each element in each molecule of a compound Molecular Formulae = n x Empirical Formulae o

n=

   

Example 1: A Compound contained (by mass) 26.67% Carbon, 2.22% Hydrogen and 71.11% Oxygen. One mole of the compound has a mass of 90.1g.

Mass of a 100g Sample

Carbon

Hydrogen

Oxygen

26.67 g

2.22 g

71.11 g

Molar Mass Moles

12.0

1.0

16.0

 = 2.22 mol 

 = 2.22 mol 

 = 4.44 mol 

2.22 mol

2.22 mol

2.22 mol

Smallest Mole



Divide



Ratio

=1

1

Empirical Formulae

= CHO2

Empirical Mass

= 12.0g + 1.0g + (2 x 16.0)g = 45.0g

n

= =

 

=1

1

 

2

       

=2 Molecular Formulae

= 2 x CHO 2 = C2H2O4

Example 2: The empirical formulae of a compound is C 2H4O. Its relative molecular mass is 88 Empirical Formulae

= C2H4O

Relative mass

= (2 x 12.0) + (4 x 1.0) + 16.0 = 44.0

n

   

=

   

=2 Molecular Formulae

= 2 x C 2 H4 O = C4H8O2


=2

4.11 Theoretical Product Yield 

The theoretical product yield of a chemical reaction can be calculated from the mass (or volume for gases) of the reactants, using a balanced equation

Example 1: 4.0g of Methane is completely burnt in excess oxygen to yield Carbon Dioxide and Water. CH4 (g) + 2 O2 (g) = CO 2 (g) + 2 H 2O(l) Number of moles of Methane

= = =

         (  )   

= 0.25 mol Mole Ratio CH4 1

: :

CO2 1

Number of moles of CO 2 to be produced = Number of moles of Methane = 0.25 mol Mass of COproduced 2

= Moles x Molar Mass = 0.25 mols x [12.0g + (16.0g x 2)] = 0.25 mols x 44.0g = 11.0 g

Volume of CO 2 produced at r.t.p = Moles x 24.0 dm 3 = 0.25 x 24.0 dm 3 = 6.0 dm

3

Mole Ratio CH4 1

: :

H2O 2

Number of moles of H2O to be produced = 2 x Number of moles of Methane = 2 x 0.25 mol = 0.5 mol Mass of H2O produced

= Moles x Molar Mass = 0.5 mols x [(1.0g x 2) + 16.0g] = 0.25 mols x 18.0g = 9.0 g


4.12 Limiting Reagent 



A limiting reagent is a reactant that causes a reaction to stop once it's completely consumed. It limits the amount of products to be formed. The limiting reagent can be identified by comparing the number of moles of each reactants with the mole ratio of the reactants in the chemical equation.

Example 1: 5.6g of iron is burnt in 6.4g of sulphur to form iron (III) sulphide. 2Fe(s) + 3S (s) = Fe2S3 (s) Number of moles of Iron

= =

         

= 0.1 mol Number of moles of Sulphur

= =

         

= 0.2 mol Mole Ratio Fe 2 0.1

: : :

S 3 0.15 < 0.2

The limiting reagent is Iron Number of moles of Fe2S3 produced

= =

  

x Number of moles of Iron x 0.1 mol



= 0.05 mols Mass of Fe2S3produced

= Moles x Molar Mass = 0.05 mols x [(56.0g x 2) + (32.0g x 3)] = 0.05 mols x 208.0g = 10.4 g

Example 2: 0.05 moles of Zinc is added to 0.075 moles of HCl. Zinc (II) Chloride and H 2 gas is produced. Zn(s) + 2HCl(aq) = ZnCl2(aq) + H2(g) Number of moles of Zinc Number of moles of HCl

= 0.05 mol = 0.075 mol

Mole Ratio Zn 1 0.05

: : :

HCl 2 0.10 > 0.075

The limiting reagent is HCl Number of moles of ZnCl2 produced

= =

   

x Number of moles of HCl x 0.075 mol

= 0.0375 mols Mass of ZnCl produced 2

= Moles x Molar Mass = 0.0375 mols x [65.0g + (32.0g x 3)] = 4.84 0.0375 mols x 129.0g = g (3 significant figures)


Chapter 5 Energy and Chemical Reactions 5.1 Chemical Energy 

  

All chemical substances store chemical energy This energy can be converted into heat, light, electrical or sound energies o A reaction that gives out heat to the surrounding is an exothermic reaction A reaction that takes in heat from the surrounding is an endothermic reaction ∆H represents the change in heat energy of the reaction. o It is the difference between the energy content of the products and the reactants. o Exothermic Reactions have a negative ∆H o Endothermic Reactions have a positive ∆H

5.2 Bond Energies  





Bond Energies measures the strength of a covalent bond. When two atoms are joined together by a chemical bond, heat energy is released. o Hence, bond forming is exothermic When a chemical bond is broken, heat energy is consumed. Hence, bond breaking is endothermic o The amount of energy consumed in breaking a chemical bond is known as the bond energy The same amount of energy is produced when the same bond is formed o

∆H (Heat of reaction) = Total Heat Energy Absorbed - Total Heat Energy Released

Covalent Bond H-H Cl - Cl C-C C-H O-H Cl - H N-H O=O C=O N≡N C=C

Bond Energy / kJmol 436 242 348 412 463 431 388 496 743 945 838

-

Example 1: Combustion is an Exothermic Reaction CH4 + 2O2 → CO 2 + 2H2O Covalent Bonds in reactants = (4 x C - H bonds) + (2 x O = O bonds) Sum of Bond Energies in Reactants (Er) = (4 x 412kJ) + (2 x 496kJ) = 2240 kJ Covalent Bonds in Products = (2 x C = O bonds) + (4 x O - H bonds) Sum of Bond Energies in Product (Ep) = (2 x 743 kJ) + (4 x 463 kJ) = 3338 kJ Overall Heat of Reaction (∆H) = Total Heat Energy Absorbed - Total Heat Energy Released = 2240 kJ - 3338 kJ = -1098 kJ


5.3 Energy Profile Diagrams

5.4 Collision Theory 





Collision Theory states that a chemical reaction occurs when reactant particles collide with each other. However, Not all collisions will result in the formation of products. A collision is effective only when the reactant particles have enough energy to overcome the activation energy of the reaction, as well as having to collide in the proper orientation. Therefore, the speed of the reaction will depend on the number of effective collisions between the reactants.

