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Chapter 18 – Organic

Macromolecules

Macromolecules (polymers) are formed by linking together many small repeating units known as monomers.

Polymerisation is the process of joining together a large number of small molecules to form a macromolecule.

2 Classes of macromolecules

Synthetic macromolecule
Man-made polymers
Eg. Polyethene, nylon, terylene

Natural macromolecule
Naturally occurring
Eg. Proteins, Fats, Carbohydrate, Cellulose, Wool, Cotton etc…

2 Types of Synthetic Polymerisation

Addition Polymerisation
is a process by which many small unsaturated molecules (monomers) are added onto one another to form one large molecule (polymer).

Condensation Polymerisation
is a process where two monomers react together to produce a large molecule, with the elimination of a small molecule (ie. Water or HCl)

Addition Polymerisation

For ONLY Unsaturated monomers
Conditions : Heat, High pressure, Catalyst
Examples:
Ethene -> Poly(ethene)
Propene -> Poly(propene

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Definitions – AMS

1. The relative atomic mass (Ar) of an element is defined as the average mass of one atom compared to 1/12 the mass of a 12C atom.

OR
The relative atomic mass (Ar) of an element is defined as the mass of one mole of atoms compared to 1/12 the mass of one mole of 12C atoms.

2. The relative isotopic mass of an isotope (of a particular element) is defined as the mass of one isotope compared to 1/12 the mass of a 12C atom.

3. The relative molecular mass of a molecule is defined as the average mass of one molecule compared to 1/12 the mass of a 12C atom.

4. The relative formula mass of an ionic compound is defined as the average mass of one formula unit compared to 1/12 the mass of a 12C atom.

OR
The relative formula mass of an ionic compound is defined as the mass of one mole of formula units compared to 1/12 the mass of one mole of 12C atoms.

5. A mole of substance is defined as the amount of substance that contains as many entities (atoms, molecules, ions, electrons or any other particles) as the number of atoms in 12g of the carbon-12.

It is equal to 6.022 X 10^23, which is called the Avogadro constant or Avogadro number.

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Chapter 18 – Atmosphere and Environment

Greenhouse Gases

Carbon dioxide is released to the atmosphere when solid waste, fossil fuels (oil, natural gas, and coal), and wood and  wood products are burned.

Methane is emitted during the production and transport of coal, natural gas, and oil. Methane emissions also result from the decomposition of organic wastes in municipal solid waste landfills, and the raising of livestock.

Nitrous oxide is emitted during agricultural and industrial activities, as well as during combustion of solid waste and fossil fuels.

What are the effects of global warming?
1. Heat waves and periods of unusually warm weather
2. Sea level rise and coastal flooding
3. Glaciers melting
4. Arctic and Antarctic warming
5. Spreading disease
6. Earlier spring arrival
7. Plant and animal range shifts and population declines
8. Coral reef bleaching
9. Downpours, heavy snowfalls, and flooding
10. Droughts and fires

What is causing the depletion of ozone?
1. Chlorofluorocarbon compounds are also known as CFCs – made of  chlorine, fluorine and carbon.
2. These compounds are unreactive and do not burn.
3. They are compressed to form liquids which are used in aerosol  propellants and as coolant fluids for refrigerators and air conditioners.
4. When an aerosol can is utilised, the CFC molecules are released into  the air. At higher altitudes in the atmosphere, these molecules are  decomposed by sunlight to produce chlorine atoms.
5. The chlorine atoms react with the ozone molecules and thus destroy  the ozone layer which protects the earth from the direct rays of the  sun.
6. Exposure to direct radiation of the sun can cause skin cancer to  humans and also destroy the agriculture.

Effects of Ozone Depletion
1. There is an increase in the temperature due to the ozone depletion as more UV rays are entering the earth’s surface.
2. There would be a no.of skin cancers.3.Crop yields would be adversely affected.Vegetation land becomes dessert.
4. North and South Pole melts causing the ocean level to rise and flood the low-lying countries such as Netherlands.5.Rapid evaporation would occur and cause droughts especially in India.Thus, CO2 dissolved in the oceans rise into the atmosphere adding further to the greenhouse effect.

