Chemical Equilibria I

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1)Reversible reaction
A reversible reaction is a chemical reaction which can take place in both directions, i.e. reactant react to form products and products can also react to form the original reactants.

2)Dynamic equilibrium
A system is said to be in dynamic equilibrium occurs when the rate of forward reaction is equal to the rate of reverse reaction and the concentration of products and reactants remain constant. It is dynamic because reactants are being continuously converted to products and products are being continuously converted back to reactants.

3)Le Chatelier's Principle
Le Chatelier's Principle states that in a system in equilibrium, when a change is made to some external factor, the position of equilibrium shifts to oppose the change.

I will use the production of ammonia as an example.

4)Effects of changes in concentration
What happens when you increase the concentration of nitrogen gas. According to Le Chatelier's Principle the equilibrium reacts to remove the extra nitrogen gas that is added. Thus the equilibrium shifts to the right and produces more ammonia.

5)Effects of changes in pressure
Changes in pressure only affects reactions involving gases. From the chemical equation, we can see that 1 mol of nitrogen gas reacts with 3 mol of hydrogen gas to form 2 mol of ammonia gas, i.e. 4 mol of gaseous reactants react to form 2 mol of gaseous products. What happens when the pressure of the system is increased? According to Le Chatelier's Principle, the equilibrium shifts to the right to reduce the pressure.

6)Effects of changes in temperature
The formation of ammonia from nitrogen and hydrogen is an exothermic reaction, i.e. the forward reaction is exothermic . What happens when the temperature is increased? Based on Le Chatelier's Principle, the equilibrium will react to oppose the increase in temperature and the equilibrium shifts to the left.

7)Effects of a catalyst
A catalyst does not change the position of the equilibrium. It only increases the rate at which the system reaches equilibrium.

8)Equilibrium Constant Kc

If a reversible reaction is allowed to reach equilibrium, the product of the concentration of products divided by the product of the concentrations of reactants has a constant value at a particular temperature.

Changes in concentration and pressure and the presence of a catalyst does not change the equilibrium constant. Only temperature change will cause a change in the equilibrium constant. For an exothermic reaction, the equilibrium constant decreases as temperature is increased . For an endothermic reaction, the equilibrium constant increase with increasing temperature.

To be continued
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The Periodic Table: Chemical periodicity II

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7) Variation of oxidation number in chlorides and oxides

The oxidation numbers of elements in their oxides is always positive because oxygen is a very electronegative atom. The maximum oxidation number of each element is the same as its group number. The group number of the element corresponds to the number of electrons in its outermost shell.

Most of the elements have their usual oxidation states. Just take note that phosphorous has an oxidation state of +5 in phosphorous oxide and +5 in phosphorous chloride. Sulphur has an oxidation state of +4 in sulphur dioxide and +6 in sulphur trioxide.

The oxidation number becomes more positive across the period because the number of valence electrons increase across the period. These valence electrons can take part in bonding with chlorine and oxygen to form chlorides and oxides.

8) Reaction of oxides with water
Metallic oxides reacts with water to form alkaline solutions.
Sodium oxide react vigourously with water to form an alkaline solution of sodium hydroxide.

Magnesium oxide reacts less readily with water to form magnesium hydroxide. Its low reactivity with wtaer is due to the high charge density of Mg ions which holds the oxygen ion more firmly.

Aluminum oxide does not react with water.

Silicon oxide does not react with water.

Non metalllic oxides react with water to form strong acids.
Phosphorous (V) oxide react with water to form phosphoric (V) acid.

Sulphur dioxide react with water to form sulphurous acid.
Sulphur trioxide react with water to form sulphuric acid.

9) Reaction of oxides and hydroxides with acid and NaOH

Metal oxides and hydroxides (Na, Mg) are ionic compounds. They are basic in nature and reacts with acid to form salt and water.

Non metal oxides and hydroxides (Si, P, S, Cl) are covalent compounds. They acidic in nature and reacts with NaOH to form salt and water.

Aluminum oxide is not souble in water but it is amphoteric in nature; i.e. it reacts with both acid and base. Aluminium oxide reacts with HCl to form aluminum chloride and water. Aluminium oxide reacts with NaOH to from an a complex ion (aluminate).