5.5 Speed of reaction  



The speed of reaction refers to how fast reactants are used up or how fast products are formed Speed of reaction =

       

Factors that affect the speed of reactions include: Concentration of chemicals involved o An increase in concentration results in a higher speed of reaction More particles in a given volume results in an increase in frequency of effective collisions Temperature o An increase in temperature results in a higher speed of reaction At higher temperatures, the particles have more kinetic energy and thus move faster, this leads to an increase frequency of effective collisions At higher temperatures, more particles have the necessary energy to overcome the activation energy needed for an effective collision. o Pressure Changes in pressures only affects reactions where gases are involved. An increase in pressure results in a higher speed of reaction The same number of particles in a smaller volume results in an increase in frequency of effective collisions Particle Size o A decrease in particle size results in a higher speed of reaction Breaking up the particles results in greater total surface area, which in turn results in more particles being able to collide per unit time o Catalysts  

 



  

 



Three experimental methods to determine the speed of a reaction include: Measuring time for reaction to be completed Measuring quantity of products formed over a fixed time interval o By volume of gas produced By mass of product Measuring quantity of reactants left over a fixed time interval o By mass of Reactants Titration For acidic or basic reactants o

 

 


5.6 Catalysts 





 

 

A catalyst is a substance that changes the rate of reaction, but itself is chemically unchanged at the end of the reaction. A catalyst changes the rate of reaction by lowering the activation energy of a reaction o Because the catalysed activation energy is lower than the uncatalysed activation energy, the reaction will take place more quickly.

Different reaction will require different type of catalysts. o Each catalyst is usually specific to a particular reaction Catalysts are usually transition metals, or transition metal compounds Enzymes are an example of biological catalysts found in living cells, used to speed up the breaking down of giant molecules such as proteins or starch. A catalyst does not change the amount of products obtained A catalyst does not change the ∆H of a reaction.

5.7 Redox Reactions  







Redox is reaction where reduction and oxidation occurs simultaneously Reduction occurs when a substance o Gains Hydrogen o Or Loses Oxygen o Or Gains Electrons o Or Decrease in oxidation state A Reducing Agent causes another substance to be reduced o Hence a substance being reduced is an oxidising agent Oxidation occurs when a substance o Loses Hydrogen o Or Gains Oxygen o Or Loses Electrons o Or Increase in oxidation state An Oxidizing Agent causes another substance to be oxidized Hence a substance being oxidized is a Reducing agent o

OIL RIG Oxidation is Lose, Reduction is Gain ©2014 Marcus Ng Chemistry Notes for Secondary School O-Levels Combined Chemistry

Example 1: H2 (g) + CuO (s) → Cu (s) + H2O (l) 



The hydrogen gas gains oxygen, and thus is oxidized. (Undergoes oxidation) o The hydrogen gas is a Reducing agent The copper oxide loses oxygen, and thus is reduced. (Undergoes reduction) o The copper oxide is an Oxidizing agent

Example 2: 4CO (g) + Fe2O3 (s) → 3Fe (s) + 4CO2 (g) 



The carbon monoxide gas gains oxygen, and thus is oxidized. (Undergoes oxidation) o The CO gas is a Reducing agent The iron (III) oxide loses oxygen, and thus is reduced. (Undergoes reduction) o The Iron (III) oxide is an Oxidizing agent

Example 3: 4H2S (g) + Cl2 (g) → 2HCl (g) + S (s) 



The chlorine gas gains hydrogen, and thus is reduced. (Undergoes reduction) o The chlorine gas is a oxidizing agent The hydrogen sulphide loses hydrogen, and thus is oxidized. (Undergoes oxidation) o The hydrogen sulphide is an reducing agent

Example 4: FeSO4 (aq) + Zn (s) → ZnSO4 (aq) + Fe (s) 



The iron gains electrons (From oxidation state +2 to 0), and thus is reduced. (Undergoes reduction) o The iron is a oxidizing agent The zinc loses electrons (From oxidation state 0 to +2), and thus is oxidized. (Undergoes oxidation) o The zinc is an reducing agent

Example 5: 2 Na (s) + Cl2 (g) → 2NaCl (s) 



The chlorine gas gains electrons (From oxidation state 0 to -1), and is reduced. (Undergoes reduction) o The chlorine gas is a oxidizing agent The sodium loses electrons (From oxidation state 0 to +1), and is oxidized. (Undergoes oxidation) o The sodium is an reducing agent

5.8 Test for Oxidizing and Reducing Agents Type Oxdizing Oxdizing Oxdizing Reducing

Reducing

Agent KMnO4 Acidified Potassium Manganate (VII) K2Cr2O7 Acidified Cl2 Chlorine KI Aqueous Potassium Iodide FeSO4 Aqueous Iron (II) Sulphate

Half Equation -

MnO4 (aq) + 8H Cr2O7

2(aq)

+

(aq) +

+

+ 14 H

Colour Change

-

2+

5e → Mn

(aq)

-

(aq)

+ 5 e → 2Cr -

3+ (aq)

-

Cl2(g) + 2 e → 2Cl (aq) -

I (aq) → I2(aq) + 2 e

Purple to colourless

+ 4H2O(l) + 7H2O(l)

Orange to Green Greenish Yellow to Colourless Colourless to

-

Reddish Brown 2+

Fe

(aq)

→ Fe

3+ (aq)

+e

-


Greenish to Yellow

Chapter 6 Acid, Bases and Salts 6.1 pH +

pH measures the concentration of H ions in a solution The pH scale ranges from 0 to 14 o Acids have a pH value of less than 7 The lower the pH, the stronger the acid Sulphuric Acid in Car Batteries (pH 1) is a much stronger acid than lemon juice (pH 2) Bases and Alkalis have a pH value of greater than 7 o The higher the pH, the stronger the base/ alkali Sodium Hydroxide in bleach (pH 13) is a much stronger base than ammonia in fertilizer (pH 11) o Neutral solutions (like water) have a pH of exactly 7

 

 

 

6.2 The pH most Indicators accurate way of measuring pH is using a pH meter 

An approximate way of measuring pH is using an indicator An indicator can tell the pH by with colour changes

 

Universal Indicator 0

1 Red

2

3 4 Orange

2

3 Red

5 6 Yellow

7 Green

8 9 Green-Blue

10

11 Blue

12

13 14 Violet

12

13

14

Litmus 0

1

4

5

6

7

8

9

10

11 Blue

6

7

8

9

10

11 12 13 Pink - Purple

14

Phenolphthalein 0

1

2

3

2

3

4 5 Colourless

Methyl Orange 0

1

4

5

6

7

8

9 Yellow

10

11

12

13

14

4

5

6

7

8

9 Green

10

11

12

13

14

3 4 Yellow

5

6

7

8

9

10

11 Blue

12

13

14

Red Screened Methyl Orange 0

1 Red

2

3 Grey

Bromothymol Blue 0

1

2


6.3 Acids Acids are substances that produces H + ions in Water.

Properties of Acids 

Acids are sour Lemon's sourness is from citric acid o Vinegar's sourness is from ethanoic acid Acids have pH of less than 7 o Acids change Blue Litmus paper Red Organic acids are acids that contain the -COOH group o Examples of organic acids are Ethanoic Acid and Citric Acid Mineral acids are acids that are not organic o







o o

Mineral acids are much stronger acids than organic acids Examples of mineral acids are Hydrochloric Acid and Nitric Acid

