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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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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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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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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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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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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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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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Definitions – Redox

1. Oxidation is a process where a chemical species loses electrons; (Oxidation Is Loss)

2. Reduction is a process where a chemical species gains electrons. (Reduction Is Gain)

3. A redox reaction refers to a reaction where oxidation and reduction occurs simultaneously.

4. An oxidising agent is a species that accepts / gains electrons (is reduced) n a reaction.

5. A reducing agent is a species that donates / loses electrons (is oxidised) in a reaction.

6. A disproportionation reaction is a redox reaction in which one species is simultaneously oxidised and reduced.

e.g. 2Cu+ —-> Cu + Cu2+

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Definitions – Atomic Structure

1. Atomic number of an element refers to the number of protons it contains. Mass number (nucleon number) refers to the sum of the protons and neutrons it contains.

2. Isotopes refer to atoms of the same element with the same number of protons but different number of neutrons.

3. An atomic orbital is defined as a region of three-dimensional space around the nucleus, whereby there is a 95% chance of locating a particular electron. Each orbital has a characteristic energy level and shape.

For ‘A’ level syllabus, you need to know the shapes of s and p orbitals.

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Definitions – Chemical Bonding

1. Valence-shell electron pair repulsion (VSEPR) theory is a model used to predict the shape of individual molecules based upon the extent of electron-pair electrostatic repulsion.

Lone pair – lone pair repulsion > Lone pair – bond pair repulsion > Bond pair – bond pair repulsion

2. Metallic bond is the electrostatic attraction between positively charged cations and the ‘sea’ of delocalised electrons.

3. Electrovalent (Ionic) bond is the electrostatic attraction between oppositely charged ions which have been formed by the transfer of one or more electrons to achieve the stable electronic configuration of a noble gas.

Coordination number of an ion in an ionic compound refers to the number of neighboring oppositely charged ions.

4. Covalent bond is the electrostatic force of attraction of the nuclei of the 2 atoms for the shared pair(s) of electrons between them.

5. Dative / Co-ordinate Covalent bond is a covalent bond in which a pair of electrons is shared between 2 atoms but ONLY ONE of them provides both electrons that make up the bond.

6. Electronegativity refers to the ability/tendency of an atom to attract electrons in a bond towards itself. Electronegativity increases across the period and decreases down a group.

7. Permanent dipole-permanent dipole interactions are a type of intermolecular forces between polar molecules (molecules with a net dipole moment) which have a simple covalent structure.

8. Instantaneous dipole-induced dipole interactions are a type of intermolecular forces between non-polar molecules (molecules with NO NET dipole moment) which have a simple covalent structure.

9. Hydrogen bonds are a special case of permanent dipole-permanent dipole interactions, whereby there is an attractive interaction of a hydrogen atom with an electronegative atom, such as nitrogen, oxygen or fluorine (typically from another molecule). Do not confuse this with a covalent bond between H and N, O or F.

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Gases

Basic Assumptions of kinetic theory of gases

• Gases consist of small particles of negligible size/volume as compared to the size of the container.
• Gas particles have negligible intermolecular forces of attraction between each other.
• Collisions between gas particles are perfectly elastic. I.e. there is no loss of kinetic energy upon collision.

Ideal Gas equation PV = nRT
Use Pa for Pressure, m3 for volume and K for temperature

Pressure 1 atm = 1.01 x 105 Pa
1 bar = 1 x 105 Pa
760 mmHg = 1.01 x 105 Pa

Volume 1 dm3 = 10-3 m3
1 cm3 = 10-6 m3

Temperature T(K) = T(C) + 273

At s.t.p Temperature = 273 K (0C)
Pressure = 1.01 X 105 Pa (1 atm)
Molar volume = 0.0224 m3

At r.t.p (standard conditions)
Temperature = 298 K (25C)
Pressure = 1.01 x 105 Pa (1 atm)
Molar volume = 0.024 m3

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Chemical Energetics – Defination

1. Hess’ law states that the change in enthalpy accompanying a reaction is independent of the path taken between the initial and final states.

2. The standard enthalpy change of reaction (Hrxn), is the enthalpy change when molar quantities of reactants (as specified by the chemical equation) react to form products under standard conditions 25C and 1 atm.

3. The standard enthalpy change of formation of a compound (Hf), is the enthalpy change when 1 mole of a pure compound in a specified state is formed from its constituent elements in their standard states, under standard conditions 25C and 1 atm.

4. The standard enthalpy change of combustion of a compound (Hc), is the enthalpy change when 1 mole of that compound is completely burnt in oxygen under standard conditions 25C and 1 atm.

5. The standard enthalpy change of neutralisation (Hneu), is the enthalpy change when an acid and a base react to form 1 mole of water under standard conditions 25C, and 1 atm.

6. The standard enthalpy change of atomisation of an element (Hatom ), is the enthalpy change when 1 mole of atoms in the gaseous state is formed from the element in its normal physical state under standard conditions 25C and 1 atm.

7. The bond dissociation energy of a bond is the energy required to break one mole of chemical bonds between two atoms in a molecule in the gaseous phase.

8. The first ionisation energy of an element (H1st I.E.), is the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms, to form 1 mole of gaseous singly charged cations. M(g)  M+(g) + e−

9. The second ionisation energy of an element (H2nd I.E.), is the energy required to remove 1 mole of electrons from 1 mole of gaseous singly charged cations, to form 1e mol of gaseous doubly charged cations. M+(g)  M2+(g) + e−

10. The first electron affinity of an element (H1st E.A), is the enthalpy change when 1 mol of electrons are added to 1 mol of gaseous atoms, to form 1 mol of gaseous singly charged anions.

