O Level Chemistry Revision Notes

Hwa Chong Institution (High School)
Secondary 4 End of Year Examinations
Chemistry Revision Notes

List of topics:

1. Atomic Structure
2. Separation Techniques
3. Chemical Bonds and Bonding (Ionic, Covalent)
4. Metals
5. Properties and structures of compounds
6. Qualitative Analysis
7. Acids, Bases and Salts
8. Chemical Periodicity
9. Chemical Calculations
10. Air and Environment
11. Electrochemistry (Including Redox reactions)
12. Thermochemistry
13. Rates of Reaction
14. Chemical Equilibrium
15. States of Matter
16. Organic Chemistry

1) Atomic Structure

(a) Model of Atom

The model of atom consists of 3 subatomic particle, namely the protons, electrons and neutrons. Of which, the protons and neutrons are found within the nucleus of the atom. As such, the nucleus of the atom has a positive charge, and most of the mass of the atom is located in the nucleus due to the insignificance of the weight of electrons. Surrounding the nucleus is a cloud of electrons, which move randomly around orbitals and sub orbitals at nearly the speed of light. Other than this, most of the atom is empty space. The overall atom is electrically neutral, as the number of protons is equal to the number of electrons, while the nucleus carries no charge.

(b) Sub atomic particles

• Protons, which carries a positive electric charge • Neutrons, which has the same mass as a proton but carried no charge • Electrons, which carries a negative electric charge

| |relative mass |relative charge |
|Proton |1 |+1 |
|Neutron |1 |0 |
|Electron |1/1836 |-1 |

(c) Behaviour of sub atomic particles in electric field

• Protons are positively charged and so would be deflected on a curving path towards the negative plate. • Electrons are negatively charged and so would be deflected on a curving path towards the positive plate. • Neutrons don't have a charge, and so would continue on in a straight line.
1) If particles have same energy: • Protons are deflected on a curved path towards the negative plate. • Electrons are deflected on a curved path towards the positive plate. • Neutrons continue in a straight line.
The amount of deflection is exactly the same in the electron beam as the proton beam if the energies are the same - but, of course, it is in the opposite direction.

2) If particles have same speed:

• Protons are deflected on a curved path towards the negative plate, but are attracted more quickly • Electrons are deflected on a curved path towards the positive plate, but are attracted more slowly • Neutrons continue in a straight line.
If the electrons and protons are travelling with the same speed, then the lighter electrons are deflected far more strongly than the heavier protons.

(d) Numbers of protons, neutrons and electrons
1) Atom number/ Proton number ( Gives the number of protons in an atom of particular element
2) Mass number/ nucleon number ( Gives the sum of number of protons and neutrons
3) Number of electrons ( As aforementioned, number of electrons is equal to the number of protons, since the atom is electrically neutral.

(e) Isotope
Isotopes are atoms of the same element which have the same atomic number but different mass numbers, as they have the same number of protons but different numbers of neutrons. The relative abundance of an element is the approximate nucleon number of an element taking into consideration all available isotopes:

[pic]
(f) Atomic Orbitals and sub orbitals
Electrons in fact inhabit regions of space known as orbitals. The Heisenberg Uncertainty Principle says that it is impossible to define with absolute precision, at the same time, both the position and the momentum of an electron.

Each orbital has a name. The orbital occupied by the hydrogen electron is called a 1s orbital. The "1" represents the fact that the orbital is in the energy level closest to the nucleus. The "s" tells you about the shape of the orbital. ‘S’ orbitals are spherically symmetric around the nucleus - in each case, like a hollow ball made of rather chunky material with the nucleus at its centre.

A p orbital is rather like 2 identical balloons tied together at the nucleus. The diagram on the left is a cross-section through that 3-dimensional region of space. Once again, the orbital shows where there is a 95% chance of finding a particular electron. Unlike an s orbital, a p orbital points in a particular direction - the one drawn points up and down the page.

At any one energy level it is possible to have three absolutely equivalent p orbitals pointing mutually at right angles to each other. These are arbitrarily given the symbols px, py and pz. This is simply for convenience - what you might think of as the x, y or z direction changes constantly as the atom tumbles in space.

In addition to s and p orbitals, there are two other sets of orbitals which become available for electrons to inhabit at higher energy levels. At the third level, there is a set of five d orbitals (with complicated shapes and names) as well as the 3s and 3p orbitals (3px, 3py, 3pz). At the third level there are a total of nine orbitals altogether.

At the fourth level, as well the 4s and 4p and 4d orbitals there are an additional seven f orbitals - 16 orbitals in all. s, p, d and f orbitals are then available at all higher energy levels as well.

(g) Electrons in box diagrams and electronic configuration
Orbitals can be represented as boxes with the electrons in them shown as arrows. Often an up-arrow and a down-arrow are used to show that the electrons are in some way different. Other than drawing electron in box diagrams, the electronic configuration of atoms can also be written out in the following form:
[pic]
Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones. Where there is a choice between orbitals of equal energy, they fill the orbitals singly as far as possible. These are hereby summarised below as the 3 rules of electronic configuration:

1) The Pauli Exclusion Principle suggests that only two electrons with opposite spin can occupy an atomic orbital.

2) Hund's rule suggests that electrons prefer parallel spins in separate orbitals of sub shells. This rule guides us in assigning electrons to different states in each sub-shell of the atomic orbitals. In other words, electrons fill each and all orbitals in the sub shell before they pair up with opposite spins.

3) The AUFBAU Principle is derived from the first 2 rules and states that no two electrons in the atom will share the same four quantum numbers, and also, electrons will first occupy orbitals of the lowest energy level. In addition, electrons will fill an orbital with the same spin number until the orbital is filled before it will begin to fill of the opposite spin number.
[pic]
(h) Ground state and excited state
The energy associated to an electron is that of its orbital. The energy of a configuration is often approximated as the sum of the energy of each electron, neglecting the electron-electron interactions. The configuration that corresponds to the lowest electronic energy is called the ground state. Any other configuration is an excited state. The electron is only and most stable in the ground state, while it tend to exist similar to radicals in the excited state.

(i) Ionisation Energy of an atom
1) First ionisation energy
The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions. This is summarised in the equation:
[pic]
Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high). All elements have first ionisation energy.

2) Factors affecting first ionisation energy
Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.
The size of that attraction will be governed by:

(a) The charge on the nucleus
The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.
(b) The distance of the electron from the nucleus
Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.
(c) The number of electrons between the outer electrons and the nucleus
Consider the electronic configuration of sodium. If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.
(d) Whether the electron is on its own in an orbital or paired with another electron.
Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.

3) General Trends of ionisation energy
Within 1 period, the first ionisation energy usually increases as we move across the period of elements from left to right. For example in the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p. These are all the same sort of distances from the nucleus, and are screened by the same 1s2 electrons, hence they have approximately the same shielding effect. The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies. In fact the increasing nuclear charge also drags the outer electrons in closer to the nucleus. That increases ionisation energies still more as you go across the period.
Down a group, the first ionisation energy usually decreases. For example in Group I, as the atomic radius increases and the negatively charged electron is further from the positively charged nucleus it is less attracted to the nucleus (electron is said to be 'shielded' due to shielding effect).
Amongst the transition metals, the first ionisation energy is approximately equal. This is as all of these elements have an electronic structure [Ar] 3dn 4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comes from the 4s orbital. As you go from one atom to the next in the series, the number of protons in the nucleus increases, but so also does the number of 3d electrons. The 3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each other out as far as attraction from the centre of the atom is concerned.
However, anomalies always exist for these 3 trends, which would be further explained the the topic chemical periodicity.
(j) Representation of electronic configuration and bonding
1) Electronic configuration
The electronic configuration of an atom is represented in this case by written forms of the electron in box diagrams, in which the orbitals and sub orbitals would be presented. It can also be presented in the shorter version:
Ar[pic]1s22s22p63s23px23py23pz2[pic][Ne]3s23px23py23pz2
2) Electron in box diagram
This is the lengthened version of the written electronic configuration, in which the electrons are drawn in boxes representing the orbitals and sub orbitals and observing the 3 rules of filling electrons:
[pic]
3) Dot cross diagrams
Dot cross diagrams are used to represent ionic and covalent bonds between elements:
[pic] [pic]
4) Lewis Structure
Lewis structure is another form of representing covalent bonds
[pic]
2) Separation Techniques

(a) Importance of separation technique
Separation techniques are important as these are the means in which purification of substances is achieved. A pure substance is a single substance not mixed with anything else. An impure substance is a mixture of substances, in which impurities and contaminants are present. Separation of intended substances from the contaminants is important and is known as purification.

To check for purity of substances, simply observe their melting and boiling points. Substances that have a fixed melting and boiling point are pure, while those in which melting and boiling points spans over a range of values are impure. This is as the melting and boiling points of pure substances are mostly unique.

(b) Filtration
Filtration is used to separate an insoluble solid from a liquid. In this technique, the mixture is poured through a filter funnel lined with filter paper, which only allows the liquid to pass through, while the solid remains on the filter paper. The resultant liquid obtained is the residue, while the resultant solid is the filtrate.

(c) Crystallisation
Crystallisation separates a dissolved solid from a solution as well formed crystals. One way to do this is to heat a solution to evaporate off most of the solvent, before allowing saturated solution to cool. It occurs because the solubility of most solutes decreases as temperature decreases. As a hot solution cools, it eventually become saturated, and hence the extra solute that cannot be dissolved separates into pure crystals, while impurities present remain in the saturated solution.

(d) Sublimation
Sublimation separates a mixture of solids, one of which is able to sublime. A few substances, such as iodine, changes from solid to vapour state directly when heated. Hence, by heating this mixture, the solid that can sublime would become vapour form, leaving the other solid.