6.4 Reactions of Acids

Reaction of Acids with Bases  

Acids will react with bases to form an inorganic Salt and water only This reaction is called Neutralization HX(aq) + ZOH(aq) → ZX (aq) + H 2O(l)

Example 1 HCl(aq) + NaOH(aq) → NaCl (aq) + H2O(l) Example 2 H2SO4(aq) + Ca(OH)2(aq) → CaSO4(aq) + 2 H 2O(l)

Reaction of Acids with Metals  

Acids will react with Metal to form an inorganic Salt and Hydrogen Gas only This is due to displacement (to be covered under the Reactivity Series in the topic of Metals) 2 HX(aq) + 2 M (s) → 2 MX (aq) + H2(g)

Example 1 2 HCl(aq) + 2 Na (s) → 2 NaCl(aq) + H2(g) Example 2 H2SO4(aq) + Mg (s) → MgSO4(aq) + H2(g)

Reaction of Acids with Carbonates 

Acids will react with Carbonates to form an inorganic salt, carbon dioxide and water. 2 HX(aq) + ZCO 3(aq) → ZX 2(aq) + CO2(g) + H2O(l)

Example 1 2 HCl(aq) + Na2CO3(s) → 2 NaCl(aq) + CO 2(g) + H2O(l) Example 2 H2SO4(aq) + CaCO3(s) → CaSO4(s) + CO2(g) + H2O(l)


Common Acids Acids Hydrochloric Acid Sulphuric Acid Nitric Acid Ethanoic Acid Phosphoric Acid Hydrofluoric Acid Hydrobromic Acid Hydroiodic Acid Carbonic Acid

Formulae HCl H2SO4 HNO3 CH3COOH H3PO4 HF HBr HI H2CO3

Cation H+ H+ H+ H+ H+ H+ H+ H+ H+

Anion ClSO4 NO3CH3COOPO43FBrICO3 -

6.5 Bases and Alkalis Bases are substances that react with an Acid to form a Salt and Water Only

Properties of Bases   

 

Edible bases taste bitter Alkalis feel slippery Bases have pH of more than 7 o Bases change Red Litmus paper Blue Bases are usually Metal oxides or Metal hydroxides Soluble Bases are called alkalis o Group I hydroxides are readily soluble o Group II hydroxides are sparingly soluble o Group III or Transition Metal hydroxides are generally insoluble

6.6 Reactions of Bases

Reaction of Bases with Ammonium Salts 

Bases will react with Ammonium Salts to form an inorganic salt, ammonia gas and water NH4X(aq) + ZOH(aq) → ZX(aq) + NH3(g) + H2O(l)

Example 1 NH4Cl(aq) + NaOH(aq) → NaCl(aq) + NH 3(g) + H2O(l) Example 2 (NH4)2SO4(aq) + Ca(OH)2(aq) → CaSO4(aq) + NH 3(g) + 2 H2O(l)

Precipitation of Insoluble hydroxides 

Alkalis are used to precipitate out insoluble hydroxides from solutions of their salt

Example 1 2 NaOHaq) + CuSO4(s) → Na2SO4(aq) + Cu(OH)2(s) (Blue precipitate) Example 2 2 NaOHaq) + MgCl2(s) → 2 NaCl(aq) + Mg(OH)2(s) (White precipitate)


Common Bases Bases Sodium Hydroxide Potassium Hydroxide Ammonium Hydroxide Calcium Hydroxide Magnesium Hydroxide Barium Hydroxide Aluminum Hydroxide Zinc Hydroxide

Formulae NaOH KOH NH4OH Ca(OH)2 Mg(OH)2 Ba(OH)2 Al(OH)3 Zn(OH)2

Cation Na+ K+ NH4+ Ca + Mg2+ Ba2+ Al3+ Zn +

Anion OHOHOHOHOHOHOHOH-

6.7 Oxides   

Oxides are formed when elements burn in Oxygen There are 4 types of Oxides: Acidic, Basic, Amphoteric and Neutral Non-Metallic oxides are acidic o They have similar properties as acids, as well as undergo similar reactions as acids o They form acids when dissolved in water Carbon Dioxide dissolves in water to form carbonic acid (H 2CO3) Metallic oxides are basic o They have similar properties as bases, as well as undergo similar reactions as bases Some Metallic oxides are amphoteric They show both acidic and basic properties o They can neutralize both acids and bases. Aluminum oxide can react with hydrochloric acid to form aluminum chloride and water Aluminum oxide can react with sodium hydroxide to form sodium aluminate and water o Some examples are Aluminum Oxide (Al 2O3), Zinc Oxide (ZnO) and Lead (II) Some Non-Metallic oxides are neutral (Pure) o They show neither acidic nor basic properties o Some examples are Dihydrogen Oxide, Carbon Monoxide (CO) and Nitrogen Oxide (NO) 





  



6.8 Solubility of Various Salts Soluble Salts All Nitrates All Chlorides except Lead (II) Chloride All Sulphate except Lead (II) Sulphate

All Group I and Ammonium Salts

Insoluble Salts

All Lead (II) Salts except Lead (II) Nitrate All Carbonates except for Group I and Ammonium Carbonates All Sulphides (s -) except for Group I and Ammonium Sulphides All hydroxides except for Group I and Ammonium hydroxides All oxides except for Group I and Ammonium Oxides

Also Insoluble: Barium Sulphate, Calcium Sulphate and Silver Chloride Sparingly soluble: Group II Hydroxides and Oxides


6.9 Preparation of Salts 

The method to prepare a salt depends on its solubility

Salt

Insoluble

Precipitation Method

Soluble

Group I or Ammonium Salt

NOT Group I or Ammonium Salt

Titration

Crystalization

React Metal with Acid

React Metal Carbonate with Acid

React Metal Oxide with Acid

6.10 Precipitation method 

The precipitation method to prepare an Insoluble salt

Step 1: Choosing the reactants. (They must be soluble) Step 2: Mix the reactants. Step 3: Wait for the insoluble salt to precipitate out. Stirring or heating may speed the reaction along. Step 4: Filter out the insoluble salt.


6.11 Crystallization method 

The crystallization method is used to prepare a soluble salt that does not contain Group I or Ammonium ions.

Step 1: Choosing the reactants. (Use the acid containing the anion, and the metal/ oxide/ carbonate.) Step 2: Mix the reactants. Let the metal/ metal oxide/ metal carbonate be in excess. Step 3: Wait for the reaction to complete. Stirring or heating may speed the reaction along. Step 4: Filter out the excess metal/ metal oxide/ metal carbonate Step 5: Heat the filtrate till saturated (when a thin layer of crystals are formed. Step 6: Leave the filtrate to cool for more crystals to form. Step 7: Filter out the crystals and dry

6.12 Titration method  

The Titration method is used to prepare a soluble salt that contains Group I or Ammonium ions. This method is based on the neutralization reaction.