11. Lattice energy is the energy evolved when 1 mole of an ionic solid is formed from its constituent gaseous ions under standard conditions 25C and 1 atm.

12. The standard enthalpy change of hydration of a gaseous ion (Hhyd), is the enthalpy change when 1 mole of hydrated aqueous ions is formed from the gaseous ions under standard conditions 25C and 1 atm. .

13. The standard enthalpy change of solution of an ionic compound (Hsoln), is the enthalpy change when 1 mole of an ionic compound is dissolved in a large excess of water under standard conditions 25C and 1 atm.

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Chemical Equilibrium – Concepts

1. Entropy (S) measures the degree of disorder in a system. The entropy of a system increases when the matter or energy in the system becomes more random in its arrangement. A system that has a high degree of disorder/randomness is said to have a large entropy. Gases have the highest entropy followed by liquids and solids.

2. The Gibbs Free Energy change, G, is the limiting maximum useful work that can be obtained from a reaction, at constant pressure. When G 0, the reaction is spontaneous.
Ecell = Ecathode – Eanode

5. Dynamic Equilibrium refers to a reversible reaction in which the forward and the backward reactions are taking place at the same rate and concentrations of reactants and product are constant.

6. Le Chatelier’s Principle states that if a system in equilibrium is subjected to a change which disturbs the equilibrium, the system will respond in such a manner as to reduce or counteract the effect of the change.

Industrial application:
Haber Process: 450 oC – 500 oC, 200 atm – 300 atm

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Ionic Equilibrium – Concepts

1.
Monoprotic or monobasic acids can donate only one proton. E.g. HCl, HNO3 and CH3COOH

2. Diprotic or dibasic acids can donate two protons. E.g. H2SO4, H2S and H2CO3

3. A Bronsted acid is a proton donor.

4. A strong acid is one that dissociates completely in aqueous solution to give H3O+ ions.
HA (aq) + H2O (l) —> H3O+ (aq) + A- (aq)

5. Weak acids only dissociate partially in aqueous solution forming ionic equilibrium systems
HA (aq) + H2O (l) —> H3O+ (aq) + A- (aq)

6. Ka provides an accurate measure of the extent to which a weak acid is dissociated.

Ka = [H3O+][A-]/[HA][H2O]

Ka is only affected by changes in temperature

7. A Bronsted base is a proton acceptor.

8. A strong base is one that dissociates completely in aqueous solution to give OH- ions.
B (aq) + H2O (l) —> BH+ (aq) + OH- (aq)

9. Weak bases only dissociate partially in aqueous solution forming ionic equilibrium systems.

B (aq) + H2O (l) —> BH+ (aq) + OH- (aq)

Kb = [BH+][OH-]/[B]

Kb is only affected by changes in temperature

10. pH is defined as the negative logarithm to base 10 of [H3O+]
 pH = – log10[H3O+]

pOH is thus the negative logarithm to base 10 of [OH-]
 pH = – log10[OH-]

pH + pOH = 14

11. Water ionizes itself to a very small extent to give H3O+ and OH- ions.

H2O (l) + H2O (l) —> H3O+ (aq) + OH- (aq) H = +ve
The equilibrium constant for the above system is given the symbol Kw and is known as the ionic product of water.

Kw = [H3O+][OH-] = 1.0 x 10-14 mol2dm-6 (at 25oC)

As the auto-ionisation of water is an endothermic process, when temperature is increased, equilibrium shifts to the right to absorb the heat. [H3O+] and [OH-] increase by the same amount, Kw increases.

Kw is only affected by change in temperature

12. An acidic/alkaline buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a mixture of a weak base and its conjugate acid. It has the property that it resists changes in pH when a small amount of acid or base is added to it.

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Reaction Kinetics – Concepts

1. Types of rate:
• Initial rate is change in concentration of reactants or product at time t = 0.
• Instantaneous rate is rate of reaction at any given time/instant.
• Average rate is total concentration of reactant used or total concentration of product formed over total time.

2. The minimum energy which colliding molecules must possess for successful collision/reaction is called the activation energy, Ea.

3. Rate law or rate equation is the mathematical relationship between the rate of a reaction and the concentration of the reactants in a reaction. E.g. A + B —> Products
The rate law is
Rate = k[A]^m[B]^n
where k is the rate constant
m is the order of reaction with respect to reactant A
n is the order of reaction with respect to reactant B
(m + n) is the overall order of the reaction
4. The order of reaction with respect to a particular reactant is the power to which the concentration of that reactant is raised in an experimentally determined rate equation / rate law.

5. The rate constant, k, is the proportionality constant in the experimentally determined rate law.

6. The half-life (t1/2 ) of a reaction is the time taken for the concentration of a reactant to fall to half its initial value. It is constant only for a first order reaction as it is independent of reactant concentrations.

7. A catalyst is a substance that increases the rate of a reaction by providing an alternative reaction pathway that has lower activation energy.

8. Biological catalysts such as enzymes are very selective in the reactions that they catalyze, and some are absolutely specific, operating for only one substance in only one reaction. For reactions that normally produce a pair of optical isomers (racemic mixture) when carried out in the lab, enzymes are able to selectively produce one optical isomer in the body.

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