(e) Decanting
Decanting is a very quick method for separating a mixture of a liquid and a heavier solid. If we want to separate a mixture of water and same, we should first allow the sand to settle on the bottom of the container. Then we poured off the water at the top. The advantage of this method is quick, but there is a disadvantage of this method which is rough. It cannot be used to separate a mixture of a liquid and a light solid, such as chalk in water. The particles of chalk are suspended in the water. They are so light that they do not sink down to the bottom for a long time. Also, not all of the water would be poured away accurately.
(f) Centrifugation
Centrifugation is used when we want to separate small amounts of suspension. The suspension of solid in liquid is poured into a centrifuge tube, then spin around very fast in a centrifuge. The spinning motion forces the solid to the bottom of the tube. Then the liquid can be poured off from the solid. Centrifugation is commonly used in dairies to separate milk from cream to make skimmed milk. It is possible because milk has less density than cream.

(g) Evaporation to dryness
This is used to separate a solution which has different boiling points. The solution is heated so that the solvent evaporates, and just leave the solid behind. Only solute can be obtained, while the solvent will evaporate.

(h) Simple distillation
Simple distillation separates a pure liquid from a solution, and can be used to obtain a pure solvent from a solution of a solute. In distillation, the liquid is changed into vapour by boiling. The vapour is a pure substance. The vapour is then cooled, hence condensing into a pure liquid called the distillate. The resultant solid left behind is also pure.

(i) Fractional Distillation
Fractional distillation separates mixtures of miscible liquids with widely differing boiling points via the use of a fractionating column. The mixture is continuously heated. The substance with the lower boiling point would boil and evaporate first. As the vapour passes through the condenser at the top of the fractionating column, it is cooled and condenses into liquid form. The substance which has higher boiling point would remain in the flask until the other substance has been distilled. Glass beads in the column ensure that the substance with higher boiling point does not evaporate to the top of the column, by condensing on these beads. These beads provide a large surface area for repeated vapourisation and condensation to occur. On the condenser, the water supply cools the condenser down constantly so that condensation would occur. Two layers of the condenser ensure purity of substance.

(j) Using separating funnel
A separating funnel is used to separate immiscible liquids. The mixture is allowed to stand, and the two liquids would soon form 2 separate layers, with the denser liquid at the bottom. The stopper is removed and the tap opened, allowing bottom layer to run off and leaving the top layer.

(k) Chromatography
Chromatography is used to separate and identify mixtures.
1) Use of chromatography • To identify mixtures of coloured substances found in food • To separate substances in drugs, urine and blood, which can be useful in tracking if an athlete has taken drugs

2) Separating mixtures of coloured substances
A drop of solution to be tested is dropped on the pencil line near the bottom of a strip of chromatography paper. The paper is then dipped to a suitable level into a tube containing a solvent with the solvent level below the spot. The solvent travels up the paper, and the dyes on the pencil line dissolve in the solvent and travel up the paper in different speeds, hence becoming separated in the chromatogram.

3) Identifying mixtures of coloured substances
A drop of dye to be tested is placed on pencil line, together with other dyes which is possibly within the dye to be tested. The paper is dipped in the solvent, and results are produced. If dyes travel the same distance up the paper, the dyes are identical.

4) Separating and identifying mixtures of colourless substances
In this case, the chromatogram is sprayed with a locating agent, which is a substance that reacts with the substances on the paper to produce a coloured product, hence showing the location of substances on the paper.

5) Alternative method of paper chromatography
Paper chromatography can also be carried out with the solvent running down the paper. This works better for longer pieces of paper as the solvent does not have to move against gravity, and thus flows more quickly.

6) Safety precautions for paper chromatography
To ensure reliability and accuracy of test, several precautions are taken. Firstly, the line is drawn using a pencil so that the solution would not be contaminated by the dyes from pen inks, hence rendering the result accurate. Also, the solvent level is below the spot of solution to be tested so that the spot would not be smudged and the solvent is allowed to travel up the paper in suitable levels, hence allowing for chromatogram to be easily read and reliable.

7) Rf values
Rf value is obtained by dividing the distance moved by the substance with the distance moved by the solvent.
(l) Summary of separation techniques
|Technique |Purpose |
|Filtration |Separate solid from liquid |
|Crystallisation |Separate dissolved solid from solution |
|Sublimation |Separate 2 solids, of which 1 can sublime |
|Decanting |Separate solid from liquid |
|Centrifugation |Separate solid suspension from liquid |
|Evaporation to dryness |Separate solution with different boiling points |
|Simple Distillation |Separate pure liquid from solution |
|Fractional Distillation |Separate 2 miscible liquids with different boiling points |
|Separating funnel |Separate 2 immiscible liquids |
|Paper chromatography |Separate and identify substances in mixtures |

3) Chemical Bonds and Bonding

(a) Formation of Ionic Bonds
It is a bond formed by a metallic and a non-metallic ion (eg. Na+ & Cl-), or rather, formed by the attraction between two oppositely charged ions. Electrons from the metal are transferred to the non-metal during this process, forming a cation and an anion, and in which electrostatic attraction is produced, which brings the two ions together, thus forming an ionic bond. It is only possible to form and ionic bond between 2 ions.

(b) Factors affecting the strength of an ionic bond

The factors refer to the charge on both the positive and negative ions, and the size (ionic radii) of the positive and negative ions.

The higher the charge, the stronger the bond.

The smaller the size, the higher the strength of the bond.

The higher the charge density, the stronger the ionic bond.

Now, if the total energy released is more than that which is absorbed, then the formation of ionic compound is favoured. The conditions that favour the formation of an ionic bond (or ionic compound) are summarized below: • Low ionisation energy of the metallic element, which forms the cation. o Ionisation energy is the amount of energy, which is required to remove the most loosely bound electron(s) from an isolated gaseous atom to form a positive ion. In forming an ionic bond, one atom must form a cation by losing one or more electrons. In general, elements having low ionisation energies have a more favourable chance to form a cation, thereby having a greater tendency to form ionic bonds. Thus, lower ionization energy of metallic elements favours the formation of an ionic bond. It is because of low ionization energy that the alkali and alkaline earth metals, form ionic compounds. • High electron affinity of the non-metallic element, which forms the anion. o Electron affinity is the amount of energy released, when an isolated gaseous atom accepts an electron to form a negative ion. The other atom participating in the formation of an ionic compound must form an anion by gaining an electron (s). Higher electron affinity favours the formation of an anion. Therefore, generally, the elements having higher electron affinity favour the formation of an ionic bond. Halogens have high electron affinities, and therefore halogens generally form ionic compounds. • Large lattice energy (the smaller size and higher charge of the ions) o When a cation, and an anion come closer to each other, they get attracted to each other due to the coulombic force of attraction. These electrostatic forces of attraction between oppositely charged ions release a certain amount of energy (when the ions come closer) and an ionic bond is formed. If the coulombic attractions are stronger, then more energy gets released and a more stable ionic bond is formed. o Lattice energy 'is the energy released when one mole of an ionic compound in crystalline form is formed from the constituent ions'. Therefore, larger lattice energy would favour the formation of an ionic bond. Lattice energy thus is a measure of coulombic attraction between the combining ions. The lattice energy of an ionic compound depends directly on the product of the ionic charges, and inversely on the square of the distance between them. Lattice Energy=q1xq2\d2.
(c) Structure of an Ionic Compound

An ionic compound consists of ions held together in a giant crystal lattice structure by ionic bonds. In most cases, the positively-charged portion is made up of metal cations and the negatively-charged portion is an anion or polyatomic ion. The ions are secured by the electrostatic force between oppositely charged bodies. Ionic compounds have high melting and boiling points and are hard and brittle. Ions that lose or gain electrons are positioned side by side to each other. The anion and cation is attracted to each other by electrostatic forces of attraction due to opposite charges.

Ionic compounds are formed when two differently charged bodies form a bond to become less reactive. Ionic compounds are arranged in a giant lattice, cubed structure. By giant, it means that the structure contains a very large and almost uncountable amount of ions packed together in that shape. The cation loses an electron and is attracted to the anion which gains 1 electron through electrostatic force. This causes the latticed structure to take shape, thus giving it its shape. However, it need not necessarily be in a cubic shape.

(d) Strength of Ionic Bond

It is measured using Lattice Energy. Lattice energy 'is the energy released when one mole of an ionic compound in crystalline form is formed from the constituent ions'. Lattice Energy actually measures the energy required to convert the solid ionic compound into gaseous compounds. In this case, the ionic bond(s) are formed, not separated, thus the Lattice Energy would be negative. Negative LE means energy is given off to the surroundings.

(e) Physical Properties of Ionic Compounds

Ions have high melting and boiling points, are very hard but brittle, are semi-conductive, have high solubility in water and forms crystals. Ionic compounds have strong bonds between the ions. These strong bonds increase the boiling and melting temperature as it requires a very large amount of energy to break these bonds. However, when a large enough force is applied to the ionic compound, the bonds does not merely break in 1 spot. The compounds break into many pieces as the ionic bonds break in many spots. When stress is applied to a side of the compound, the part where stress is applied to shifts slightly to the side. This causes similar ions to touch each other, thus repelling themselves away from each other and splitting the compound. Thus, ionic compounds are brittle.

As for electric conductivity, ions are semi conductive. When in molten or aqueous form, the ionic compound conducts electricity. However, in solid form, an ionic compound does not conduct electricity. Ionic compounds in the solid state do not conduct electricity, as the electrons in the compound are not free to move due to the ions in the compound already sharing them. However, upon becoming molten or when they are dissolved, the ions are separated and gain electrical conductivity as they are now mobile and able to move freely within the lattice structure.

The more complex and larger the structure is, the more bonds there will be, and the more energy will be needed to break the bonds between the ions, making the structure stronger.