Step 1: Choosing the reactants. (Use the acid containing the anion & the hydroxide containing the cation.) 3

Step 2: Pipette out 25 cm of one of the reactants into a conical flask. Add a few drops of indicator Step 3: Add the other reactant into a burette Step 4: Add the reactant in the burette into the conical flask drop by drop. Stop when the colour changes. Step 5: Heat the filtrate till saturated (when a thin layer of crystals are formed. Step 6: Leave the filtrate to cool for more crystals to form. Step 7: Filter out the crystals and dry


Chapter 7 Periodic Table 7.1 Periodic Trends (An Overview) 





The Periodic Table is an arrangement of elements with an increasing number of Protons Number of valence electrons increases across a period (Left to right) from 1 to 8. o The Periodic Table is arranged in vertical groups and horizontal periods. o Elements of the same period have the same number of electron shells. Elements of the same group have the same number of valence electrons. o o Elements of the same group have similar chemical properties & form compounds with similar chemical formulae. A zig-zag diagonal line divides the metals and non metals. Elements near the line are called metalloids & have characteristics of both metals & non-metals. o

7.2 Electronegativity



Electronegativity refers to the ability to gain electrons Less electronegative elements have a greater metallic character. Elements become more electronegative across the period (Left to right) Elements become less metallic across the period o Elements become less electronegative down a group (Top to Bottom)



Elements more metallic element, down a group Chlorine is thebecome most electronegative while Francium is the least Electronegative.

  

o

Name of Various Groups of Elements

Note: Only Group I (Alkali Metals), Group VII (Halogens), Group VIII (Noble Gases) and Transition Metals are examinable.


7.3 Group I Metals: Alkali Metals 

Group I Elements reacts with water to form Alkalis and hydrogen gas (Thus the name Alkali Metals). 2 Na(s) + 2 H2O(l) → 2 NaOH (aq) + H2(g)





Group I Metals are strong reducing agents Physical properties of Group I Metals include: o They easily react with cold water and air, and thus have to be stored in oil or vacuum. o They have low densities. The densities increases down the group. o They have low melting points. The melting point increases down the group. They are shiny and silvery solids. o They are very soft and can be easily cut with a knife or razor. o 



o

They are good conductors of electricity and heat.

7.4 Group VII: Halogens 



Group VII Halogens form diatomic molecules with a single covalent bond Trends of Group VII Halogens include: The melting and boiling point increases down the group. o Fluorine and Chlorine are gaseous at room temperature Bromine is liquid at room temperature Iodine and Astatine are solids at room temperature o The colours of the Halogens get darker down the group. Fluorine is pale yellow Chlorine is yellowish green Bromine is reddish brown Iodine and Astatine are black o The Halogens get less reactive down the group Fluorine is the most reactive, and astatine is the least reactive   

   







Group VII Halogens are strong oxidizing agents Halogens undergo displacement reactions. o A halogen in a salt can be displaced by a more reactive halogen For example; Fluorine is the most reactive, and astatine is the least reactive, hence the astatine in an astatine salt can be displaced by fluorine gas 

2 NaAt(s) + F2(g) → 2 NaF (s) + At2(s) 



Physical properties of Group VII Halogens include: o They have low boiling and melting points. They do not conduct heat or electricity in any state. o They are sparingly soluble in water o They are soluble in organic solvent like CCl 4. o Some uses of Halogens o Fluoride is used in toothpaste to prevent tooth decay o Chlorine and Iodine is used to kill bacteria Iodine is needed by the human body for proper thyroid gland function o


7.5 Group VIII: Noble Gases 





Group VIII Noble Gases are chemically unreactive due to their stable octet structure (With fully filled valence shells) Noble Gases exists as monoatomic gases at room temperature. o Noble gases have very low boiling and melting points Some uses of Noble Gases o Argon is used to fill light bulbs o Neon is used to fill coloured glowing tubes o Helium is used to fill weather balloons

7.6 Transition Metals 

Properties of Transition Metals include: They are strong hard metals with high boiling and melting points. o o o

They have high density They form coloured compounds Iron (II) oxide is green, while Iron (III) Oxide (Rust) is reddish brown. They form ions with variable charges 2+ 3+ Iron can form Fe and Fe ions. They are used as catalyst They are used to make alloys Steel is an alloy comprising of Iron and Carbon 

o



o o



Summary (Trends)


Chapter 8 Metals 8.1 Physical Properties of Metals 





High Boiling and Melting point. General exception to these are Mercury and Group I and Group II Metals o Good Conductor of Heat and Electricity o Due to sea of delocalized electrons Malleable (Ability of being flattened) & Ductile (Ability to be pulled into wires) o Due to metallic bonding, in which the layers of atoms can easily slide over each other.

8.2 Alloys 

 



An alloy is a mixture of metal with another element This second element may be both either a metal or a non metal o Pure metals are usually too soft to be used. Alloys strengthens metals to be used by disrupting the orderly arrangement of the metal atoms with foreign atoms of different sizes. Some metals, like iron, oxidize or rusts easily. Hence alloys of these metals may be used in place of the metals due to their resistance to o oxidization or corrosion.

Examples: Alloys Bronze Brass Pewter Industrial Steel Stainless Steel Chromium Steel High Speed Steel

Constituent Elements Copper and Tin Copper and Zinc Tin, Antimony and Copper Iron and Carbon Iron, Chromium and Nickel Iron and Chromium Tungsten and Vanadium

Uses Trophies Musical Instruments and Electrical plug pins Dinnerware like plates and teapots Scaffoldings Cutlery and surgical instruments Ball Bearings High Speed Drills

8.3 Reactivity Series  



Metals differ greatly in their chemical reactivity Very reactive metals are unstable as a metal, but form very stable compounds o These metals are not found uncombined in nature Less reactive metals are more stable as a metal. These metals can be found uncombined in nature o

Most Reactive Potassium

Sodium

Calcium

Magnesium

Aluminum

Carbon

Zinc

Iron

K

Na

Ca

Mg

Al

C

Zn

Fe

Potato

Salad

Can

Make

A

Cunning

Zebra

Itchy

Tin

Lead

Hydrogen

Copper

Mercury

Silver

Gold

Sn

Pl

H

Cu

Hg

Ag

Au

These

Large

Helicopters

Can

Make

Some

Giddy Least Reactive


8.4 Chemical Reactions of Metal 

All metals undergo displacement reactions A more reactive metal is able to displace a less reactive metal from its compounds o E.g. Displacement of the less reactive copper by the more reactive zinc o Zn(s) + CuSO 4 (aq) → Cu(s) + ZnSO 4 (aq) o