10) Reaction of chlorides with water

Simple ionic chlorides like NaCl and magnesium chloride simply dissolve in water. The solutions of ionic chlorides are neutral.

Aluminum chloride reacts with water to form a complex ion. The aluminum complex can polarize water molecule due to its high charge density. The highly charged aluminium ion draws electrons away from surrounding water molecules, causing them to give up H+.

Silicon chloride reacts with water to form silicon dioxide and HCl.

Phosphorous (V) chloride reacts with water to form phosphoric (V) acid and HCl.

Take home message:
Compounds of Na and Mg are ionic compounds Thus they have high melting points and form basic oxides. Ionic compounds do not react with water, they simply dissolve.

Aluminum compounds are ionic with strong covalent character and form amphoteric oxides. Aluminum compounds react with water to form acidic solutions.

Si compounds has a giant molecular structure and have high melting points. They form acidic oxides.

Phosphorus and sulphur compounds have simple molecular structures and low melting points. They form acidic oxides that react with water giving rise to acidic solutions.

Side note: The syllabus for this section looks very demanding and difficult. However this section isnt really being asked in long answer type of questions. So dont panic

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The Periodic Table: Chemical Periodicity I

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In this chapter we would concentrate on the elements found in Period 3 of the Periodic Table. They are : Na, Mg, Al, Si, P, S, Cl and Ar

1) Variation of atomic radii across Period 3
Atomic radii decreases across the period. Across the period, electron are added to same electron shell (Principal quantum number =3). Shielding effect is similar across the period. However nuclear charge increases across the period due to the increase in number of protons. Hence effective nuclear charge increases and atomic radii decreases.

2) Variation of ionic radii across Period 3
Ionic radii decreases across the period from Na+ to Si4+. Across the period, the outer most electron of the ion is in the 2p subshell, hence the outer most electron experiences the same shielding effect. However nuclear charge increases across the period due to the increase in number of protons. Hence effective nuclear charge increases and ionic radii decreases.

Note that we have to leave out P, S, Cl and Ar in the comparison as their outermost electron is found in the 3p subshell.

3) Variation of melting point across the period
The melting point of an element depends on the bonds present and the structure of the element.

Metals
Na, Mg and Al are metals. The bonds present are metallic bonds defined as the electrostatic forces of attraction between the positive metal ions and sea of delocalised electrons. The metal ions are pack together to form a metallic lattice. Hence the melting point of metals are high and increases from Na to Al as the nuber of delocalized electrons increase.

Si
The atoms of Si are bonded to each other by covalent bonds in a giant molecular structure. When Si melts, all the covlent bonds have to be broken. This is unlike metals where some metallic bonding still remains in the liquid metal. Hence Si has a melting point that is very much higher than metals

P, S, Cl and Ar
Phosphorous, sulphur and chlorine has simple molecular structures. Within the molecule, the atoms are bonded together by strong covalent bonds. Van der Waals interactions exists between molecules. These intermolecular interactions are weak, thus P,S and Cl have low melting points.

4) Variation in the first ionization energy
First IE is the energy required to remove the outermost electron from one mole of atoms in the gas phase. First IE of the elements in period 3 increases across the period. Across the period, electron are removed from the same electron shell (Principal quantum number =3). Shielding effect is similar across the period. However nuclear charge increases across the period due to the increase in number of protons. Hence effective nuclear charge increases and first IE increases.

5) Reaction with oxygen
Na reacts very vigorously with oxygen to form sodium oxide.
Mg reacts very vigorously with oxygen to form MgO.
Al reacts vigorously with oxygen to form aluminium oxide.
P reacts vigourously with oxygen to form phosphorous (V) oxide.
S reacts slowly with oxygen to form sulphur dioxide and sulphur trioxide.

6) Reaction with chlorine
Na reacts very vigorously to giveNaCl.
Mg reacts vigorously to form magnesium chloride.
Al reacts vigorously to form aluminium chloride.
Si reacts slowly to form silicon chloride.
P reacts slowly to form phosphorous (V) chloride.

The Cambridge International Examinations Website has put up the questions for A Levels Nov 2006 Paper 2. These questions are available for download. Students can download the question paper and try question 3. This question is about chemical periodicity. You can try it and check your answers against the mark scheme provided.