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Chemical Energetics – Definition

1. The enthalpy change of formation, Hf , of a compound is defined as the enthalpy change when 1 mole of the compound is formed from its elements under standard conditions of 298 K and 1 atm.

2. The standard enthalpy change of combustion, Hco , is defined as the enthalpy change when one mole of a compound is completely burnt in oxygen under standard conditions of 298K and 1 atm.

3. The standard enthalpy change of hydration, Hhydo, of an ion is defined as the enthalpy change when 1 mole of the gaseous ions is dissolved in a large amount of water under standard conditions of 298 K and 1 atm

4. The standard enthalpy change of solution, Hsolo , is defined as the enthalpy change when 1 mole of a substance dissolves in such a large volume of solvent that addition of more solvent produces no further heat change under standard conditions of 298 K and 1 atm.

5. The standard enthalpy change of neutralisation, Hno, is defined as the enthalpy change when 1 mole of water is formed in the neutralisation between an acid and an alkali, the reaction being carried out in aqueous solution under standard conditions of 298 K and 1 atm.
Always negative (exothermic reaction)

6. The standard enthalpy change of atomisation, Hato, is defined as the enthalpy change when 1 mole of separate gaseous atoms of the element is formed from the element under standard conditions of 298 K and 1 atm.

7. Bond energy is the energy required to break the covalent bond between 2 atoms in the gaseous state (for dissociation).It is usually measured in kJ mol-1 and is an average value.

8. Ionisation energy is the energy required to remove one electron from each atom in a mole of gaseous atoms producing one mole of gaseous cations

9. The electron affinity is the energy change associated with the formation of an anion from the gaseous atom, measured in kJ mol-1.

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Chemical Bonding FAQ 3

Do you know?

The fact that ice is less dense than water causes water to freeze downwards. This helps living organisms in a body of water to survive freezing conditions!

As the temperature of water near the surface drops, the density of water increases. Cold water sinks while warmer water, which is less dense, rises. This convection motion of water continues until the temperature of the water reaches 4 °C. Below this temperature, the density of water
decreases with decreasing temperature so that water colder than 4 °C no longer sinks. On further cooling, the water begins to freeze at the surface. The ice layer formed does not sink as it is less dense than water.

This layer of ice helps to insulate the water underneath from further heat loss, thus keeping the water below it from freezing solid. Hence, living things can survive in ponds and rivers even when the temperature falls below freezing.

Question 1: Explain why the boiling point of propane, dimethyl ether and ethanol deviates so greatly even though their electron cloud sizes are similar.

Since the electron cloud sizes of the three compounds are similar, the magnitudes of instantaneous dipole–induced dipole (id–id) interactions in the three compounds are similar.

Propane is non–polar so that the intermolecular forces between propane molecules are due only to id–id interactions.

Dimethyl ether is polar so that besides id–id interactions, permanent dipole–permanent dipole (pd–pd) interactions also exist between the molecules. Since boiling involves breaking intermolecular forces, more energy is required to overcome the stronger pd–pd interactions between dimethyl ether molecules compared to the weaker id–id interactions between propane molecules. Hence the boiling point of dimethyl ether is significantly higher than propane.

Ethanol contains a hydrogen atom covalently bonded to the small and highly
electronegative O atom so that hydrogen bonding exists between ethanol molecules. Since ethanol contains hydrogen bonding besides pd–pd interactions and id–id interactions, intermolecular forces between ethanol molecules are the strongest,hence the energy required to overcome intermolecular interactions in ethanol is greater than in propane and dimethylether. Hence ethanol has a much greater boiling point than both propane and dimethylether.

Question 2: Explain why iodine has a higher boiling point than water even though iodine has only instantaneous dipole–induced dipole interactions and water can form hydrogen bonds.

The strength of instantaneous dipole–induced dipole interactions depends on the size of the electron cloud. Iodine has a much larger electron cloud than water so that the instantaneous dipole–induced dipole interactions between iodine molecules are significantly stronger than the hydrogen bonding between water molecules which have much smaller number of electrons.

Since boiling involves overcoming the intermolecular attractions between molecules, a greater amount of energy is required to overcome the id−id forces between iodine molecules compared to the id−id forces and the hydrogen bonds between water molecules.