E.g. Displacement of the less reactive lead by the more reactive magnesium Mg(s) + PbO (s) → Pb (s) + MgO



(s)

Most Metals can react with water to produce hydrogen gas and either hydroxides or metal oxides. o E.g. Reaction of Sodium with cold water 2Na(s) + 2H 2O(l) → 2NaOH(aq) + H2 (g) o

E.g. Reaction of Aluminum with steam 2Al(s) + 3H2O(l) → Al2O3 (aq) + 3H2 (g)



Most Metals can react with acids to produce salts and hydrogen gas o E.g. Reaction of Calcium with Hydrochloric Acid Ca(s) + 2HCl (aq) → CaCl2 (aq) H2 (g) Metals

Reaction with Water

Potassium Sodium

K Na

Calcium

Ca

Magnesium

Mg

Aluminum Zinc Iron Tin Lead Copper Mercury Silver Gold

Al Zn Fe Sn Pb Cu Hg Ag Au

Reacts vigorously with cold water to produce hydroxides and hydrogen gas Reacts slowly with cold water to produce hydroxides and hydrogen gas Reacts with Steam to produce metal oxides and hydrogen gas

Reaction with Acids Reacts explosively with acids to produce salts and hydrogen gas Reacts vigorously with acids to produce salts and hydrogen gas Reacts rapidly with warm acids to produce salts and hydrogen gas

Reacts slowly with hot salts and hydrogen gasacids to produce

Does not react to water or steam Does not react with acids


8.5 Chemical Reactions of Metal Compounds 

The oxides of the less reactive metals can be reduced by carbon to produce the metal & CO 2 only. o E.g. Reduction of Zinc oxide C(s) + 2ZnO(s) → 2Zn(s) + CO 2 (g)



The oxides of the less reactive metals can be reduced by hydrogen to produce the metal & H 2O only. o E.g. Reduction of Lead oxide PbO(s) + H2(g) → Pb (s) + H 2O (l)



Some of the carbonates of the less reactive metals can be decomposed upon heating. o E.g. Decomposition of Copper Carbonate CuCO3(s) → CuO o

(s)

+ CO2 (g)

E.g. Decomposition of Silver Carbonate 2Ag2CO3(s) → 4Ag Reduction of Oxides by Carbon

Metals Potassium Sodium Calcium Magnesium Aluminum Zinc Iron Tin Lead Copper Mercury Silver

K Na Ca Mg Al Zn Fe Sn Pb Cu Hg Ag

Gold

Au

(s)

+ O 2 (g) + 2CO2 (g)

Reduction of Oxides by Hydrogen

Decomposition of Metal Carbonates Does not decompose

Does not reduce

Does not reduce

Metal carbonate decomposes upon heating into metal oxides and carbon dioxide gas Metal Oxides get reduced by carbon to form metal and carbon dioxide only

Metal Oxides get reduced by hydrogen to form metal and water only Carbonate decomposes upon heating into metal, O 2 and CO2 gas

8.6 Extraction of Metals 

Metals high up in the reactivity series do not exist in the free state, but can be found in the form of metal ores or metal salts These ores are normally in the form of oxides, sulphides or carbonates. o o Some common ores include Metal Aluminum Iron Zinc







Ore Bauxite Haematite Zinc Blende

Compound in Ore Aluminum Oxide Iron (III) Oxide Zinc Sulphide

Very reactive metals form very stable compounds, hence they can only be extracted by electrolysis. o All metals above carbon on the reactivity series can only be extracted by electrolysis. Less reactive metals can be extracted through reduction by carbon, hydrogen or more reactive metals o All metals below carbon on the reactivity series can only be extracted through reduction Metals with low reactivity can be found usually in the free state, or can be extracted through heating. o This refers to metals like mercury, silver and gold.


8.7 Extraction of Iron 

1. 2. 3.

Iron is extracted from haematite (Fe 2O3) through reduction in a blast furnace.

Haematite (Iron (III) Oxide), Limestone (Calcium Carbonate) and Coke (Carbon) are fed into the blast furnace Hot air is fed into the bottom of the furnace The Coke burns in the hot air to form carbon dioxide. C (S) + O2 (g) → CO 2 (g)

4.

The carbon dioxide is further reduced to carbon monoxide CO2 (g) + C

5.

(S)

→ 2CO (g)

The carbon monoxide reduces the haematite to iron CO (g) + Fe2O3 (s) → 2Fe (s) + CO2 (g)

6. 7.

The molten iron form is filled with sand particles, which can be removed using the limestone The Limestone is decomposed by heat to produce calcium oxide and carbon dioxide. CaCO3 (g) → CaO (s) + CO 2 (g)

8.

The calcium oxide reacts with the sand to form slag (calcium silicate) CaO (g) + SiO 2 (s) → CaSiO3 (s)

9. The slag is less dense than iron. Hence it floats on the molten iron & is removed from the top tap 10. The molten iron is removed from the bottom tap


8.8 Recycling of metals  



Metal ores resources are finite and limited. Hence it is important to recycle metals Some advantages of recycling include o Recycling saves energy required to extract metals from ores o Recycling reduce emission of greenhouse gases like carbon dioxide, produced in the extraction of metals like iron from ores o Recycling preserves scarce non-renewable raw materials Recycling reduces environmental air pollution and water pollution. o Recycling reduced the amount of land required for the disposal of metals through landfills o It is cheaper to recycle some metals like aluminum, than to extract them from the earth's crust o Some disadvantages of recycling include Recycling is a time consuming process o Recycling takes up a high amount of effort and human resources. o

8.9 Rusting of Iron  

In the presence of water and oxygen, Iron rusts Some methods of preventing rust include: Coating the iron with a substance to prevent air and water from coming into contact with the metal o surface. These substances include Paint Oil or grease Electroplating the iron with a less reactive metal like tin or copper o Sacrificial protection.   

