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FAQ – Group VII

For Group VII halogens (e.g. Cl2, Br2, I2)

1. Volatility decreases down the group (i.e. boiling point increases)

Explanation:
• Increase in electron cloud size hence more easily polarised
• leads to stronger dispersion forces between molecules

2. Oxidising power decreases down the group

X2 + 2e ⇔ 2X−

Explanation:
• Decrease in effective nuclear charge, hence electron affinity decreases
• Hence less easily accepts electrons (i.e. less easily reduced)
• Eθ value decreases (Quote from Data Booklet)

Two important proofs of this are:
• Any halide (e.g. I−) can be displaced by the halogen (e.g. Br2) above it.
• Reaction with sodium thiosulfate (explain change in oxidation state of S)

3. Reactivity with H2 decreases down the group
X2 + H2 → 2HX

Explanation:
• Atomic size of X increases
• Hence H – X bond becomes longer and weaker
• Product formed is less and less stable, hence reactivity decreases.

Note: Quote bond energy values to explain this, NOT Eθ values

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Salt Preparation

Preparation of salt by reaction between acid and an insoluble base.

Example: Preparation of copper(II) sulfate from copper(II) oxide(insoluble base)

Step 1: Put sulfuric acid in a beaker.

Step 2: Add copper(II) oxide until no more can dissolve.
Copper(II) sulfate solution & unreacted copper(II) oxide

Step 3: Filter to remove the excess copper(II) oxide

Step 4: Collect the filtrate.

Step 5: Heat up the filtrate to evaporate the water to about 1/3 the original volume.

Step 6; Allow the filtrate to cool to room temperature so that copper(II) sulfate crystals.

Step 7: Wash the crystal with a little cold distilled water.

Step 8: Dry the crystal with filter paper

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Acids and Bases – Key Points

1. An acid is a substance that produces hydrogen ions, H+, when dissolved in water.

2. The strength of an acid refers to the extent to which the acid molecules dissociates when dissolved in water.

3. A base is a substance that reacts with an acid to form a salt and water only.

4. An alkali is a base that is soluble in water.

5. Neutralisation reaction is the reaction between an acid and a base to from a salt and water only.

6. The term concentration tells us how much a substance is dissolved in 1 dm^3

7. The term strength refers to how an acid or an alkali dissociates when dissolved in water.

8. The pH scale is a set is a set of numbers used to whether a solution is acidic, neutral or alkaline.

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FAQ – Group II

For Gp II metals (e.g. magnesium, calcium, barium)

1. Reactivity (with water or other substances) increases down the group

Explanation:
• Increase in shielding effect outweighs increase in nuclear charge down the group due to increase
in number of electron shells.
• Hence ionization energy decreases
• Metals are able to form ions more easily.

For Gp II ionic compounds (e.g. carbonates, nitrates and hydroxides),

2. Melting point decreases down the group

Explanation:
• Increase in cation size
• leads to decrease in magnitude of lattice energy.
• Hence, ionic bonds are weaker and more easily broken.

3. Thermal decomposition temperature increases down the group

Explanation:
• Increase in cation size,
• leads to decrease in charge density of cation.
• Hence, electron cloud of anion is less distorted and hence less easily decomposed.

Note: Quote ionic radius values from Data Booklet to explain this. Calculate charge density if necessary.

4. Solubility of sulfates decreases down the group (not in syllabus but
may still be tested)

Explanation:
• Increase in cation size
• leads to significant decrease in magnitude of hydration energy (as compared to slight decrease in magnitude of lattice energy)
• Hence, enthalpy of solution (= LE – ΔHhyd) is less exothermic.

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TO Master to PERFECTION before A’levels (Part 1)

Standard Definitions (Don’t Memorize. But appreciate and understand why key terms are important)

– Relative atomic, isotopic, molecular and formula mass, based on the 12C scale (just give mathematical expression)

– Mole in terms of the Avogadro constant
– VSEPR (2 assumptions)
– Basic assumptions of the kinetic theory as applied to an ideal gas
– Standard enthalpies (11 of them)
– Hess’ Law
– Entropy
– Standard electrode potential and standard cell potential
– Dynamic Equilibrium, LCP
– Strong and weak acids and bases
– Kc, KP, Ka, Kb, Kw, KSP,pH etc. (m. expression)
– Rate of reaction; rate equation; order of reaction; rate constant; (m. expression)
– Half life of a reaction
– Rate-determining step
– Activation energy
– Catalysts
– Transition metal, ligands, complex, coordination number
– Proteins 1o,2o,3o structure, Denaturation

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Organic Chemistry – Concepts

1. Empirical formula is the simplest formula that shows the ratio of each kind of atom in a molecule. e.g. C2H5 is the empirical formula for C4H10

2. Molecular formula shows the actual number of each kind of atoms in a molecule. e.g. C4H10

3. Structural formula shows how the atoms are connected to each other in a molecule. e.g. CH3CH2CH2CH3

4. Displayed/full formula shows all the bonds and relative placing of all the atoms in a molecule.

5. Homologous series are compounds have the same general formula and functional group and each homologue differs from its neighbor by a fixed group of atoms (e.g.–CH2). As we go down a homologous series, the chemical properties remain unchanged but there is a gradual change in physical properties. Examples of homologous series are alkanes, alkenes, alcohols…..

6. Structural isomerism refers to compounds with the same molecular formula but different structural formula. E.g. CH3COOCH3 and C2H5COOH

7. Stereoisomerism refers to compounds that have the same molecular formula but with different spatial arrangements.

• Geometric isomers have same carbon skeleton with double bonds restricting free rotation. For geometric isomerism to exist, there must be two different groups of atoms bonded to each side of the C=C bond.