FAQ: Explain why ionic compounds do not dissolve in non-polar solvents.

Answer: The ion−solvent interaction is much too weak to overcome the
strong electrostatic attraction between the ions in the crystal lattice.

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Chapter 17 – Speed of Reaction

1. Different chemical reactions take place at different speeds

2. Speed of reaction = change in amount of reactant or product/Time taken

3. Speed of reaction = change in volume gas/Time taken

4. For reactions involving solutions, an increase in the concentration of a reactant increases the speed of reaction.

5. For reactions involving gases, an increase in the pressure of a gas increases the speed of reaction.

5. Reactions take place faster when the solid is broken into smaller pieces.

6. The higher the temperature, the faster the movement of the particles and the greater the number of collisions. Hence, reactions take place faster when the temperature is increased.

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Chemical Bonding FAQ 2

FAQ: Why does a stream of chloromethane (polar substance) running from a burette be deflected towards a negatively charged rod?

The polar molecules will align themselves such that the d+ ends of the molecules will face the negatively charged rod. The electrostatic force of
attraction causes the stream to be deflected towards the rod.

FAQ: Explain the trend in the boiling points of the halogens.

The halogens from F2 to I2 are non-polar so that the intermolecular attraction between their molecules is due to instantaneous dipole–induced dipole (id−id)interactions.

The number of electrons of the halogens increases down the group from F2 to I2.\ Ease of polarisability of electron clouds and strength of id–id interactions also increase down the group, with I2 having electron clouds which are most easily polarised resulting in greatest id–id interactions.
Since boiling involves overcoming intermolecular forces, the boiling points of halogens increases down the group.

Ease of donation of lone pair of electrons on Y
FAQ Why must Y be N, O or F?

To interact strongly with the d+ H on H-X, Y needs to be highly electronegative and the lone pair of electrons on Y must not be too diffuse in space. Hence, Y also needs to be small
must be N, O or F.

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Chapter 15 – Electrolysis

1. An electrolyte is a compound that conducts electricity in the molten state or in aqueous state.

2. A cathode is a negatively charged electrode. An anode is a positively charged electrode.

3. Positively charged ions are called cations and negatively charged ions are called anions.

4. The decomposition of a compound by electricity is called electrolysis.

5. During the electrolysis, cations move towards the cathode while the anions move towards the anode.

6. Redox reactions take place at the electrodes. Oxidation occurs at the anode and reduction at the cathode

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Chapter 15 – Electrolysis

1. An electrolyte is a compound that conducts electricity in the molten state or in aqueous state.

2. A cathode is a negatively charged electrode. An anode is a positively charged electrode.

3. Positively charged ions are called cations and negatively charged ions are called anions.

4. The decomposition of a compound by electricity is called electrolysis.

5. During the electrolysis, cations move towards the cathode while the anions move towards the anode.

6. Redox reactions take place at the electrodes. Oxidation occurs at the anode and reduction at the cathode

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Chemical Bonding FAQ 1

What are Chemical bonds?
– Binding forces of attraction between particles (atoms, ions or molecules) resulting in a lower energy arrangement.
– The formation of a bond involves the re-distribution of the outer electrons of the atoms concerned.

What is the Octet Rule?
Atoms tend to lose, gain or share electrons until they are surrounded by eight valence electrons. Atoms try to achieve the same number of electrons as the noble gases closest to them in the Periodic Table.

Explain which of the following ionic compound has the stronger ionic bond.
(i) sodium fluoride or sodium chloride
Sodium fluoride has stronger ionic bond as fluoride is smaller than chloride so that the shorter distance between Na+ and F– resulted in stronger ionic bond.

(ii) sodium fluoride or magnesium fluoride
MgF2 has stronger ionic bond as Mg2+ is more highly charged than Na+, resulting in greater electrostatic attraction in MgF2.

Why is it that both NCl3 and PCl3 exist, but only PCl5 exist and not NCl5?
Such expansion of octet is observed in some compounds formed by elements of Period 3 (and beyond) This is due to the availability of vacant, low-lying orbitals. The energy required to promote an electron from 3s or 3p to 3d is not very large.