Chapter 9 Electrolysis Not part of the O-Level Combined Chemistry Syllabus


Chapter 10 Organic Chemistry 10.1 Introduction to Organic Chemistry 







Organic Chemistry is the study of Carbon based compounds o Except for Carbon Monoxide, Carbon Dioxide and metal Carbonates Most organic compounds also have hydrogen, and some also have oxygen. Organic compounds with carbon and hydrogen atoms are called hydrocarbons o Important reminder: Carbon Atoms only can form 4 bonds around each one In organic chemistry, there are many compounds with similar chemical properties, and have a general formula This family of compounds is known as a homologeous series o All compounds in a homologus series typically have a common functional group, and differ by a o o



CH2 unit. Compounds in a homologeous series have similar chemical properties but different physical properties like boiling and melting points

4 Main homologeous series to be taught include Alkanes o Has no functional group Has the suffix -ane Has the general formula of C nH2n + 2 o Alkenes has a double bond between 2 carbon atoms Has the C=C functional group Has the suffix -ene Has the general formula of C nH2n o Alcohols Has the -OH functional group Has the suffix -anol Has the general formula of C nH2n + 1OH o Carboxylic Acid   

   

  

  

Has the -COOH functional group Has the suffix -anoic Acid Has the general formula of C nH2nO2 or C n-1H2n-1COOH

1 MethAlkanes CnH2n + 2

Alkenes CnH2n

Alcohol

Carboxylic Acid CnH2nO2

4 But-

Methane

Ethane

Propane

Butane

CH4

C2 H6

C3H8

C4H10

Ethene

Propene

Butene

C2 H4

C3H6

C4 H8

Methanol

Ethanol

Propanol

Butanol

CH3OH

C2H5OH

C3H7OH

C4H9OH

Ethanoic

Propanoic

Acid CH3COOH

Acid C 2H5COOH

-ane

-ene

Functional Groups CnH2n + 1OH

No. Of Carbons 2 3 EthProp-

-anol

Methanoic Acid - Acid anoic

HCOOH

Butanoic Acid C 3H7COOH


10.2 Petroleum and Crude Oil 



Crude oil is a mixture of many thousands of different hydrocarbons with different properties. To make crude oil useful, batches of similar compounds with similar properties need to be sorted o and separated by fractional distillation. o In fractional distillation, the crude oil is heated to make it vaporise. The vapour is then cooled. Different fractions of the oil are collected at different temperatures. o

The larger hydrocarbons are not as useful as the smaller hydrocarbon Cracking is a process that can be used to break larger hydrocarbons into smaller ones o Cracking is done by passing the vaporised hydrocarbon over a solid catalyst o Cracking produces alkenes o


10.3 Isomerism 

Organic compounds with same chemical formula but different structural formula are known as isomers

Example 1: Isomers of Butane Chemical Formula: C4H10

Example 2: Isomers of Butanol Chemical Formula: C 4H9OH

Example 3: Propane Chemical Formula: C 3H8

(EXAMPLE OF WHAT'S NOT A SET OF ISOMERS) ALL 3 are the same structure


10.4 Alkanes   

Alkanes are organic compounds with only Carbon and Hydrogen atoms with NO Functional groups Alkanes have the general formula of C nH2n + 2 Alkanes are generally unreactive. Methane

Ethane

Propane

Butane CH3CH2CH2CH3

HCH3

CH3CH3

CH3CH2CH3

Or CH3CHCH3 CH3

10.4a Combustion 

Alkanes undergo Combustion o Alkanes burn in Oxygen to form Carbon Dioxide and Water Vapour

Example 1: Combustion of Methane, CH 4

CH4 + 2O2 →

CO2

+ 2H2O

Example 2: Combustion of Butane, C 4H10

2C4H10 + 13O2

→

8CO2 + 10H2O

10.4b Substitution Reaction 

In the presence of light, Alkanes undergo Substitution reaction with halogens Observation: The coloured halogens will decolourise o

Example 1: Substitution Reaction of Methane, CH 4 with chlorine gas, Cl2

CH4 + Cl2

→ CH3Cl

+ HCl

Example 2: Substitution Reaction of Butane, C 4H10 with bromine, Br 2

C4H10 + Br2 →

C4H9Br

+ HBr

Note: The substitution reaction could proceed further

C4H10 + Br2 → C4H9Br + HBr C4H9Br + Br2 → C4H8Br2 + HBr C4H8Br2 + Br2 → C4H7Br3 + HBr


10.5 Alkenes   

Alkenes are organic compounds with a double bond between 2 carbon atoms Alkanes have the general formula of C nH2n Alkenes are unsaturated organic compounds (Has 1 or more double bonds) o Alkanes are saturated organic compounds (Has no double bonds) Ethene

H2C=CH2

Propene

Butene H2C=CHCH2CH3

Or

Or

CH3CH=CHCH3

H2C=CHCH3

H2C=CCH3 CH3

10.5a Combustion 

Alkenes undergo Combustion o Alkenes burn in Oxygen to form Carbon Dioxide and Water Vapour

Example 1: Combustion of Butene, C 4H8

C4H8 + 6O2 → 4CO2 + 4H2O 10.5b Substitution Reaction 

In the presence of light, Alkenes can undergo Substitution reaction with halogens o Observation: The coloured halogens will decolourise

10.5c Addition Reaction with Halogens 

Alkenes can undergo addition reactions with halogens in the absence of light o Observation: The coloured halogens will decolourise

Example 1: Addition of aqueous bromine, Br2 to Ethene, C2H4

C2H4 + Br2

→

C2H4Br2

Note: This can be used as a test to differentiate Alkanes from Alkenes. Alkenes can decolourise bromine in the absence of light, while Alkanes cannot decolourise bromine in the absence of light Example 2: Addition of aqueous bromine, Br2 to Butene, C4H8

C4H8 + Br2

→

C4H8Br2

Example 3: Addition of chlorine gas, Cl2 to Butene, C4H8

C4H8 + Cl2

→

C4H8Cl2


10.5d Addition Reaction with Hydrogen Gas 

Alkenes can undergo addition reactions with Hydrogen Gas This process is known as hydrogenation o Hydrogenation is used to change vegetable oil into margarine Hydrogenation is used to change Alkenes to Alkanes o Reaction Conditions: 200 C with Nickel Catalyst  



Example 1: Addition of Hydrogen Gas, H2 to Ethene, C2H4 o

Temperature: 200 C Pressure: 1 atm Catalyst Used: Nickel Catalyst

C2H4 + H2 → C2H6

10.5e Addition Reaction with Water Vapour 

Alkenes can undergo addition reactions with Water vapour This process is known as Hydration o Hydrogenation is used to change Alkenes to Alcohol o Reaction Conditions: 300 C, 60 atm with Phosphoric Acid Catalyst 



Example 1: Addition of Water Vapor, H 2O to Ethene, C 2H4 o

Temperature: 300 C Pressure: 60 atm Catalyst Used: Phosphoric Acid (H 3PO4)

C2H4 + H2O → C2H5OH 10.5f Addition Polymerisation 

Alkenes can undergo addition reactions with itself to form long chains of polymers This process is known as Polymerisation o

Example 1: Polymerisation of Ethene to form polyethene

nC2H4 → - [C2H4]n-



Polyethene can be used to make plastic items like plastic bags, plastic bottles, etc...