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Group II Elements

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The Group II elements included in the syllabus are Magnesium(Mg), Calcium (Ca), Strontium (Sr) and Barium (Ba)

1) Reactions of elements with oxygen
Group II elements react with oxygen to form oxides

2) Reactions with water
Reactivity of group II elements with water increases down the group
Mg reacts slowly with cold water to form MgO and H2. Mg reacts rapidly with steam.
Calcium reacts steadily with water to form calcium hydroxide and hydrogen.
Strontium and barium reacts explosively with water to form hydroxides and hydrogen.

The syllabus only require students to describe the reaction. If you are interested to know why reactivity of group II elements increases down the group, chemguide.co.uk has a good explanation.

3) Behaviour of oxides with water
Reactivity of group II oxides with water increases down the group.
Group II oxides react with water to form hydroxides.

4) Thermal stability of Group II nitrates
Thermal stability of Group II nitrates increases down the group
The thermal stability of Group II nitrates can be explained in terms of the charge density of the metal cation and the polarisability of the nitrate ion.

Down the group, the ionic radii of the Group II elements increases. The charge of the metal ion remains the same, hence charge density decreases. The ion with a lower charge density is less able to polarize the nitrate ion. Hence more heat have to be supplied to decompose the nitrate to metal oxide and NO2.

5) Make predictions
There would e some questions in the exam that would require you to make some predictions. For example, thye woould say an unknown metal X is a group II element. How would its metal hydroxide react with water? So you are expected to know that it would react to form a metal hydroxide.

End of notes for Group II elements
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The Periodic Table: Chemical Periodicity I
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Group II Elements

Group VII: The Halogens

Chemistry of Transition Elements I
Chemistry of Transition Elements II

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Group VII: The Halogens

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The halogens refer to chlorine, bromine and iodine.

1) Colour and volatility
Chlorine is a dense green gas.
Bromine is a brown volatile liquid.
Iodine is a shiny black solid

Down the group, the number of electrons in the halogen molecule increases, hence the strength of van der Waals forces between the halogen molecules increases. This explains why chlorine exists as a gas, bromine exists as a liquid and iodine exists as a solid.

2) Relative reactivity of halogens as oxidizing agents
Chlorine is the strongest oxidizing agent followed by bromine and iodine. Recall that an oxidizing agent causes oxidation and it itself is reduced, i.e. an oxidizing agent has a tendency to gain electrons.


The electrode potential values become less positive down the group, indicating that down the group the halogens has a reduced tendency to gain electrons, hence the oxidizing power of halogen decreases.
3) Reaction of halogens with hydrogen
Chlorine reacts explosively in sunlight with hydrogen. It also reacts slowly in the dark with hydrogen.
Bromine reacts with hydrogen only at high temperatures
Iodine reacts with hydrogen to form an equilibrium mixture of hydrogen, iodine and HI.

4) Relative thermal stability of hydrides (hydrogen halides)
Thermal stability of hydrogen halides decreases down the group.
Down the group, the atomic radii of the halogen increases and the bond dissociation enthalpy of hydrogen halides decreases.
Bond energy of H--Cl : 431 kJ/mol
Bond energy of H--Br: 366 kJ/mol
Bond energy of H--I: 299 kJ/mol

Note: In the exam, students may be asked to quote bond energy values from the Data booklet. You do not have to memorize these bond energies. they are in the data booklet.

5) Reaction of halide ions with aqueous silver ions followed by aqueous ammonia
Chloride ions form a white precipitate with silver ions. White ppt is soluble in aqueous NH3
Bromide ions form a pale yellow ppt with silver ions. Yellow ppt is sparingly soluble in aqueous ammonia.
Iodide ions form a deep yellow ppt with silver ions. Deep yellow ppt is insoluble in aqueous ammonia.

Silver ions is usually supplied as silver nitrate. This is a common test for halides. Aqueous ammonia is added to differentiate between silver chloride and silveer bromide as it is very hard to differentiate between the white and pale yellow ppt.

Aqueous ammonia has to be added slowly as the pale yellow silver bromide is soluble in excess aqueous ammonia. Aqueous ammonia forms a complex with silver ions, shifting the equlibrium to the left, and increases the solubility of AgCl. AgCl is most soluble followed by AgBr and AgI, thus the addition of ammonia can fully solubilise AgCl. AgBr is only soluble in concentrated ammonia while ammonia has no effect on the solubility of AgI.