Thus, the general properties of an ionic compound are summarized below:

|Criteria |Properties |Reason |
|Melting and Boiling Points |High |Strong bonds, large amount of energy required to |
| | |break them |
|Material Strength |Strong | |
|Brittleness |Brittle |Stress applied, similar ions touch each other, repel|
| | |away |
|Electric conductivity (solid) |Does not conduct electricity |Ions are immobile |
|Electric conductivity (molten/aqueous) |Conducts electricity |Ions are separated and mobile |
|Solubility in water |Highly Soluble |Water molecules can pull hard enough to break the |
| | |lattice structure |
(f) Formation of a covalent bond

A covalent bond is formed when atoms share a pair of valence electrons when their atomic orbitals overlap. It is formed by the sharing of a pair of electrons of opposite spins between 2 or more atoms so as to achieve a noble gas electronic configuration. Compounds are covalently compounded when the difference in the electronegativities between the atoms are not significantly large (especially between non-metals where transfer of electrons is difficult due to the high ionization energies involved). When the atomic orbitals of the bonding atoms overlap, molecular orbitals are formed in which the shared pair of electrons can exist in. The force of repulsion between the nuclei is offset by the attractive forces between the electron cloud and the nuclei. As non metals tend to have similar high electronegativity, neither atom can take electrons from the other, forcing them to share electrons.

(g) Sigma and Pi Bonds The different kinds of orbital overlap are by sigma bonds and pi- bond. A sigma bond is a way of orbital overlap where it is due to head-on overlap of atomic orbitals. What makes each of these a sigma bond is that the orbital overlap occurs directly between the nuclei of the atoms. A pi- bond is a way of orbital overlap where it is due to side-on overlap of atomic orbitals. Therefore, a sigma bond is stronger than a pi- bond as in Sigma bonds the orbitals from 2 different atoms have a head-to-head alignment with each other and therefore have a larger overlap area than Pi bonds where the 2 orbitals are parallel to each other and have a smaller overlap area. A single bond is made up of a sigma bond. A double bond is made up of a sigma bond and a pi-bond. A triple bond is made up of a sigma bond and two pi-bonds.

(h) Strength of covalent bonds

Strength of covalent bond is measured by its bond dissociation energy, a measure of how much energy is required to break the covalent bond, otherwise known as bond enthalpy. It reflects the average amount of energy required to break 1 mole of covalent bonds in gaseous molecules. The larger the bond energy, the stronger the bond.

To determine strength of covalent bond, we look at the covalent bond length, the distance between the centres of the nuclei of the two atoms joined by a covalent bond. Half of the bond length of a single bond of two similar atoms is called covalent radius. Covalent bond length is influenced by the size of the two atoms joined by the covalent bond and the number of bonds present. The longer the bond length, the weaker the covalent bond.
Covalent bonds are formed as a result of the sharing of one or more pairs of bonding electrons. The electro negativities (electron attracting ability) of the two bonded atoms are either equal or the difference is no greater than 1.7. As long as the electro-negativity difference is no greater than 1.7, the atoms can only share the bonding electrons.
(i) Electronegativity in covalent bonds

Electronegativity is defined as the ability of an atom in a particular molecule to attract electrons to itself. It causes polarity in a covalent bond because it may cause the atoms not to be equal in their attraction towards the electrons.
So it can be concluded also that the higher the value of electronegativity for one atom, the more the electrons are attracted towards a particular atom and a polar covalent bond is established. The greater the difference in electronegativity between two covalently bonded atoms, the more polar the bond.

(j) Polarity in covalent bonds

If the molecule is polar, the electrons will be attracted more towards a particular atom because the electrons are attracted differently by the atoms.

When a covalent bond is formed between 2 similar atoms from the same element, the electron pair that is being shared will lie exactly midway between the 2 atoms. This is because the attraction towards each of the atoms is exactly the same.

Dipole is formed when the atoms of different elements combine and shares the electrons, thus the atom of the element with a higher electronegativity attracts the shared pair of electrons more towards itself causing a separation of a positive and a negative charge.

(k) Forces of attraction in covalent bonds

Three types of intermolecular covalent bonds:

Van der Waals forces, consisting of:

1. Temporary dipole-dipole interactions

▪ The weakest force among the 3. Molecules do not actually have an even electrical charge across the molecule. Electrons are mobile and move within their electron shell, so there will be instances where one region of the molecule contains more electrons, making that end more negatively charged while the other end is more positively charged.

▪ This creates a dipole, which is a separation of the positively and negatively charged sides. Thus, the more positively charged side of a covalent molecule will pull the negatively charged electrons of another molecule towards it, also causing an imbalance in the placement of electrons in the second molecule, inducing a dipole in it as well. This creates a covalent bond.

▪ Temporary dipole-dipole interactions are the weakest intermolecular force as the polarity of a side of a molecule can quickly change as the electrons move, so the forces can be easily overcome.

▪ Permanent dipole-dipole interactions

▪ Stronger than temporary dipole-dipole interactions, but weaker than hydrogen bonds. Only occurs between molecules that have permanent dipoles (polar molecules). Such molecules have their component atoms arranged such that one side of the molecule has a positive charge while the other side has a negative charge due to an imbalance of the number of electrons on each side. Because one side of the molecule is positively-charged while the other side is negatively charged, the molecules will thus be attracted to each other due to electrical attraction.

▪ Permanent dipole-dipole interactions are stronger than temporary dipole-dipole interactions as the molecules in this interaction are permanent dipoles, compared to the molecules in dispersion forces where they are only temporary dipoles.

2. Hydrogen Bonds

▪ The strongest force among the 3 and is a special case of the permanent dipole-dipole interactions. Occurs only in molecules containing hydrogen atoms and oxygen, fluorine or nitrogen atoms, and only if the hydrogen atoms are directly bonded to the oxygen, fluorine or nitrogen atoms. It is a special case of the dipole-dipole interaction as a hydrogen atom is the only atom that does not have a core shell of electrons, only a valence shell. (Helium is a noble gas and rarely bonds with other chemicals)

▪ When hydrogen atoms bond with oxygen, fluorine or nitrogen atoms, the positively charged nucleus is exposed and is easily attracted to the very electronegative (able to attract electrons to itself) atoms of the other molecules.

▪ Hydrogen bonds are the strongest among the three bonds (intermolecular forces of attraction) as the hydrogen nucleus is positively charged and very small and is attracted to the very electronegative atoms of oxygen, fluorine or nitrogen in the other molecules so the attraction will be greater than the other 2 bonds.

(l) Determining Forces of Attraction

First, we check if the given covalent molecule is polar or non-polar. If it is non-polar, the intermolecular attraction is due to temporary dipole-dipole interactions. If the molecule is polar, we check if there are any hydrogen atoms present in the molecule, and whether they are bonded to either an oxygen, nitrogen or fluorine atom. If they are, the molecules are attracted due to hydrogen bonds. If not, they are attracted by permanent dipole-dipole interactions.

(m) Trends in physical properties of Covalent Substances – Halogens (Fluorine, Chlorine, Bromine, Iodide)

1) Generally have much lower melting and boiling points than ionic compounds.

2) Tend to be more flammable than ionic compounds.

3) Do not conduct electricity as they are formed by sharing electrons between non-metals.

4) Aren't usually very soluble in water.

Trends down the group of halogens

1) Increasing melting point

2) Increasing boiling point

3) Decreasing reactivity

4) Decreasing solubility

5) Increasing in density

6) Darkening in colour

(n) Trends in physical properties of covalent substances- Hydrogen molecules (Hydrogen fluoride, Hydrogen chloride, Hydrogen bromide, Hydrogen Iodide)

Solubility in water
HCl, HBr, HI compounds are bonded by temporary dipole - dipole interactions/attractions. The attraction forces are weaker and tend to dissociate in aqueous solution. HF compound have molecules bonded by forces of attraction of hydrogen bonding. Hydrogen Bonding is a very strong type of covalent bonding. Fluorine is also a very electronegative element. Therefore bond overlap would be strongest and bonds are hard to overcome as Hydrogen and Fluorine atoms would be attracted very closely to each other.

Melting Point
Generally covalent substances are held by weak van der Waals forces, and therefore they melt at very low temperatures, compared to ionic compounds held strongly in their lattice structures.

Conductivity of Electricity
Able to conduct electricity to the free electrons dissociated from the following compounds in aqueous solution.

Material Strength
Generally weak due to weak van der Waals forces.

(o) Trends in physical properties of covalent substances- Gases and water (Hydrogen, Oxygen, nitrogen, water, ammonia, carbon dioxide)

|Melting and Boiling points |Relatively lower as compared to ionic compounds, hence most covalent substances are liquids and gases, if |
| |not, are solids with low melting points |
|Electrical Conductivity in water |Mostly no |
|Material strength |Relatively softer than ionic compounds |
|Solubility in Water |Generally insoluble and barely soluble |
|Flammability |Relatively more flammable than ionic compounds |
|Reactivity |Unreactive |

(p) Trends in physical properties of covalent substances- Giant Covalent Molecules (Graphite, diamond, sand, silicon)

|Physical Properties |Graphite |Diamond |Sand |Silicon |
|Structure |Giant covalent |Giant covalent |Ordered molecular |Crystal structure |
| |structure |structure |structure | |
|Conduction of electricity |Able to conduct |Unable to conduct |Unable to conduct |Unable to conduct |
| |electricity |electricity |electricity |electricity |
|Melting Point |High melting point |High melting point |High melting point |High melting point |
|Solubility in water |Insoluble in water |Insoluble in water |Insoluble in water |Insoluble in water |
|Material strength |Soft and brittle |Hard and brittle |Very brittle |Brittle |

What is a Giant Covalent Structure?
•Contains many non-metal atoms joined together by covalent bonds to form a giant lattice
•High melting and high boiling points
•Extremely strong structures because of many bonds formed (exception of graphite)
•For graphite, each carbon atom forms three covalent bonds, forming giant molecules in hexagonal layers
•For diamond however, each carbon atom forms four covalent bonds

Solubility
1) Weak attractions between the carbon atoms and the water molecules
2) Carbon atoms are bonded very tightly to each other; water molecules are unable to pass through Melting/ Boiling Point
•For all the four covalent substances, covalent bonding between the atoms is very strong
•A lot of heat energy is required to break the covalent bonding to cause a change in state
•Therefore, the boiling/melting point of the substances is very high
(q) General properties of covalent substances

|Criteria |Properties |Exceptions |
|State |Usually liquid or gases |Some are in solid |
|Volatility |Volatile |- |
|Melting/Boiling Points |Low |Giant covalent substances |
|Solubility in water |Insoluble |- |
|Electric Conductivity |Do not conduct electricity |Non-polar molecules (small amount) |
|Material Strength |Weaker than ionic compounds |- |
|Hardness |Soft |- |
|Brittleness |Not brittle/Sturdy |- |

(r) Dative bonds

A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

4) Metals

Metals are giant structures of atoms held together by strong metallic bonds and closely packed together, that is, they fit as many atoms as possible into the available volume. Each atom in the structure has 12 touching neighbours such a metal is described as 12-co-ordinated. However, some metals (notably those in Group 1 of the Periodic Table) are packed less efficiently, having only 8 touching neighbours. These are 8-co-ordinated, and they also have varying numbers of delocalized electrons.