10.6 Alcohol    



Alcohols are organic compounds with a -OH functional group Alcohols have the general formula of C nH2n+1OH Like Alkanes and Alkenes, Alcohol can undergo combustion and substitution reactions. Alcohol can be prepared through the hydration of alkenes (see 10.5e), or through the fermentation (see 10.6a) Alcohols can be used as an organic solvent, or as fuel. Methanol

Ethanol

Propanol

Butanol

CH3CH2CH2OH 4 Isomers CH3OH

Or

CH 3CH2OH

Check 10.3 Example 2 CH3CHCH3 OH

10.6a Fermentation 

Alcohol can also be prepared through the fermentation Enzymes are added to break down the glucose in sugar or starch to produce ethanol and carbon o dioxide. The fermented mixture is then filtered and the alcohol is extracted through fractional distillation. o Note: Fermentation must be carried out in an oxygen free environment to prevent the alcohol from o turning into carboxylic acids. o

Temperature: 37 C Pressure: 1 atm Catalyst Used: Enzymes in Yeast

C6H12O6

→

2C2H5OH + 2CO2

10.6b Oxidation Reaction 

Alcohols oxidize to form carboxylic acids and water For the oxidation, Acidified Potassium Dichromate (VI), Acidified Potassium Manganate (VII), or o even atmospheric oxygen can be used.

Example 1: Oxidation of Ethanol using Acidified Potassium Dichromate (VI) Conditions: r.t.p Observations: Orange Acidified Potassium Dichromate (VI) changes to green

C2H5OH + Cr2O72- + 10H+ → CH3COOH + 2Cr3+ + 6H2O Example 2: Oxidation of Propanol using Acidified Potassium Manganate (VII) Conditions: r.t.p Observations: Purple Acidified Potassium Manganate (VII) decolourises

C3H7OH + MnO4-

→

C2H5COOH + MnO2 + H2O

Example 3: Oxidation of Butanol using Atmospheric Oxygen

C4H9OH + O2

→

C3H7COOH + H2O

10.6c Condensation Reaction: Esterification 

See 10.7b


10.7 Carboxylic Acid   

Carboxylic Acids are organic compounds with a -COOH functional group Carboxylic Acids have the general formula of C nH2nO2 or C n-1H2n-1COOH Carboxylic Acids are weak acids + o Weak acids only partially dissociate in water to form H ions o As a (weak) acid, carboxylic acids also undergo all acid reactions Methanoic Acid

Ethanoic Acid

Propanoic Acid

Butanoic Acid CH3CH2CH2COOH

HCOOH

CH3COOH

Or

CH3CH2COOH

CH3CHCH3 COOH

10.7a Acid Reactions 





Carboxylic Acids can undergo Neutralization o Carboxylic Acids reacts with bases to form salt and water only Carboxylic Acids can undergo reactions with metals o Carboxylic Acids reacts with metals to form salt and hydrogen gas only Carboxylic Acids can undergo reactions with metal carbonates o Carboxylic Acids reacts with metal carbonates to form salt, water and carbon dioxide

10.7b Condensation Reaction: Esterification  



Esterification is a condensation reaction Alcohols and carboxylic acids react to form Esters o Esters have the functional group -OCOEster are named after the alcohol and the carboxylic acid used to prepare the ester o The first term follows the alcohol used Methanol used = methyl Ethanol used = ethyl  

o

The second term follows the carboxylic acid used Methanoic Acid used = methonate Ethanoic Acid used = ethanoate o Hence: E.g. Ethanol and butanoic acid produces Ethyl Butanoate Esters are sweet-smelling compounds Reaction Conditions: Sulphuric Acid Catalyst  

 

Example 1: Preparation of Methyl Ethanoate Reagents: Methanol and Ethanoic Acid Conditions: Sulphuric Acid as Catalyst Observations: A sweet smell is produced

CH3OH + CH3COOH

→

CH3OCOCH3 + H2O

Example 2: Preparation of Methyl Propanoate Reagents: Methanol and Propanoic Acid Conditions: Sulphuric Acid as Catalyst Observations: A sweet smell is produced

CH3OH + C2H5COOH

→

CH3OCOCH2CH3 + H2O


10.8 Macromolecules 



Macromolecules are giant molecules that are formed when thousands of smaller units of identical molecules or atoms are joined together This smaller units are called monomers o The process of joining these monomers into a macromolecule is called polymerisation o There are 2 types of macromolecules Natural Macromolecules o Synthetic (man-made) Polymers o Synthetic polymers can be made through addition polymerisation (addition polymers) or condensation polymerisation (Condensation Polymers) 

10.8a Addition Polymers  

Addition polymers are made from unsaturated monomers through an addition reactions Alkenes undergo addition polymerisation to form polyalkenes

10.8b Condensation Polymers 

Condensation polymers are made from condensation reactions A small molecule like water is produced through condensation reaction Monomers used in condensation reactions have 2 functional group on both ends o A Diol has 2 -OH functional group on either sides A Dicarboxylic Acid has 2 -COOH functional group on either sides o o A Diammine has 2 -NH 2 functional group on either sides o



Diol HO

Dicarboxylic Acid OH

HOOC

COOH

Diammine H2 N

NH2

= -[CH2]n

2 condensation reactions that can be used to produce condensation polymers are Esterification (Between diols and dicarboxylic acids) The -OCO- functional group between the alcohol and the acid is called the ester linkage

o

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Terrylene is a polymer with ester linkages, made from diols and dicarboxylic acids. Amide Condensation (Between diamine and dicarboxylic acids) The -CONH- functional group between the acid and the amine is called the amide linkage Nylon is a polymer with amide linkages, made from diamines and dicarboxylic acids. 

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Chapter 11 The Environment 11.1 Air 

Air comprises of o ≈ 78% Nitrogen Gas (N2) o ≈ 21% Oxygen Gas (O 2) o ≈ 1% Argon Gas (Ar) o Very small amounts of Carbon Dioxide and other rare gases.

11.2 Carbon Cycle & the Greenhouse Effect 

The carbon cycle shows how carbon is circulated around the world o All living creatures, plants, animals and humans, release carbon dioxide as part of respiration All living creatures, plants, animals and humans, also release carbon dioxide through decay and o decomposition Plants consume carbon dioxide during photosynthesis o Animals and Humans consume carbon in the form of food (be it in the form of both plants or other animals) Animals and Humans release carbon in the form of methane o These two gases, Carbon Dioxide and Methane, are major contributors to the green house effect. o

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Greenhouse gases, such as methane and carbon dioxide, are responsible for the green house effect, which traps heat in our earth's atmosphere. The green house effect is essential to sustaining life as the earth would otherwise be too cold to survive. However, too much greenhouse gas leads to global warming, which results in o Melting of polar caps o Rising sea levels causing floods in low lying land o Changing weather patterns such as increase in rainfall in some areas, and possibly causing floods o Changing weather patterns such as decrease in rainfall in some areas, resulting in an increase in number of deserts, as well as possible famine due to crop distruptions. Some causes of the increase of Greenhouse gases are o Increase in use of fossil fuel o Deforestation o Decay of vegetation due to deforestation o increased farming of rice fields

11.3 Carbon Monoxide 

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Besides carbon dioxide, the burning of fossil fuels also produces Carbon Monoxide, especially when there is insufficient oxygen. The major source of carbon monoxide is from the burning of petrol in vehicles o Carbon Monoxide is harmful as it is a poisonous gas It binds with the haemoglobin in our blood and prevents it from carrying the oxygen that the body o needs. The release of Carbon Monoxide can be reduced by the use of catalytic converters in vehicles. o The catalytic converter converts the carbon monoxide to carbon dioxide.