• Optical isomers are non-superimposable mirror images of each other (enantiomers). Isomers have at least one chiral C atom, i.e. there are four different groups attached and have no plane of symmetry. An equal proportion of enantiomers forms a racemic mixture which is optically inactive.

8. The primary structure of a protein shows the exact order (or unique sequence) of the -amino acids held by peptide/amide linkages along the polypeptide chain. The primary structure determines what the protein is, how it folds and its function.

9. The secondary structure refers to the detailed configurations of the polypeptide chain. In a protein molecule, the long chain of amino acid units may be coiled into an -helix or folded into a -pleated sheet. Both structures are stabilized by hydrogen bonds between the N-H group of one amino acid residue and the C=O group of another along the main chain.

10. The tertiary structure of the protein refers to the overall 3-dimensional shape of the entire protein involving folding or coiling of the chains. It shows how protein molecules are arranged in relation to each other.

There are four types of R group interactions which hold the tertiary structure in its shape.

 van der Waals’ forces (induced dipole-induced dipole bonding) exist when non-polar R groups (e.g. alkyl or aryl groups) come close together. They are usually found on the inside of globular proteins where, because they are hydrophobic, they do not interfere with solubility.

 hydrogen bonding between polar groups (e.g.. –CH2OH, -COOH and –NH2 groups).

 ionic bonding eg. –COO-, -NH3+, and >NH2+.

 disulfide linkages eg. –SH or –CH2-S-S-CH2- groups.
Quaternary structure of proteins refers to the spatial arrangement of its protein subunits. It shows how the individually folded protein subunits are packed together to yield large structures. This only applies to proteins that contain two or more polypeptide chains. The individual polypeptide chains are called the subunits. E.g. haemoglobin contains 4 subunits, each containing a haem group.

11. It is stabilized by the same R-group interactions that stabilise the tertiary structure.

12. Denaturation is the loss of biological activity of a native protein. When proteins are denatured, the secondary and tertiary structures are disrupted i.e. the R group interactions are broken or destroyed. Note that the primary structure remains unaffected.
Factors that can lead to denaturation include extremes in pH, temperature, ionic salts, heavy metal compounds, presence of organic solvents etc.

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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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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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Chemical Calculations – Key Points

1. Spectator ions are ions that are not involved in a chemical reaction.

2. Stoichiometetry of the reaction is the relationship between the amounts of reactants and products involved in a chemical reaction.

3. Limiting reactant is the reactant that is completely used up in a reaction and determines the amount of products formed.

4. The concentration of a solution is given by the amount of a solute dissolved in a unit volume of the solution.

5. Molar concentration is the concentration of a solution expressed in mold/dm^3

6. The theoretical yield is the calculated amount of products that would be obtained if the reaction is completed.

7. Actual yield is the amount of pure products that is actually prodcued in the experiment.

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The Mole – Key Points

1. Both relative atomic mass and relative molecular mass have no units.

2. Empirical formula an molecular formula may or may not be the same.

3. The total percentage composition of the elements in the compound must be 100%.

4. 1 dm^3 = 1000 cm^3

5. When calculating the Mr of a substance in the reaction, do not include the coefficient.

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Chapter 12 – Acids and Bases

Q: What is an acid?
A: An acid is a substance which produces hydrogen ions, H+, when it is dissolved in water.

Q: What are the physical properties of acids?
A: 1. Acids have a sour taste.
2. Acids dissolve in water to form solutions which conduct electricity.
3. Acids turn blue litmus paper red.

Q: What are the chemical properties of acids?
A: 1. Acids react with reactive metals to form hydrogen gas and a salt.
metal + acid —–> salt + hydrogen

2. Acids react with carbonates to form a salt, carbon dioxide and water.
carbonate + acid —–> salt + water + carbon dioxide

3. Acids react with metal oxides and hydroxides to form a salt and water only.
metal oxide + acid —–> salt + water

metal hydroxide + acid  salt + water

Q: Do all metals react with acid?
A: No, When unreactive metals such as copper or silver are added to dilute acids, there is no reaction.

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Chapter 11 – Salt

Key ideas

1. Soluble salts are prepared by the following methods:
a) Acid + a metal (excluding potassium, sodium, calcium, copper and silver)
b) Acid + an insoluble base
c) Acid + an insoluble carbonate
d) Acid + an alkali (titration method)

2. Insoluble salts are prepared by the precipitation reaction of two soluble salt solutions

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Chapter 10 – Chemical Calculations

What does an equation tell us?

A balanced chemical equation shows important facts about a reaction
a) The reactants
b) The products
c) The ratio of the amounts (in moles) of the reactants and the products
d) The state of each reactants/products if indicated

It is the relationship between the amounts (measured in moles) of reactants and products involved in a chemical reaction.

Exam based questions will be discussed in the lessons

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Chapter 9 – The Mole

1. A mole of any substance contains 6 x 10 power 23 particles. This number is called Avogadro’s constant

2. Number of mole of atoms = mass of the element(g)/Ar

3. Number of mole of substance = Mass of the substance(g)/Mr

4. Molar mass refer to the mass of one mole of the substance.

From : O Level Chemistry Tutor

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Chapter 8 – Writing Chemical Equations

3 Basic steps

Step 1 : Write down the chemical formulae of the reactants and products to get the chemical equation.

Step 2 : Check the number of atoms of each element in the formulae on both sides of the equation are balanced.

Step 3 : Add the state symbols

From : O Level Chemistry Tutor

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