Solubility product Ksp

Solubility product is used to compare the solubility of sparingly soluble ionic solids. I would use AgCl as an example to illustrate solubility product.

Ksp(AgCl) = [Ag+] [Cl-]

The solubility product of AgCl is the concentration of soluble silver ions at equilibrium multiplied by the concentration of soluble chloride ions at equilibrium. Students should note the difference between solubility product and ionic product. Ionic product is simply the product of concentration of ions at a given time, i.e the concentration of ions may not be at equilibrium.

Note that the solubility product concept applies to only sparingly soluble ionic compounds, i.e. it cannot be used for soluble compunds like NaCl. The solubility product is a modified equilibrium constant hence it is affected by temperature.

Common ion effect
The solubility of a sparing soluble ionic compound AB can be reduced by the presence of A+ or B- from a second source. For example we can compare the solubility of AgCl in water and in NaCl solution. The common ion is the chloride ion. Intuitively we know that AgCl is more soluble in water than in NaCl solution. Why is this so? Consider the solubility equilibria of AgCl:


In a solution of pure water, there are no chloride ions, so AgCl dissolves until the ionic product equals to Ksp. In a solution of NaCl, less AgCl can dissolve, as the ionic product will reach Ksp sooner as there are chloride ions already present in solution.

Its quite hard to explain in words but in the A level examinations, students will be asked to do some calculations. Based on the calculations, the effect of the common ion will be more apparent.

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Bicarbonate buffer: Controlling blood pH

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The buffer in blood is made up of carbonic acid (weak acid) and hydrogen carbonate (conjugate base).

When protons is released into the blood, the protons combine with hydrogen carbonate (conjugate base) to form carbonic acid. As a result, the concentration of protons in the blood only increases slightly and the resulting pH change is small.

When hydroxide ions are released into the blood, the alkali that is added reacts with carbonic acid to form salt an dwater. As a result, the hydroxide ions are removed from blood, and the pH of blood only changes slightly.

The explanation that I have given above is actually a oversimplified one. For a more physiological and thorough explanation, please refer to the following website on Blood, Sweat and Buffers. For the purpose of the A Level examinations, my explanation will suffice.

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Buffers and pH control

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Buffers are solutions which resist changes in pH on addition of small amount of acid or alkali. Buffers usually consist of a weak acid and its conjugate base (acid buffer) or a weak base and its conjugate acid (basic buffer).

How does an acid buffer work?

As mentioned above, an acid buffer consists of a weak acid and its conjugate base. An example is ethanoic acid and its conjugate base ethanoate ion. The conjugate base is usually supplied as a salt. The buffer is most effective at resisting pH changes when the concentration of the undissociated acid molecule (ethanoic acid) is similar to the concentration of the conjugate base (ethanoate ion).

When a small amount of acid is added, the protons combine with the conjugate base (ethanoate ions) to form the weak acid (ethanoic acid). The H+ that is added is removed, hence the concentration of H+ in solution and thus the pH changes only slightly.

When a small amount of base is added, the base reacts with the weak acid to form salt and water. The hydroxide ions are removed from the solution in the form of water and the pH of the solution only changes slightly.

Calculating the pH of buffer solution

The pH of a buffer solution can be calculated using the Handerson-Hasselbach equation.

Students should note that this equation can only be used for calculating the pH of buffer solutions. Also the equation above can only be used to calculate the pH of acid buffers. For base buffers, use the equation below

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Acids and bases II

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3) pH, dissociation constants (Ka, Kb) and ionic product of water (Kw)

The pH of a solution is the negative logarithm to base ten of the molar concentration of hydrogen ion. Students should note that pH only tells you how acidic or basic a solution is, it does not tell you about the strength of the acid or base. A concentrated solution of a weak acid will have a low pH , but it does not mean that the acid is a strong acid

In order to measure the strength of the acid , chemists use the acid dissociation constant. The acid dissociation constant is an equilibrium constant. Basically it is the ratio of concentration of protons and conjugate base to the concentration of undissociated acid molecules. Since acid dissociation constants are equilibrium constants, they are not affected by concentration, unlike pH which is concentration dependent. The greater the Ka, the stronger the acid.


Dissociation of water
Water dissociates to a very small extent to form hydrogen ions and hydroxide ions. This is represented by the ionic product for water Kw, which is the [H+] X [OH-], 10e-14 for pure water at 25 degree celsius. Note that like acid dissociation constant, the ionic product for water is effectively an equilibrium constant, hence it is affected by temperature, i.e. Kw will vary with temperature.

4) Indicators for acid base titration
The most common indicators are litmus, methyl orange and phenolpthalein. Methyl orange is red at acidic pH, orange at pH 3 to 4 and yellow at pH 5 onwards. Litmus paper is red at acidic pH and blue at pH values of 7 and above. Phenolpthalein is colourless at at acidic pH (pH 0 to 7 ) and red at alkalike pH (pH 9 and above). Students shouldnote that not all indicators change colour at pH 7.