11.4 ChloroFluoroCarbons (CFCs) 

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The earth is protected by a layer of ozone which absorbs dangerous Ultra-Violet rays from the sun. The UV rays would otherwise cause severe damage to vegetations, as well as higher risk of skin o cancer. The ozone layer is constantly being destroyed by CFCs used in aerosols, refrigerators and cleaning solvents As a result of this depletion, use of CFCs is banned in many countries


11.5 Sulphur Dioxide and Sulphuric Acid 

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Sulphur Dioxide is a strong Reducing Agent It is a good bleaching agent (The reducing properties reduces the coloured dyes) o It is a good disinfectant (The reducing property effectively kill bacteria) o It is a good food preservative (The reducing property effectively kill bacteria) o It is used to manufacture Sulphuric Acid o Sulphuric Acid is a strong mineral acid It is used to make fertilizers o It is used to make detergents o It is used to make dyes o It is used as en electrolyte in car batteries o Although Sulphur Dioxide has many uses, it is also a serious pollutant when released to the environment o

As it is water soluble, it can dissolve to form acid rain 2SO2 (g) + O 2 (g) + 2H 2O(l) → 2H2SO4(aq)

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Acid rain is harmful to the environment as it Corrodes metal structures like bridges and vehicles Corrodes limestone buildings Endangers marine life as many aquatic life cannot survive in acidic waters Reduces the pH of soil, which causes plant life to easily die Sulphur Dioxide also irritates the eyes and cause breathing difficulties    

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The main source of Sulphur Dioxide is from the combustion of fossil fuels o Factories burning fossil fuels containing sulphur impurities o Petrol and diesel engines burning fossil fuels containing sulphur impurities Volcanos are also another secondary source of Sulphur Dioxide We can cut down the emission of Sulphur Dioxide by o Not using fossil fuels containing sulphur impurities o Spray factory exhaust chimneys with CaO or Ca(OH) 2 to absorb the Sulphur Dioxide CaO(s) + SO2 (g) → CaSO3 (s) Ca(OH)2 (s) + SO2 (g) → CaSO3 (s) + H2O(g) o

Solid calcium hydroxide or calcium oxides can be added to water bodies and soil to counter the effects of acid rain

11.6 Nitrogen Oxide and Nitrogen Dioxide 

At high temperatures, the nitrogen and oxygen in the air combine to form nitrogen oxide and/or nitrogen dioxide. These high temperatures can be due to lightning, forest fires, or at industrial factories and cars. o N2 (g) + O2 (g) → 2NO(g) 2NO(g) + O2 (g) → 2NO2 (g)

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These oxides are air pollutants as they can o Damage lungs o React with other air pollutants to form ozone, which irritates eyes and damage vegetation o

Dissolve in water to form acid rain


11.7 Summary of Air Pollutants Pollutant CO

Source - Cars internal combustion engine - Forest fires

Hazard It is a poisonous gas that combines with haemoglobin, reducing the efficiency to transport oxygen, leading to cell death.

Preventive measures - Install catalytic converters in cars - Reduce number of cars on road - Create efficient engines to ensure complete combustion

- Combustion of fossil fuels containing sulphur impurities - Volcanic eruptions

- Lung irritant - Eye irritant - Acid rain

- Prevent using fuels containing sulphur impurities, e.g. coal - Use hydrated CaO/ Hydrated Ca(OH)2 to absorb SO2 before it’s released into the atmosphere - Add CaO to soil and water bodies to neutralize acid rain

- Lightning activity - Forest fires - Internal combustion engines - Power stations

- Eutrophication - Lung damage - Acid rain

- Install catalytic converters in cars - Design car engines which run at lower temperatures - Add CaO to soil and water bodies to neutralize acid rain

Methane

- Decomposition of vegetable matter - Rice field - Cattle ranching - Natural gas - Mines

- Highly Flammable - greenhouse gas

- Cattle and other ruminant animals should be given improved diet - Animal manure and rotting vegetation can be used as biomass fuel

Unburnt hydrocarbons

- Internal combustion engines

- Carcinogenic - Forms photochemical smog - Greenhouse Gases contributing to global warming.

- Install catalytic converters in cars - Reduce number of cars on road - Create efficient engines in cars to ensure complete hydrocarbon combustion

Ozone

- It is formed when

- It reacts with unburnt

an electrical spark passes through air.

hydrocarbons to form photochemical smog that causes headache, eye, nose and throat irritation. - It corrodes and kills plants and trees

CFCs

- CFC based aerosol products

- Destroys the Ozone layer

Dust and Smoke

- Building work - Mining activities - Forest fires - Incomplete combustion of fuels.

- Irritate lungs, causing bronchitis and other lungrelated diseases.

Lead compounds

- Combustion of leaded petrol in car engines - lead compounds are added to petrol

- when breathed in can build up inside the body and are toxic and poisonous - Causes lead poisoning which leads to brain damage

Carbon Monoxide

SO2 Sulphur Dioxide

NO & NO2 Oxides of Nitrogen

- Use CFC-free products

to make heavier so that ititdoes not ignite too soon.


11.8 Treatment of Water 

The main steps in the treatment of raw water are: The water is first screened to remove large solids impurities Alum is added to cause fine suspended particles to clump together and settle in the sedimentation tank Lime is added to reduce acidity Addition of activated carbon to remove foul smells and taste o Filtration removes any remaining solid particles o Chlorination is carried out to disinfect the water by killing the harmful bacteria Fluoride is sometime added to prevent tooth decay o

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Raw water is screened

Mixing Chamber •

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Lime and Alum is added

Sedimentation Tank Activated Carbon is added

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Sand and Gravel filter beds

Clear water tanks •

Chlorine and fluoride is added

Clean water can also be prepared through desalination Desalination is the process of removing dissolved salts from seawater

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Two methods of desalination commonly used: Distillation: Seawater is evaporated and the pure water vapour formed is condensed o Reverse Osmosis: Pure water is extracted from seawater using a semi-permeable membrane under high pressure o

11.9 Control of pH in agriculture  

Most plants need a soil pH of 6.5 to 7.5 to grow well If the ground is too acidic, slaked lime - Ca(OH) 2 can be added to neutralise the acid. o This process is called liming the soil o Slaked lime is used as it is cheap and easily available o Slaked lime is sparingly soluble. Once the acid is neutralized, the excess base will remain as a solid in the soil and not dissolve to make the soil too alkaline.



Olevels Chemistry Notes - Combined Chemistry

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