(a) Physical properties of metals

• Metals have high melting and boiling • Good conductors of heat and electricity • Large majority are malleable and ductile

(b) What it affects?

• Melting and Boiling Points of water
The strength of the metallic bonds affects the Boiling/Melting point, since more/less energy would be required to break/form these bonds. Metals tend to have high melting and boiling points because of the strength of the metallic bond. The strength of the bond varies from metal to metal and depends on the number of electrons which each atom delocalises into the sea of electrons, and on the packing.

• Conduction of Electricity
The number of delocalized electrons affects the conduction of electricity. Since more delocalized electrons would mean the metal conducts electricity with greater efficiency. The delocalised electrons are free to move throughout the structure in 3-dimensions. They can cross grain boundaries. Even though the pattern may be disrupted at the boundary, as long as atoms are touching each other, the metallic bond is still present.

• Material Strength
The ability of the atoms to roll over each other into new positions without breaking the metallic bonds affects material strength. Metals are described as malleable (can be beaten into sheets) and ductile (can be pulled out into wires).

Thermal Conductivity
Metals are good conductors of heat. Heat energy is picked up by the electrons as additional kinetic energy (it makes them move faster). The energy is transferred throughout the rest of the metal by the moving electrons.

(c) Comparing and Contrasting Group 1 vs Group 2 Metals

Group I metals (eg. Li, Na, K, Rb…)

•the physical properties are similar because they are in the same group
•low melting and boiling points
•each atom has only one electron to contribute to the bond
•less efficiently packed as compared to other metals
•relatively large atoms
•react violently with water.

Na vs. Mg; K vs. Ca (ie. Group I vs. Group II metals)

•alkaline earth metals that are considered highly reactive, but not as reactive as metals in Group I
•harder and denser, and have higher melting points
•have two valence electrons on each atom, while Group I metals have one
•the former has stronger metallic bonding.

Na, K, Ca etc vs. Fe, W, Cu (Main group metals vs. transition metals)

•high melting points and boiling points
•can involve the 3d electrons in the delocalization as well as the 4s
•the force of attraction is stronger
•more energy would be required to break the forces of attraction
•boiling and melting points of the metals would naturally be higher

Physical Properties of Mercury

• only metal that is in liquid form at room temperature and pressure • poor conductor of heat energy but fair conductor of electricity

• melting point of mercury is -38.83oC, while boiling point is 356.73 oC • insoluble in water and has 11 isotopes • Reacts with sulphur. • combination of mercury vapour and noble gases results from an electrical discharge • occurs naturally as mercuric sulfide

5) Structures and properties of substances

(a) Molecular Geometry of substances

1) Rules governing molecular geometry

The rule governing molecular geometry would be the Valence Shell Electron-Pair Repulsion Theory, which focuses on the positions taken by the groups of electrons on the central atom of a simple molecule. The theory states that the shape of a molecule or ion is governed by the arrangement of the electron pairs around the central atom. All you need to do is to work out how many electron pairs there are at the bonding level, and then arrange them to produce the minimum amount of repulsion between them. You have to include both bonding pairs and lone pairs.

Electron clouds are negatively charged since the electrons are negatively charged, so electron clouds repel one another and try to get as far away from each other as possible. Lone pairs of electrons exert a greater repelling effect than bonding pairs do, while lone pair-bonding pair repulsion is greater than bonding pair-bonding pair repulsion. Hence, lone pair-lone pair repulsion > lone pair-bonding pair repulsion > bonding pair-bonding pair repulsion.

2) Finding out the molecular geometry of a molecule

(1) Write down the Lewis dot structure for the molecule.
(2) Count the number of bond pairs and lone pairs around the central atom.
(3) Decide on the electron pair orientation based on the total number of electron pairs.
(4) Consider the placement of lone pairs and any distortions from "regular" shapes.
(5) Name the shape based on the location of central atom.

3) Molecular Geometry

|Total number of electron pairs|Number of bond pairs |Number of lone pairs |Molecular Geometry |
|2 |2 |0 |Linear |
|3 |2 |1 |Bent |
|3 |3 |0 |Trigonal Planar |
|4 |2 |2 |Bent |
|4 |3 |1 |Trigonal Pyramidal |
|4 |4 |0 |Tetrahedral |
|5 |2 |3 |Linear |
|5 |3 |2 |T-shaped |
|5 |4 |1 |Seesaw |
|5 |5 |0 |Trigonal Bipyramidal |
|6 |4 |2 |Square Planar |
|6 |5 |1 |Square Pyramidal |
|6 |6 |0 |Octahedral |

(b) Structure and Properties of Ionic Compounds

1) Structure
Ionic compounds exist in a giant, ionic regular lattice with alternative anions and cations held together by strong electrostatic forces of attraction.

An ionic compound consists of ions held together in a giant crystal lattice structure by ionic bonds. In most cases, the positively-charged portion is made up of metal cations and the negatively-charged portion is an anion or polyatomic ion. The ions are secured by the electrostatic force between oppositely charged bodies. Ionic compounds are arranged in a giant lattice, cubed structure. By giant, it means that the structure contains a very large and almost uncountable amount of ions packed together in that shape. The cation loses an electron and is attracted to the anion which gains 1 electron through electrostatic force. This causes the latticed structure to take shape, thus giving it its shape.

2) Properties
(a) Physical hardness
Ionic compounds are hard and crystalline solids with flat sides and regular shapes. This is as the ions are arranged regularly, forming a large crystal. The ions are held in place by strong electrostatic forces of attraction, which make the crystals hard.
(b) Melting and Boiling Points
Strong electrostatic forces of attraction exist between ions in the ionic compound, and a high amount of energy is required to overcome these forces, hence ionic compounds have high melting and boiling points.
(c) Volatility
Ionic compounds are non-volatile as they cannot evaporate easily due to the strong electrostatic forces of attraction between ions.
(d) Electrical Conductivity
Solid ionic compounds cannot conduct electricity as ions are immobile. However, when in molten or aqueous form, they can conduct electricity as mobile ions are available to carry the electric current.
(e) Solubility
Many ionic compounds are soluble in water, as the ions attract water molecules which disrupt the crystal lattice structure, causing ions to separate.

(c) Structure and Properties of Simple Covalent Compounds

1) Structure
Simple covalent compounds exist as simple, discrete covalent molecules with weak van der Waals forces of attraction existing between molecules.

2) Properties
(a) Physical Hardness
Most covalent substances are liquids or gases in room temperature, as van der Waals forces are weak, hence molecules are not held together tightly and can move about freely.
(b) Melting and Boiling Point
Simple covalent compounds have low melting and boiling points, as little amount of energy is required to overcome the weak van der Waals forces of attraction between molecules (However, strong covalent bonds between atoms in each molecule are not broken).
(c) Volatility
Many simple covalent compounds are highly volatile and evaporate easily to give a smell, as molecules can easily separate from each other due to the weak van der Waals forces.
(d) Electrical conductivity
Simple covalent compounds do not conduct electricity, as they do not contain any mobile ions or electrons.
(e) Solubility
Many simple covalent compounds are insoluble in water due to insufficient interaction between water molecules and covalent molecules (except for molecules which exhibit hydrogen bonding). However, they are soluble in organic solvents.

(d) Structure and Properties of Giant Covalent Compounds

1) Structure
Giant covalent compounds exist in a giant regular covalent lattice, which consists of a large network of atoms held together by strong covalent bonds. Strong covalent bonds exist between atoms of giant covalent compounds.

2) Properties
(a) Physical state
All giant covalent compounds are solids at room temperature, as strong covalent bonds exist between atoms of giant covalent compounds, which make them very hard.
(b) Melting and boiling points
All giant covalent compounds have very high melting and boiling points, as strong covalent bonds exist between atoms of giant covalent compounds, hence a higher amount of energy is required to break the covalent bonds between atoms.
(c) Electrical conductivity
All giant covalent compounds do not conduct electricity, except graphite. This is as all valence electrons of these compounds are used to form covalent bonds, hence there are no free electrons and electrical conduction does not occur.

3) Examples
(a) Diamond
Diamond is the hardest natural substance in the world, and this is due to the strong covalent bonds which exist between carbon atoms of the diamond compound. Therefore, it is used in drill bits, rocks drills and saws to cut up other hard solids.
(b) Graphite
Graphite has different properties from diamond due to different structure. Firstly, graphite is soft as the forces between carbon layers are weak and so layers can easily slide past each other, making graphite soft and slippery. As layers of carbon atoms can slide off the pencil onto the paper, graphite is used as pencil lead. Also, it is used as a lubricant as it does not decompose at high temperatures due to strong covalent bonds between carbon atoms.
Also, graphite conducts electricity. In every molecule of graphite, 1 carbon atom forms bonds with 3 other carbon atoms, hence leaving 1 free electron per carbon atom. As such, these electrons are mobile and hence can conduct electricity, thus it is used as electrodes.