In a titration of strong acid against a strong alkali, at equivalence point when the acid is completely neutralized, the pH changes by a large extent from around pH 4 to 10. Thus both methyl orange and phenolpthalein can be used to indicate that equivalenc epoint has been reached.

In the titration of a strong acid with a weak base, the pH change is less drastic, from around pH 4 to 8, hence phenolpthalein is not a suitable indicator as it changes colour at around pH 8. Methy orange will be a suitable indicator.

In the titration of a weak acid with a strong base, the pH change at equivalence point is from around pH 7 to 11. Phenolpthalein will be a suitable indicator as it changes colour at around pH 8.
In the titration of a weak acid with weak base, the pH change at equivalence point is very small. hence there are no suitable indicators. You can probably detect the pH change using a pH meter.

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Acids and bases I

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1) Bronsted-Lowry Theory of Acids and Bases

According to the Bronsted Lowry theory, an acid is defined as a substance that can donate proton to another substance, i.e. a proton donor. An example of an acid is hydrochloric acid. Students also need to know what a conjugate base is. An acid whhich has donated its proton becomes a conjugate base. Using hydrochloric acid as an example, hydrochloric acid is the acid and the chloride ion is the conjugate base. Using ethanol as an example, if ethanol behaves as a Bronsted acid, it loses its proton and becomes ethanoate ion. Therefore the conjugate base is the ethanoate ion.


A base is a substance which accepts a proton from another substance, i.e. a proton acceptor. An example is ammonia. Ammonia is a Bronsted base, it accepts a proton and becomes the ammonium ion. In this case ammonia is the base and the ammonium ion is the conjugate acid. Note that a bronsted base is related to its conjugate acid and a bronsted acid is related to its conjugate base.

Note that certain acids and bases do not conform to the bronsted acid and base definition. Take for example sodium hydroxide, it is well know that sodium hydroxide is a base. However NaOH is not a proton acceptor hence it is not a Bronsted base.

There is another definition of acids and bases proposed by Arrhenius. According to the Arrhenius definition, an acid dissociates in water to form hydrgen ions and a base dissociates in water to produce hydroxide ions. NaOH will fit the Arrhenius definiton of a base. There are other definitions such as the Lewis acid definition. For more information students can refer to the wikipedia entry.

2) Differences between a weak and strong acid.

A strong acid such as hydrochloric acid dissociates fully in water to form hydrogen ions and chloride ions. A weak acid dissociates partially in water. An example is ethanoic acid. When ethanoic acid is dissolved in water, some of the ethanoic acid molecules will dissociate to form the ethanoate ion and hydrogen ions. However not all of the ethanoic acid molecules dissociate so there will be ethanoic acid molecules in the wtaer in addition to ethanoate ions and protons.

Students must be careful not to mix up concentration and strength of acid. A concentrated acid might not be a strong acid. A 5 M ethanoic acid is a concentrated acid solution but ethanoic acid is not a strong acid.

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Updates

I have added a navigation bar near the top of the blog. You can access the sitemap via this bar

For the chapter on Chemical periodicity, I have also added a question from the actual A Levels Examinations. Students are encouraged to try the question to get an idea of how this chapter can be tested in the examinations.

The chapters on Chemical Periodicity has been rewritten. Hope it is more clear now. The textbook by Hill and Holman, Chemistry in Context has a good chapter written on this topic.

Chemical Equilibria II

9）The equilibrium in terms of partial pressure Kp
Kp is generally used when the reaction involves gaseous reactants and products. For example, reaction between hydrogen and iodine to form hydrogen iodide.


10）Haber Process

Haber process involves the production of ammonia from nitrogen and hydrogen gas. Recall that the reaction is exothermic and that 1 mol of nitrogen gas reaacts with 3 mol of hydrogen gas to produce 2 mol of ammonia gas. By Le Chatelier's Princple high pressure and low temperature will increase the proportion of ammonia at equilibrium.

However industrially a high temperature of 450 degrees and a pressure of 250 atm is used. Although the yield of ammonia is lower at higher temperature, using a low temeprature reduces the rate of reaction making the process uneconomical. Fe is also used as a catalyst to speed up the rate of reaction.

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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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Electrochemistry I
Electrochemsistry II
Simple Electrochemical Cell
Limitations of Standard Cell Potentials
Fuel Cells

Electrolysis I
Electrolysis II
Calculations involving electrolysis
Electrolysis in industrial process

Reaction kinetics I
How to determine the order of reaction?

Chemical Equilibria I
Chemical Equilibria II

Acids and Bases I
Acids and Bases II
Buffers and pH control
Bicarbonate buffer: Controlling blood pH
Solubility Product: Ksp

The Periodic Table: Chemical Periodicity I
The Periodic Table: Chemical Periodicity II

Group II Elements

Group VII: The Halogens

Chemistry of Transition Elements I
Chemistry of Transition Elements II

Worked examples

N2007/II/5
N2007/III/1
N2007/III/2
N2007/III/5

Exam Advice
A Level chemistry Syllabus
Cambridge International Examinations Website
How to use the worked examples?