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Chapter 7 – Covalent and Metallic Bonding

1. A covalent bonding is bond formed by sharing of electrons between atoms
2. A single covalent bond consists of one shared pair of electrons
3. A double covalent bond consists of two shared pair or electrons.
4. A molecule is formed when two or more atoms are joined together by covalent bonds.
5. In a molecule, each atom has the electronic configuration of a noble gas.

From : O Level Chemistry Tutor

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Chapter 6 – Ionic Bonding

1) Noble gases are unreactive and do not form compounds because they have octet structure
2) Metals form postively charged ions (Cations)
3) Non-Metals usually form neagtively charged ions (Anions)
4) Metals react with non-metals to form ionic compounds.
5) An ionic bond is formed when electrons are transfeered from a metallic atom to a non-metallic atom.

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Chapter 5 – Atomic Structure

1) Atoms are made up of protons, neutrons and electrons.
2) Proton number or atomic number is the number of protons in an atom.
3) Nucleon number is the number of protons and neutrons in an atom.
4) Isotopes are atoms of the same element with the same number of protons but different number of neutrons.

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Elements, Compounds and Mixtures – Key Points

1. All metals exist as atoms. Most non-metals as molecules, examples hydrogen and oxygen

2. Tap water is a mixture that contains water and dissolved substances such as chlorine and other minerals.

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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.

At standard state conditions,
Gθ = Hθ – TSθ

3. The standard electrode potential of an element is the potential difference between the element and its aqueous ion of 1.00 mol dm–3 relative to that of the standard hydrogen electrode at 1 atm and 298 K.

4. The standard cell potential, Ecell, is the potential difference between two standard half cells measured under standard conditions. When Ecell > 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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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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Experimental Techniques – Key Points

1. Fractional distillation is used to separate a mixture of liquids with different boiling points

2. The starting line in paper chromatography should be drawn with pencil and not ink. Ink contains dyes that will dissolve in the solvent.

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Experimental Techniques – Common Error

Wrong Concept
The accuracy of burette is +- 0.01

Correct Concepts
The accuracy of bureete is +-0.05.
Thus the burette reading can be 23.80 or 23.85 cm^3 but not 23.84cm^3

Wrong Concept
Lemon juice is a pure substance

Correct Concepts
It is a mixture of different substance such citric acid and water

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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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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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Kinetic Particle Theory – Common Error

Wrong Concept
During melting, the temperature rises because heat is absorbed

Correct Concept
During melting, the temperature remains constant at the melting point because the heat absorbed is used to overcome the attractive forces between the particles.

Wrong Concept
Diffusion does not occur in solid

Correct Concept
Diffusion does take place in solids, it occurs extremely slowly in solids because the solid particles are less energetic and move with less speed

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Purification of Substances

Stop and Think, are the following questions True or False

1. In filtration, the filtrate is always a pure liquid.

2. Drinking water can only be obtained from seawater by distillation.

3. The fractional distillation of miscible liquids is only possible if the liquids have different boiling points.

4. Paper chromatography is a physical method for separating mixutres.

5. Mixtures have fixed melting and boiling points.

Ans: FFTTF

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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 – 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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Stop and Think, are the following questions True or False

1. Temperature, volume and mass are physical quantities.

2. All measurements in chemistry are made using SI units.

3. It is not possible to measure precise volumes of liquids.

4. Clocks and watches are accurate to only 1 second.

5. Many physical quantities can be measured with senors connected to
computers.

Ans: TFFFT

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