(e) Structure and Properties of Metallic Substances

1) Structure
Metal exists in a giant metallic structure, which consist of a regular lattice of metal cations surrounded in a sea of delocalised, mobile electrons.

2) Properties
(a) Physical State (Malleability and Ductility)
Metals are malleable and ductile as layers of atoms in a metal can slide over each other easily when a force is applied, but does not break as the electrons hold the atoms together.
(b) Electrical conductivity
Metals are good conductors of electricity due to the presence and movement of free, mobile delocalised electrons through the metal that can carry the electric charge.
(c) Thermal Conductivity
Metals are good conductors of heat due to both lattice vibration and movement of free, mobile delocalised electrons through the metal via free electron diffusion. When heat is applied to one end of the metal, delocalised electrons obtain more energy and hence move faster, thus colliding with neighbouring electrons and transferring the heat energy to the neighbouring electrons, as a result conducting heat.
(d) Melting and Boiling Points
Strong metallic bonding exists within the metal and hence a high amount of energy is required to overcome the strong metallic bonding.
6) Qualitative Analysis
(a) Solubility of substances

1) All sodium, potassium, ammonium and nitrate salts are soluble in water.

2) All chloride and iodide salts are soluble in water except silver and lead (II) chloride and iodides.

3) All sulfate salts are soluble in water except lead (II), calcium and barium sulfate.

4) All oxides and hydroxides insoluble in water except barium, potassium and sodium oxides and hydroxides. (Calcium oxide and hydroxide are sparingly soluble.)

5) All carbonates are insoluble in water except sodium, potassium and ammonium carbonates.

(b) Reactivity of substances

|Most reactive |Potassium |
| |Sodium |
| |Calcium |
| |Magnesium |
| |Aluminium |
| |Zinc |
| |Iron |
| |Tin |
| |Lead (II) |
|Point of inflexion |Hydrogen |
| |Copper |
| |Mercury |
| |Silver |
| |Gold |
|Least reactive |Platinum |

(c) Test for anions

The principle behind most of the QA for anions is based on the insolubility of the salts formed between the metal cation and the anion to be tested. All the anion tests (except nitrate) require acidification with dilute nitric acid to ascertain the unknown is not a carbonate. If no effervescence is noted when the acid is added (the conclusion one can draw is that the carbonate anion is absent), the aqueous reagents are then added to separate portions of the unknown.

General Instructions:

1) Perform test for carbonate anion. To a 2cm3 portion of unknown, add an equal volume of dilute nitric acid. If there is a brief effervescence of colourless, odourless gas which gives a white precipitate when bubbled into limewater observed, the anion is carbonate.

2) Conduct test for sulfate anion. To a 2cm3 portion of unknown, acidify the unknown with an equal volume of nitric acid before adding aqueous barium nitrate (or the other reagents) solution dropwise. If white precipitate is formed, anion is sulfate (or corresponding anions).

3) Conduct test for nitrate anion. To a 2cm3 portion of unknown, add an equal volume of sodium hydroxide solution, before placing some aluminium foil and warming it gently. If an evolution of colourless, pungent gas which turns damp red litmus paper blue is observed, the anion of nitrate.

|Anion |Reagent |Observation |
|Carbonate |Dilute nitric acid |Brief effervescence of colourless, |
| | |odourless gas which gives a white |
| | |precipitate when bubbled into limewater |
| | |observed |
|Sulfate |Dilute nitric acid |Formation of a white precipitate in a |
| |Aqueous Barium Sulfate |colourless solution |
|Chloride |Dilute nitric acid |Formation of a white precipitate in a |
| |Aqueous Silver Nitrate |colourless solution |
|Iodide |Dilute nitric acid |Formation of a yellow precipitate in a |
| |Aqueous Lead (II) Nitrate |colourless solution |
|Nitrate |Aqueous Sodium Hydroxide |Evolution of colourless, pungent gas |
| |Aluminium Foil |which turns damp red litmus paper blue |
| |Warm Gently | |

(d) Cation Test

This test makes use of the solubility of hydroxides as well as the colour of the hydroxide precipitate formed to determine the identity of the cation in solution. Excess sodium hydroxide solution / aqueous ammonia is added to check if the metal ion is one that will form soluble complexes. Cation tests must always be conducted when the unknown is in solution form.

General Instructions:

1) Conduct the cation test for ammonium. To 2cm3 of unknown, add an equal volume of aqueous sodium hydroxide before warming the solution gently. If an evolution of colourless, pungent gas which turns damp red litmus paper blue is observed, the cation is ammonium.

2) Conduct the cation test for calcium. To 2cm3 of unknown, add aqueous sodium hydroxide dropwise until in excess. To another test tube of 2cm3 unknown, add aqueous ammonia dropwise until in excess. If the cation is calcium, a white precipitate that is insoluble in excess aqueous NaOH would be observed for aqueous sodium hydroxide test, while for aqueous ammonia test, no (or slight white precipitate) would be observed.

3) Conduct the cation test for aluminium and lead (II). To 2cm3 of unknown, add aqueous sodium hydroxide dropwise until in excess. To another test tube of 2cm3 unknown, add aqueous ammonia dropwise until in excess. If the cation is aluminium or lead (II), a white precipitate that is soluble in excess aqueous sodium hydroxide giving a colourless solution would be observed for aqueous sodium hydroxide test, while for the aqueous ammonia test, a white precipitate that is insoluble in excess aqueous ammonia would be observed.

4) Conduct test to determine if cation is aluminium or lead(II). Add aqueous sodium chloride solution dropwise into the test tube. If a white precipitate is formed, the cation is lead (II). Otherwise, it is aluminium.

|Cation |Reagent |Observation |
|Ammonium |Aqueous Sodium Hydroxide |Evolution of colourless, pungent gas |
| |Warm gently |which turns damp red litmus paper blue |
| |Aqueous Ammonia |- |
|Calcium |Aqueous Sodium Hydroxide |White precipitate in colourless solution,|
| | |insoluble in excess aqueous NaOH |
| |Aqueous Ammonia |No (or slight white) precipitate |
|Aluminium |Aqueous Sodium Hydroxide |White precipitate in colourless solution,|
| | |soluble in excess aqueous sodium |
| | |hydroxide giving a colourless solution |
| |Aqueous Ammonia |White precipitate in colourless solution,|
| | |insoluble in excess aqueous ammonia |
| |Aqueous sodium chloride |No observable change |
|Lead (II) |Aqueous Sodium Hydroxide |White precipitate in colourless solution,|
| | |soluble in excess aqueous sodium |
| | |hydroxide giving a colourless solution |
| |Aqueous Ammonia |White precipitate in colourless solution,|
| | |insoluble in excess aqueous ammonia |
| |Aqueous sodium chloride |Formation of a white precipitate in a |
| | |colourless solution |
|Zinc |Aqueous Sodium Hydroxide |White precipitate in colourless solution,|
| | |soluble in excess aqueous sodium |
| | |hydroxide giving a colourless solution |
| |Aqueous Ammonia |White precipitate in colourless solution,|
| | |soluble in excess aqueous ammonia giving |
| | |a colourless solution |
|Copper (II) |Aqueous Sodium Hydroxide |Light blue precipitate in colourless |
| | |solution, insoluble in excess aqueous |
| | |sodium hydroxide |
| |Aqueous Ammonia |Light blue precipitate in colourless |
| | |solution, soluble in excess ammonia |
| | |giving a dark blue solution |
|Iron (II) |Aqueous Sodium Hydroxide |Pale green precipitate in colourless |
| | |solution, insoluble in excess aqueous |
| | |sodium hydroxide |
| |Aqueous Ammonia |Pale green precipitate in colourless |
| | |solution, insoluble in excess aqueous |
| | |ammonia |
|Iron (III) |Aqueous Sodium Hydroxide |Reddish brown precipitate in colourless |
| | |solution, insoluble in excess aqueous |
| | |sodium hydroxide |
| |Aqueous Ammonia |Reddish brown precipitate in colourless |
| | |solution, insoluble in excess aqueous |
| | |ammonia |

(e) Test for Gases

General Instructions:

1) Perform test to test for presence of hydrogen gas. Firstly, place a damp litmus paper near the mouth of the test tube. If the damp litmus paper does not change colour, preliminary test suggests it coudl be hydrogen or oxygen gas. Place a lighted splint near the mouth of the test tube. If the lighted splint is extinguished with a 'pop' sound, the gas is hydrogen gas.

2) Perform test to test for presence of oxygen gas. Firstly, place a damp litmus paper near the mouth of the test tube. If the damp litmus paper does not change colour, preliminary test suggests it coudl be hydrogen or oxygen gas. Place a glowing splint near the mouth of the test tube. If the glowing splint is rekindled, the gas is oxygen gas.

3) Perform test to test for presence of carbon dioxide gas. Firstly, place a damp blue litmus paper near the mouth of the test tube. If the damp blue litmus paper turns red, preliminary test suggests that it could be carbon dioxide gas. Bubble the colourless, odourless gas into limewater solution. If a white precipitate is formed, the gas is carbon dioxide gas.

4) Perform test to test for presence of chlorine gas. Place a damp blue litmus paper near the mouth of the test tube. If the damp blue litmus paper turns red before becoming bleached after some time, the gas is chlorine gas.

5) Perform test to test for presence of ammonia gas. Place a damp red litmus paper near the mouth of the test tube. If the damp red litmus paper turns blue, the gas is ammonia gas.