6) Reaction of halide ions with concentrated sulphuric acid
Concentrated sulphuric acid can act as a proton donor, i.e. it supplies H+ ions which reacts with the halide ion to form hydrides (HCl, HBr and HI). The hydrides appear as steamy white fumes



Concentrated sulphuric acid can also act as an oxidizing agent.
Recall that iodine is the weakest oxidizing agent, hence it is the most easily oxidized. Concentrated sulphuric acid can oxidize iodide ions to iodine, hence you will see purple vapour
Concentrated sulphuric acid can also oxidize bromide ions to bromine gas.
Concentrated sulphuric acid cannot oxidize chloride ions to chlorine gas.

7) Reaction of chlorine with NaOH


Chlorine reacts with cold NaOH to form NaCl and NaClO (sodium hypochlorate)
Note that in this reaction one atom of chlorine is undergoing oxidation and the other atom of chlorine is undergoing reduction.
Chlorine is oxidized. The oxidation number of chlorine increases from 0 in Chlorine gas to +1 in NaClO. Chlorine ia also reduced. The oxidation number of chlorine decreases from 0 in Chlorine gas to -1 in NaCl.



Chlorine reacts with hot NaOH to form NaCl and NaClO3 (sodium chlorate)
In this reaction one atom of chlorine is undergoing oxidation and the other atom of chlorine is undergoing reduction.
Chlorine is oxidized. The oxidation number of chlorine increases from 0 in Chlorine gas to +5 in NaClO3. Chlorine ia also reduced. The oxidation number of chlorine decreases from 0 in Chlorine gas to -1 in NaCl.

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Chemistry of transition elements II

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7) Transition elements in redox systems

These transition metals were highlighted in the syllabus. For e.g. Fe3+ can act as a reducing agent while it itself is oxized to Fe2+.

8) Transition metals react with ligands to form complexes
A ligand possess one or more lone pairs of electrons and donates its electrons to the transition metal ion

A complex is formed by the coordination of lone pairs of electrons from a ligand to a cation which has empty orbitals to accomodate them. Transition metal ions have incompletely filled 3d subshells to accomodate the lone pair donated by the ligand.

A coordinate bond is a covalent bond in which the shared paired of electrons is provided by only one of the bonded atoms. In the case of transition ion complexes, the shared paired of electrons is provided by the ligand

9) Ligand exchange
A stronger ligand can displace a weaker ligand from a complex ion.
The order of ligand strength is CN->NH3>Cl->H2O

CO and oxygen are ligands for Fe in haemoglobin
Haemoglobin contains an Fe2+ ion. Oxygen can act as a ligand and form coordinate bonds with Fe2+. This bonding is reversible to allow haemoglobin to release oxygen where it is needed. Carbon monoxide is a stronger ligand than oxygen. It bonds strongly to Fe2+ . This prevents Fe2+ and thus haemoglobin from bonding with oxygen. This prevents the transport of oxygen in the body eventually leading to death.

10) Transition metals can act as catalyst

Transition metals can act as heterogenous catalysts. Reactant molecules are adsorbedon the surface of the catalyst . The 3d electrons of transition metals enable the transition metal to form temporary bonds with reactant molecules, facilitating the breaking of "old" bonds and the formation of "new" bonds leading to the product formation. Examples include iron (catalyst for the Haber process), Ni (catalyst for hydrogenation).

Transition metals can also act as homogenous catalysts due to the variable oxidation states of transition metals.


For example, the oxidation of iodide ions by peroxodisulphate ions is energetically favourable but in reality the reaction takes place very slowly without a catalyst due to the fact that both reactants are negative ions thus they repel each other.

Fe2+ can act as a catalyst for this reaction. First Fe2+ reduces peroxodisulphate ions to sulphate ions and it itself is oxidized to Fe3+.


Subsequently, Fe3+ oxidizes iodide ions to iodine and Fe2+ is regenerated.

11) Transition metal complexes are coloured
In an isolated tansition metal aton, the 3d orbitals are at the same energy level (degenerate) . However in a complex ion , the d orbitals have slightly different energy due to the presence of ligands (d orbital splitting). An electron can be promoted to a higher energy level by absorbing a photon of light. The color of the complex metal ion is the frequencies of light that are not absorbed.



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