However elements in Period 2 (e.g. O and N) do not have low-lying vacant orbitals for expansion of octet. Promotion of electrons to the next quantum shell requires too much energy and hence Period 2 elements can accommodate only a maximum of eight valence electrons.

Point to note
A single covalent bond consists of one s bond.
• A double covalent bond consists of one s bond and one pie bond.
• A triple covalent bond consists of one s bond and two pie bonds.

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Apr Lesson Plan – J2

Nitrogen Compounds – Lecture Outline

1. Amines
1.1 Introduction and Classification of Amines
1.2 Nomenclature
1.3 Physical properties of amines
1.3.1 Physical Properties of aliphatic amines
1.3.2 Physical Properties of Phenylamine
1.4 Basicity of Amines
1.4.1 Relative Basicity of Ammonia
1.5 Preparation of amines
1.5.1 Reduction of Nitrile Compound
1.5.2 Nucleophilic Substitution of Halogenoalkane
1.5.3 Reduction of nitrobenzene
1.6 Reactions of Amines
1.6.1 Reaction of amine as a base
1.6.2 Acylation of amines
1.6.3 Reaction of Phenylamines with Bromine

2. Amides
2.1 General properties
2.2 Preparation of Amides
2.3 Reactions of Amides
2.3.1 Hydrolysis of Amides
2.4 Chemical test: Distinguishing Amides, Ammonium salts and Amines

3. Amino Acids
3.1 General Properties of amino acids
3.2 Physical properties
3.3 Separation of Amino Acids
3.4 Peptide Formation
3.5 Hydrolysis of Proteins or Polypeptides

4. Proteins
4.1 Introduction
4.2 Classification of Amino acids
4.3 Types of R group interactions of Amino-acids
4.4 Peptides Formation
4.4.1 Formation of Peptide Bond
4.4.2 Hydrolysis of Peptide Bond
4.5 Structure of Proteins
4.5.1 Primary Structure of Proteins
4.5.2 Secondary Structure of Proteins
4.5.3 Tertiary Structure of Proteins
4.5.4 Quaternary Structures of Proteins
4.6 Denaturation of proteins

Electrochemistry Part 1 – Lecture Outline

1. Electrolytic cell
1.1 Definition
1.2 Set-up of electrolytic cell

2. Electrolysis of Compound
2.1 Electrolysis of Molten ionic Compound
2.2 Electrolysis of Aqueous Solution

3. Preferential Discharge: Factors affecting the discharge of ions
3.1 Position of ion in Redox Series (Electrode Potential)
3.2 Concentration of Ions
3.3 Nature of Electrodes

4. Faraday’s Law of Electrolysis

5. Calculation using Faraday’s Law

6. Industrial Application of Electrolysis
6.1 Electrolysis of brine (Saturated NaCl) using a diaphragm cell
6.2 Anodising of Aluminium
6.3 Electrolytic purification of Copper
6.4 Electroplating

Electrochemistry Part 2 Electrochemical Cell – Lecture Outline

1. Electrochemical Cell
1.1 Set-up
1.2 Cell Diagram / Cell Notation

2. Electrode Potential
2.1 Definition of Electrode Potential
2.2 Factors affecting Electrode Potential

3. Standard Electrode Potential
3.1 Definition of standard electrode potential:
3.2 Standard Hydrogen Electrode (S.H.E)
3.3 Measuring standard electrode potential, Eq
3.3.1. To determine Eq of metal – metal ion half-cell
3.3.2. To determine the of Eq non-metal (gaseous) – non-metal ion half
cell
3.3.3 To determine the of Eq of ion – ion half-cell
(ions of the same element in different oxidation states)

4. Standard Cell Potential (Eq cell or cell emf)

5. Redox Series (Electrode potential)

6. Application of REDOX Series (Electrode potential)
6.1. Determine the emf of Cell
6.2 Predict the reactivity of elements
6.3 Determine the strength of oxidizing and reducing reagents
6.4 Predict the relative stabilities of metallic ions in different oxidation states
6.5 Predict the feasibility of Redox Reactions