6) Perform test to test for presence of sulfur dioxide gas. Bubble the gas through aqueous acidified potassium dichromate (VI) solution. If the orange aqueous acidified potassium dichromate (VI) solution turns green in colour, the gas is sulfur dioxide gas.

|Gas |Test (Preliminary/ Confirmatory) |Observation |
|Hydrogen Gas (Colourless, odourless gas) |Damp Litmus Paper |No change in colour of damp litmus paper |
| | |(no observable changes) |
| |Lighted Splint |Extinguishes lighted splint with a ‘pop’ |
| | |sound |
|Oxygen Gas (Colourless, odourless gas) |Damp Litmus Paper |No change in colour of damp litmus paper |
| | |(no observable changes) |
| |Glowing Splint |Rekindles glowing splint |
|Carbon dioxide gas (Colourless, odourless|Damp Blue Litmus Paper |Damp blue litmus paper turns red in |
|gas) | |colour |
| |Bubble gas into limewater |Formation of a white precipitate in |
| | |colourless solution |
|Ammonia Gas (Colourless, pungent gas) |Damp Red Litmus Paper |Damp red litmus paper turns blue in |
| | |colour |
|Chlorine Gas (Greenish yellow, pungent |Damp Blue Litmus Paper |Damp blue litmus paper turns red in |
|gas) | |colour before becoming bleached |
|Sulfur Dioxide Gas (Reddish brown, |Bubble the gas through aqueous acidified |Orange aqueous acidified potassium |
|pungent gas) |potassium dichromate (VI) solution |dichromate (VI) solution turns green in |
| | |colour |

7) Acid, Bases and Salt

(a) Definition of acids and bases

1) Arrhenius’ Definition of acid and bases
This definition of acid states that an acid is a substance that produces hydrogen ions when dissolved in water. Therefore, according to this definition, all acids are soluble in water and can form the aqueous form. This definition of base also states that a base is a substance that produces hydroxide ions when dissolved in water. This theory is limited in the sense that many acids and bases which can be included in the other 2 definitions are left out simply because they do not produce hydrogen or hydroxide ions in water. One such case would be the reaction of ammonia with an acid.

2) Bronsted-Lowry Definition of acid and bases
This definition of acid states that an acid is a proton (hydrogen ion) donor, while a base is a proton (hydrogen ion) receiver. This theory simply expands the scope of the first definition by including more substances in the list of acid and bases. This definition is also used to determine the strength of acids and bases. Using the above example of reaction of ammonia and hydrochloric acid, this is now an acid-base reaction as in this case, ammonia accepts a proton (a hydrogen ion) from hydrochloric acid molecules. The hydrogen becomes attached to the lone pair on the nitrogen of the ammonia via a dative bond.

The term conjugate pair is also derived from this definition of acid and bases. A conjugate acid is the acid member of a pair of two compounds that transform into each other by gain or loss of a proton. A conjugate acid can also be seen as the chemical substance that donates in the forward chemical reaction of a reversible reaction. A conjugate base is the base produced due to the acceptance of a proton from the acid.

3) Lewis Definition of acid and bases
The Lewis definition of acid and bases does not focus on the transfer of hydrogen ion, as in the former 2 definitions. In this case, the Lewis definition talks about the donating or receiving of an electron pair from the other party involved in the reaction. Hence, a Lewis base is a compound that donates an electron pair during the reaction, while a Lewis acid is one that accepts this electron pair. This removes the requirement of presence of hydrogen ion in acid-base reactions. This further expands the scope of acid and bases- for example, by the second definition, boron trifluoride is not an acid as it is not a donor of hydrogen ion. However, in this definition, it is a Lewis acid by accepting the nitrogen's lone pair.

(b) Acids and acid reactions
As mentioned, acids are either donors of hydrogen ions, or substances that produce hydrogen ions when dissolved in water. All acids have a sour taste (degree of sourness depends on strength of acid), are hazardous to the human skin by causing it to redden and blister when in contact, and can change the colour of pH indicators. Acid have a pH of less than 7.

1) Reaction of acid with metals
Most dilute acids react with metals above hydrogen in the reactivity series to produce hydrogen gas and a metal salt. The validity of this test can be confirmed by testing for hydrogen gas produced. If hydrogen gas is produced, burning splint put near mouth of test tube in which reaction is occurring would extinguish with a ‘pop’ sound.
[pic]
2) Reaction of acid with carbonate
In this reaction, acid reacts with a carbonate to produce the resultant salt, water and carbon dioxide. The validity of this test can be confirmed by testing for carbon dioxide gas produced, in which a white precipitate would be formed when gas is bubbled into limewater. Reaction of acid with hydrogen-carbonates is the intermediate step of this reaction.
[pic]
3) Reaction of acid with metal oxides and hydroxides
In this reaction, acid reacts with metal oxides and hydroxides to produce the resultant salt and water. This is a relatively slow process.
[pic]
[pic]
4) As such, acids are used for various purposes. Sulfuric acid acts as an important industrial catalyst in the formation of ester as an organic reaction. It is also used in the manufacture of agricultural fertilisers, as well as other products such as detergents, paints and so on. Acids can also be used to remove rust, which is iron (III) oxide which can dissolve in acids.
5) Basicity of an acid
The basicity of an acid is defined by the maximum number of hydrogen ions that can be produced by an acid molecule. A monobasic acid, such as hydrochloric acid, is an acid which can form 1 hydrogen ion per acid molecule when dissolved in water.
6) Strength of an acid
The strength of an acid depends on the proportion of acid molecules which dissolve in water to produce hydrogen ions. A strong acid is one in which all the acid molecules fully dissolve in water to produce hydrogen ions. A strong acid would generally have a low pH, as they would produce a huge amount of hydrogen ions in water. Examples of strong acids include sulphuric acid, nitric acid and hydrochloric acid. A weak acid is one in which only a small proportion of acid molecules dissolve in water to produce hydrogen ions, some of which include citric acid. A weak acid would have a higher pH compared to a strong acid, as they would produce only a small proportion of hydrogen ions. Of course, strong acids react more vigourously than weak acids due to the availability of much more hydrogen ions to react.
7) Strength vs concentration
The strength of an acid is the measure of the proportion of acid molecules which dissolve in water to produce hydrogen ion. However, the concentration of an acid is the number of moles of acid molecules present within a fixed amount of acid solution.

(c) Bases and alkali
Bases are oxides and hydroxides of metals. Within the bases, there is a small group of substances called alkali, which is unique as it dissolves in water to produce hydroxide ions. Hence, an alkali is always a base, but not vice versa. Alkalis are generally slippery, and are also hazardous when in high concentrations. They also change the colour of pH indicators, and they have a pH of higher than 7 (7 to 14).

1) Reaction of alkali with acid
Alkali reacts with acid in the process called neutralisation, which produces salt and water. Heat is also liberated in the process.
[pic]
2) Reaction of alkali with ammonium compounds
Alkali reacts with ammonium compound under heat to produce water, ammonia gas and the resultant salt.
[pic]
3) Reaction of alkali with solutions of metal ions
Alkali reacts with solutions of metal ions, such as copper(II) ions, to produce the resultant metal hydroxide and salt. This is called a precipitation reaction, which can be used to identify metal ions.
[pic]
4) Use of alkalis
Alkalis are found in window cleansers, detergents, toothpaste and medicines, amongst other substances used for cleaning. The purpose of alkali in such cases would be to neutralise the acid, such as in toothpaste and antacids. Another purpose would be to dissolve grease.
5) Strength of alkali
The strength of an acid depends on the proportion of alkali molecules which dissolve in water to produce hydroxide ions. A strong base is one in which all the acid molecules fully dissolve in water to produce hydroxide ions. A strong acid would generally have a high pH, as the concentration of hydrogen ions in water would significantly lessen with more hydroxide ions. Examples of strong alkali include potassium hydroxide and sodium hydroxide. A weak alkali is one in which only a small proportion of alkali molecules dissolve in water to produce hydroxide ions, some of which include aqueous ammonia. A weak base would have a lower pH compared to a strong alkali. Of course, strong alkali react more vigourously than weak alkali due to the availability of much more hydroxide ions to react.

(d) pH and indicators pH is a measure of the acidity of a solution, and has a range of 0 to 14. The value of the pH (how high it is) is inversely proportionate to the concentration of hydrogen ions present in the solution. A substance of pH 7 is neutral, and one of such substance is pure water. A solution with pH less than 7 is acidic, as there is an increase in concentration of hydrogen ions which is produced by acid dissolving in water. A solution with pH more than 7 is alkaline, as hydroxide ions is produced which decreases the concentration of hydrogen ions in solution. The equation to measure pH is as follows:
[pic]
The pH of a solution is predicted via the use of indicators. An indicator is a substance that shows whether a substance is acidic or alkaline by changing to different colour when in different acidity of solutions. Below is a list of common indicators and their colour change:
|Indicator |Acidic (7) |
|Litmus Paper |Red |Blue |
|Methyl Orange |Red |Yellow |
|Phenolphthalein |Colourless |Pink |

(e) Neutralisation
Neutralisation is the process involving the reaction between an acid and a base, in which the resultant salt is produced, together with water as the only 2 products of the reaction. Heat is liberated in this exothermic reaction. The basis of this reaction is that an acid produces hydrogen ions while alkali produces hydroxide ions, both of which react to form water.

Neutralisation is important in our daily life and can be used for many purposes. One of such purpose would be to control the pH of the soil, treating indigestion/gastric problems, treating acidic insect stings, treating industrial wastewater and in toothpaste.

(f) Oxides
1) Acidic oxides
Acidic oxides are oxides of non metals, and they react with water to produce acids, as well as react with alkali to form salt and water. Some examples of acidic oxides include SO2, SO3, CO2 and P4O10.
2) Neutral oxides
Neutral oxides are also oxides of non metals, and they do not react with acids and bases, hence they do not form salts. Examples include H2O, CO and NO.
3) Amphoteric Oxides
Amphoteric oxides are metallic oxides which can react with both acids and with alkalis to form salt and water. 3 of such oxides are Al2O3, ZnO and PbO.
4) Basic Oxides
Basic oxides are metallic oxides which react with acid to form salt and water. Basic oxides can be soluble or insoluble in water. Examples include CuO and CaO.

(g) Ionic Equations
An ionic equations shows the actual reaction taking place between the reactants. Ions that are present but do not participate in the reaction are removed when writing ionic equations. To write an ionic equation, write the balanced chemical equation of the reaction first. Following that, break all compounds up into their base ions. Keep all non-aqueous compounds on both sides of the equation and remove those not participating in the reaction (hence not in the product). Ensure the final ionic equation is balanced.

(h) Salt Preparation
The method to be adopted depends on the solubility of the salt to be prepared, as well as the solubility of the starting base used.

1) Titration (Soluble base, soluble salt)
(This method makes use of the neutralisation reaction between alkali and acid to produce salt and water. The water would then be evaporated before crystallising the salt so as to obtain salt crystals.)
(1) A known volume of acid (20 cm3) is pipetted into a conical flask and 2 drops of methyl orange indicator is added. The acid is titrated with the alkali from the burette.
(2) The alkali is added until the indicator turns orange, pH 7 neutral. This means all the acid has been neutralised to form the salt. End point is reached.
(3) The volume of alkali needed for neutralisation is then noted, this is called the end point volume. (1)-(3) are repeated with both known volumes mixed together but without the universal indicator.

(4) The solution is transferred to an evaporating dish and heated to evaporate the water causing crystallisation to occur.
(5) The crystals can be carefully collected and dried between 2 pieces of filter paper.

2) Reaction of acid with excess metal carbonate/hydroxide/oxide (Insoluble base, soluble salt)
(This method makes use of the neutralisation reaction between base and acid. Excess base is used to ensure that all the acid has been reacted and neutralised. The resultant would be filtered and filtrate collected and dried to obtain salt crystals.)
(1) The required volume of acid (20 cm3) is measured out into the beaker with a measuring cylinder. The insoluble metal, oxide, hydroxide or carbonate is added into the acid in the beaker with stirring in excess.
(2) When no more of the solid dissolves it means all the acid is neutralised.
(3) The mixture is filtered to remove the excess base and the filtrate is collected..
(4) The filtrate is then heated before being left to cool and crystallise in a evaporating dish.
(5) Collect and dry the crystals between 2 pieces of filter paper..

3) Precipitation (Soluble base, insoluble salt)
(This method makes use of the neutralisation reaction between base and acid. The resultant salt is insoluble and hence can be easily obtained via filtering and collecting the residue.)
(1) Mix an equal volume of 2 solutions, 1 containing the cation of the salt and the other containing the anion, in a beaker. Stir the resultant mixture with a glass rod A precipitate of the insoluble salt would be formed.
(2) Filter the resultant mixture with filter funnel and filter paper and obtain the residue.
(3) Wash the residue with deionised water.
(4) Dry the residue between 2 sheets of filter paper

8) Chemical Periodicity
(a) Basis of chemical periodicity – effective nuclear charge
1) Nuclear charge
The nuclear charge is the total charge of all the protons in the nucleus. The increase in nuclear charge is proportionate to the increase in the number of protons in the nucleus of an atom, which increases the total positive charge of the atom.
2) Shielding effect
The shielding effect is the effect of the electrons shielding the valence electrons from the electrostatic forces of attraction exerted by the positive charge in the nucleus. This is caused by the shielding electrons, which are the electrons in the orbitals between the valence electrons and the nucleus. The shielding effect increases with the number of core electron shells, as in this case, there are more electrons to provide the shielding effect.
3) Effective nuclear charge
Effective nuclear charge is the charge felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus. In mathematical terms, effective nuclear charge is the result obtained when the nuclear charge of an atom is subtracted by the shielding effect acting on the nucleus. Hence, effective nuclear charge is proportionate to the number of protons and inversely proportionate to the number of core electron shells in an atom.

(b) Periodic Trends – Ionisation energy
The first ionisation energy is the energy required to remove the outermost electron from one mole of gaseous atoms to produce 1 mole of gaseous ions.
Across the period, the ionisation energy generally increases. Across the period, the number of core electron shells is equal as electrons are added to the same shell. Hence, there is only a slide, negligible increase in shielding effect across the period. The major difference is the increasing number of protons in the nucleus across the period. That causes greater attraction between the nucleus and the electrons and so increases nuclear charge, which also pulls the valence electrons in closer to the nucleus. Thus, with significant increase in nuclear charge and negligible increase in shielding effect, effective nuclear charge increases across the period. As such, more energy is required to remove the outermost electron in an atom, hence ionisation energy increase across the period.
However, anomalies occur in 2 cases, 1 between Beryllium and Boron, the other between Magnesium and Aluminium. In both cases, there is a slight dip in the ionisation energy. The explanation lies with the structures of boron and aluminium. The outer electron is removed more easily from these atoms than the general trend in their period would suggest. You might expect the boron value to be more than the beryllium value because of the extra proton. Offsetting that is the fact that boron's outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, and the electron is, on average, to be found further from the nucleus. This has two effects. Firstly, the increased distance results in a reduced attraction and so reduced ionisation energy. Also, the 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy. The explanation for the drop between magnesium and aluminium is the same, except that everything is happening at the 3-level rather than the 2-level.

Down the group, the ionisation energy decreases. The increase in nuclear charge down the group, due to increase in number of protons, is offset by the greater increase in shielding effect, due to increase in number of core electrons down the group. As a result, effective nuclear charge decreases, less ionisation energy is needed to remove outermost electron.

(c) Periodic Trends – Atomic Radii
The atomic radius is a measure of the size of its atoms, usually the mean or typical distance from the nucleus to the boundary of the surrounding cloud of electrons. As the boundary is ill-defined, the atomic radius is measured by halving the distance between the nuclei of 2 touching atoms.
The atomic radii of elements decrease across the period. Across the period, electrons are added to the same electron shell, hence the number of core electron shell is the same, and thus shielding effect is approximately equal for all elements across a period. The increasing number of protons in the nucleus as you go across the period increases the nuclear charge. Hence, effective nuclear charge increases, pulling the valence electrons more tightly towards the nucleus, thus decreasing the atomic radii of atoms across the period.
The atomic radii of elements increase down the group. Down the group, the number of core electron shell increases, and thus shielding effect increases significantly for all elements down the group. This offset the increase in nuclear charge, which is brought about by the increasing number of protons in the nucleus down the group. Hence, effective nuclear charge decreases down the group, thus pull of valence electrons towards the nucleus weakens down group, thus decreasing the atomic radii of atoms down the group.

(d) Periodic Trends – Ionic Radii
The ionic radius is a measure of the size of its ions, usually the mean or typical distance from the nucleus to the boundary of the surrounding cloud of electrons. As the boundary is ill-defined, the ionic radius is measured by halving the distance between the nuclei of 2 touching ions.
Across the period, the ionic radius generally decreases. From Group I to III, ions in the same period are said to be isoelectronic, hence they have the same amount of electrons and so same number of core electron shells, so shielding effect is approximately equal. However, the number of protons in the nucleus of the ions is increasing, so nuclear charge increase. Hence, effective nuclear charge increases, pulling the valence electrons more tightly towards the nucleus, thus decreasing the ionic radii of atoms across the period.
From Group V to VII, ions in the same period are also isoelectronic, hence they have the same amount of electrons and so same number of core electron shells, so shielding effect is approximately equal. However, the number of protons in the nucleus of the ions is increasing, so nuclear charge increase. Hence, effective nuclear charge increases, pulling the valence electrons more tightly towards the nucleus, thus decreasing the ionic radii of atoms across the period.
Between the cations and anions, there is an increase in the ionic radii. This is due to an additional core electron shell added for anions due to addition of electrons instead of removal (as in the case of cations), which increases the shielding effect. Hence, effective nuclear charge for cation is larger than anion, so ionic radii are smaller in cations due to stronger pull of valence electrons towards nucleus.
Down the group, the ionic radii increase. Down the group, the number of core electron shell increases, and thus shielding effect increases significantly for all elements across a period. This offset the increase in nuclear charge, which is brought about by the increasing number of protons in the nucleus down the group. Hence, effective nuclear charge decreases down the group, thus pull of valence electrons towards the nucleus weakens down group, thus decreasing the ionic radii of atoms down the group.
(e) Periodic Trends – Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons towards itself in formation of bonds.
The electronegativity of elements increase across the period. Across the period, electrons are added to the same electron shell, hence the number of core electron shell is equal, so shielding effect is approximately the same. Increase in number of protons across the period causes increase in nuclear charge, so effective nuclear charge increase. Positive charges would attract the negative charges (electrons), hence increase in effective nuclear charge across period would increase electronegativity of elements.
Down the group, the ionisation energy decreases. The increase in nuclear charge down the group, due to increase in number of protons, is offset by the greater increase in shielding effect, due to increase in number of core electrons down the group. As a result, effective nuclear charge decreases, less positive charge to attract the negative charge of electrons, so electronegativity decrease down group.

(f) Physical Properties – Melting and boiling point
Across the period, the melting and boiling point of elements increase from elements of Groups I to IV, before decreasing from Groups V to 0. Metallic bonding exists within elements from Groups I to III, which exists as a metallic lattice of metal cations surrounded by a sea of electrons. Hence, strong metallic bonds exist within molecules, thus high amount of energy is required to break the metallic bonds, so melting and boiling points are high. Elements in Group IV exists as giant covalent compound with strong covalent bonds between atoms of compounds, hence high amount of energy is required to break the covalent bonds, hence high melting and boiling points. Elements from Groups V to VII exist as simple discrete covalent molecules with weak van der Waals forces of attraction between molecules, hence small amount of energy required to overcome these forces, boiling and melting points are low. Group 0 elements do not form bonds and hence have very low melting and boiling points.

(g) Physical Properties - Electrical conductivity
Across the period, elements in Group I to III are able to conduct electricity, while elements from Group IV to 0 is unable to conduct electricity. Metallic bonding exists within elements from Groups I to III, which exists as a metallic lattice of metal cations surrounded by a sea of electrons. They are good conductors of electricity due to the presence and movement of free, mobile delocalised electrons through the metal that can carry the electric charge. Elements in Group IV exists as giant covalent compound with strong covalent bonds between atoms of compounds All giant covalent compounds do not conduct electricity, except graphite. This is as all valence electrons of these compounds are used to form covalent bonds, hence there are no free electrons and electrical conduction does not occur. Elements from Groups V to VII exist as simple discrete covalent molecules with weak van der Waals forces of attraction between molecules. Simple covalent compounds do not conduct electricity, as they do not contain any mobile ions or electrons.