7. Limitation of Standard Electrode Potential

8. Batteries and Fuel Cell
8.1 Two main types of batteries
8.2 Fuel Cell

9. Comparing Electrochemical Cell and Electrolytic Cell

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Apr Lesson Plan – J1

Chemical Bonding – Lecture Outline

1. Introduction
2. Ionic Bonds
3. Metallic Bonds
4. Covalent Bonds
5. Shapes of Molecules
6. Partial Ionic and Partial Covalent Character
7. Polar and Non-polar Molecules
8. Intermolecular Forces of Attraction

Gases – Lecture Outline

1 Introduction to Gases

2 The Gas Laws
2.1 Boyle’s Law
2.2 Charles’s Law
2.3 Combined Gas Law

3 Ideal Gas Law

4 Avogadro’s Law

5 Dalton’s Law of Partial Pressure

6 Kinetics Theory of Gases

Chemical Kinetics – Lecture Outline

1 Rate of Reaction

2 Rate Equations & Orders of Reaction
2.1 The Rate Law
2.2 Zero-Order Reactions
2.3 First-Order Reactions
2.4 Second-Order Reactions

3 Experiments for Studying Kinetics
3.1 Monitoring Concentration Changes
3.2 Deducing Order of Reaction
(a) Initial Rate method
(b) Inspection Method and Calculation Method

4 Reaction Mechanisms
4.1 Elementary & Non-Elementary Reactions
4.2 Reaction Mechanisms

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Chapter 14 – Metal

1. The physical properties of metals are as follows:
– Usually have high densities, melting points and boiling points.
– Can be bent, stretched or beaten into very thin sheets without breaking
– Good conductors of heat and electricity

2. Any alloy is a mixture of a metal with one or few other elements

3. There are four main reasons for making alloys
– To improve the strength and hardness of metals.
– To improve the appearance of metals.
– To improve the resistance of metals against corrosion.
– To lower the melting points of metals.

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Hi JC/A level/O level Chemistry Tuition Students

Chapter 13 – Salts

1. Salts can be prepared using the following methods:

Solubility of the salt in water – soluble
Solubility of the starting materials in water – one is insoluble
Method of preparation – reaction of aids, with metals, insoluble bases or insoluble carbonates
_________________________________________________________________________

Solubility of the salt in water – soluble
Solubility of the starting materials in water – soluble
Method of preparation – titration of solids with alkalis or soluble carbonates
__________________________________________________________________________

Solubility of the salt in water – insoluble
Solubility of the starting materials in water – soluble
Method of preparation – precipitation
__________________________________________________________________________

2. Filtration and crystallisation are important laboratory techniques used for the separation and purification of salt crystals

Contact Mr Ong @9863 9633 for much key ideas on salts

 

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Chemical Bonding – Part 1 FAQ

What are Chemical bonds?

– Binding forces of attraction between particles (atoms, ions or molecules) resulting in a lower energy arrangement.
– The formation of a bond involves the re-distribution of the outer electrons of the atoms concerned.

What is the Octet Rule?

Atoms tend to lose, gain or share electrons until they are surrounded by eight valence
electrons. Atoms try to achieve the same number of electrons as the noble gases closest to them in the Periodic Table.

All noble gases (with the exception of helium) have eight valence electrons. They have very stable electronic arrangements. Evidence of stability of noble gases:
– high ionisation energy
– low affinity for additional electrons
– general lack of reactivity

However, there are many exceptions to the octet rule. Nevertheless, it provides a useful framework for introducing many important concepts of bonding.

Why is it that both NCl3 and PCl3 exist, but only PCl5 exist and not NCl5?

Such expansion of octet is observed in some compounds formed by elements of Period 3 (and beyond) This is due to the availability of vacant, low-lying orbitals. The energy required to promote an electron from 3s or 3p to 3d is not very large.

However elements in Period 2 (e.g. O and N) do not have low-lying vacant orbitals for expansion of octet. Promotion of electrons to the next quantum shell requires too much energy and hence Period 2 elements can accommodate only a maximum of eight valence electrons.

Explain which bond is stronger, C—H or Si—H.

C—H bond is stronger since C is smaller than Si so that valence
orbital of C is less diffuse and overlap of its valence orbital with that
of H is more effective.

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