(h) Physical Properties - Properties of oxides formed
Across the period, elements in Group I and II form basic oxides, elements in Group III form amphoteric oxides, elements from Group IV to VII form acidic oxides. This is as elements in Group I and II tend to lose electrons to form ions for energy efficiency, while elements from Group IV to VII tend to gain electrons.

(i) Trends in Group I – Alkali Metals
The elements in group 1 are called the alkali metals. They belong to the left-hand column in the periodic table. They are very reactive.
1) Physical Properties of alkali metals • Low melting and boiling points • Very soft • Low densities • Reactive with water • Hydroxides and oxides dissolve in water to form alkaline solutions
2) Trends down group • Melting and boiling points decrease (Due to decrease in effective nuclear charge) • Densities increase (More significant increase in mass than atomic radius/volume) • Become softer • Become more reactive (Due to decrease in ionisation energy which makes it easier to remove the outermost electron from atom, and weaker metallic bonds, leading to lower activation energy, hence easier to form positive ions, more reactive)
3) Reactions with water
All of these metals react vigorously or even explosively with cold water. In each case, a solution of the metal hydroxide is produced together with hydrogen gas.
4) Reactions with oxygen gas
The alkali metals react with oxygen in the air very quickly to form the metal oxide. We say the metal tarnishes in air rapidly.

(j) Trends in Group VII – Halogens
The elements in group 7 are called the halogens. They belong to the column second from right in the periodic table. The halogens are all toxic, but this is a useful property. Chlorine is used to sterilise drinking water and water in swimming pools. Iodine is used in antiseptics to treat wounds.
1) Physical Properties of Halogens • Non metals • Low melting and boiling points • Brittle in solid form • Poor electrical and thermal conductivity • Coloured vapours • Diatomic molecules
2) Trends down the group • State at room temperature turns from gas to solid (Due to increase in strength of weak van der Waals forces which is proportionate to increase in number of electrons which pulls molecules more strongly together – Gas, Gas, Liquid, Solid, Solid) • Colour of solution gets darker (yellow, green, reddish orange, reddish brown, black) • Reactivity decreases down the group (Halogens at top displaces those at bottom due to increase in distance between nucleus and valence electrons) • Melting and boiling point increase (Due to increase in number of electrons, resulting in proportionate increase in van der Waals forces, hence more energy required)
3) Displacement reaction
Displacement reaction can occur in reactions involving 2 halogens, with the more reactive one (at top of table) displacing the less reactive one (bottom of table). This is oxidation.
4) Reaction with metals
They react with metals spontaneously to make ionic compounds in which the halide ions have a -1 charge.
5) Reaction with hydrogen gas
The halogens react with hydrogen gas to form hydrogen halides. The hydrogen halides dissolve in water to make acidic solutions.
6) Reaction with water
The more reactive halogens (fluorine, chlorine) readily dissolve in water to form the acids.

(k) Trends in Group 0 – Noble Gas
The elements in group 0 are called the noble gases. They belong to the right-hand column in the periodic table. The noble gases are all chemically unreactive which means they are inert. The highest occupied energy levels (the outer shells) of the atoms in Group 0 are full. This means that noble gas atoms have no tendency to lose or share electrons to form ionic compounds, or to share electrons to form covalent bonds with other atoms. Because their highest occupied energy levels are full, the noble gases are unreactive (inert) and exist as single atoms (they are monatomic).
1) Physical Properties of Halogens • Non metals • Very unreactive • Colourless • Monoatomic • Low boiling and melting point • Low densities
2) Trends down the group • Increase in melting and boiling points (stronger electrostatic forces of attraction) • Increase in density (increase in mass is more significant than increase in volume)

(l) Trends in Transition Metals
A transition metal is one which forms one or more stable ions which have incompletely filled d orbitals. The transition metals belong to the large block of elements between Groups II and III in the periodic table. They are commonly known as the d-block elements as most of these elements have partially filled d-orbitals.
1) Physical Properties in transition metals • High melting and boiling points, • High densities, • Hard, tough and strong. • Good conductors of heat, • Good conductors of electricity, • Malleable • Form coloured compounds
2) Chemical Properties in transition metals • Exhibit variable oxidation states (Due to partially filled d-orbitals) • Less reactive than alkali metals
3) Use of transition metals • Catalysts in many processes, such as manufacturing of ammonia • Structural materials
9) Chemical Calculations
(a) The Mole
One mole of substance is approximately 6.022 x 1023. This is also known as the Avogadro’s Constant, NA.
(b) Calculating number of moles
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(c) Relative Atomic Mass, Ar The relative atomic mass of an element is referenced to the carbon-12 atom. Most of the atomic mass in the periodic table are whole numbers as isotopic abundance of one of the major isotopes is approximately 100%.
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(d) Calculating Isotopic Abundance
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(e) Relative Molecular Mass, Mr
The relative molecular mass of a molecule is the average mass of one molecule of a substance when compared with 1/12 of the mass of one carbon-12 atom. In other words, it is the sum of all the relative atomic mass of all atoms in the molecule.
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(f) Relative Formula Mass, Mr
The relative formula mass of a molecule is the average mass of a formula unit of molecule when compared with 1/12 of the mass of one carbon-12 atom in ionic compounds. In other words, it is the sum of all the relative atomic mass of all atoms in the formula unit.
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(g) Percentage Composition of an element in compound
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(h) Percentage Composition of water in compound
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This is used to calculate the number of water of crystalliation in hydrated salts.
(i) Empirical Formula
The empirical formula of a compound is the formula of a compound in its simplest form, expressed in ratios of elements in the lowest terms.
To find empirical formula:
| |Element |Element |Element |
|Mass/100g | | | |
|No. of moles | | | |
|Mole Ratio | | | |

(j) Molecular Formula
The molecular formula of a compound is the actual unit formula of a compound, expressed in multiples of the empirical formula. To find molecular formula, divide the given molecular mass of molecular formula by that of the empirical formula.
(k) Avogadro’s Principle and Molar Volume
Avogadro’s Principle states that at any given temperature and pressure, the volume of a sample of any gas is proportional to the number of molecules in the sample.
At room temperature and pressure, 1 mole of any gas occupies a volume of 24 dm3.
At standard temperature and pressure, 1 mole of any gas occupies a volume of 22.4 dm3.
(l) Percentage Yield of Reaction
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(m) Percentage Purity of Sample
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(n) Concentration
The concentration of a sample refers to the amount of solute dissolved per unit volume of solvent. It can either be given in g/dm3 or mol/dm3 (M). To convert from mol/dm3 to g/dm3, simply divide the concentration in mol/dm3 by the relative molecular mass of compound.
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(l) Limiting Reagent
A limiting reagent is the chemical reagent that is present in the least number of moles according to the stoichiometric ratio. It places a limit on the amount of product that can be formed even if the other reactants are present in excess.
(m) Chemical Equations
Procedure for answering questions related to chemical equations:
1) Write a balanced chemical equation for the reaction.
2) Find out the embedded information in the reaction that helps to solve the problem.
3) If 2 set of data are given (2 reactants), find out the limiting reagent first and use the number of moles of the limiting factor to carry out the calculations.
4) Compare the stoichiometric ratio to obtain the number of moles of compound of interest
5) Calculate what is asked for in the question (Mass of product? Percentage yield?)

10) Air and Environment
(a) The Atmosphere and air
The Earth is surrounded by a layer of air called atmosphere. Air consists mainly of 79% nitrogen, 20% oxygen, 1% Noble gas, 0.03% carbon dioxide and varying amount of water vapour.
(b) Air Pollution and air pollutants
The presence of substances in the atmosphere that are harmful to living things and environment contributes to air pollution. These substances are called air pollutants – they come from both natural sources and human activities. Volcano eruptions and forest fires are examples of natural sources of air pollutants, while pollution from human activities are mainly due to burning of carbon containing fuels in motor vehicles and factories.
1) Sulfur dioxide
Large amounts of coal and petroleum are burnt arounf the world in power stations to generate electricity and in industries to provide energy. Both fuels contain sulphur as an impurity, which, during combustion, would oxidise to form sulphur dioxide. Diesel fuel in vehicles also contains a little amount of sulphur.
Sulfur dioxide irritates the eye and causes breathing problems. It enters leaves and affects plant growth. It is also the main cause of acid rain, as sulphur dioxide rises and reacts with oxygen and water vapour in air to form sulphuric acid, a source of acid rain.
To reduce sulphur dioxide emission in power stations, flue gas desulfurisation is adopted. Powdered limestone, which is basic calcium carbonate, is added to hot gases produced from burning of fuels at power station. The heat decomposes the limestone to give calcium oxide. The calcium oxide then removes sulphur dioxide as calcium sulfite. This method is cheap and removes 95% of sulphur, hence is also efficient. Also, cleaner fuels are used – fuels in Singapore now cannot contain more than 0.05% sulphur by weight. Power stations are now increasingly burning natural gas, while vehicles are encouraged to use compressed natural gas instead of diesel. Natural gas contains no sulphur and hence would not produce sulphur dioxide in motor vehicle, thus reducing sulphur dioxide emission.
2) Oxides of nitrogen
At high temperature, nitrogen and oxygen in air combine to form nitrogen monoxide. The nitrogen monoxide can also combine with more oxygen to become nitrogen dioxide. These reactions occur naturally in lightning. However, more significant sources to these oxides are the engines of vehicles, as well as burners in power stations, factories and incinerators.
These gases irritate and damage the lungs. Nitrogen dioxide is also a cause of acid rain, as it rises and reacts with oxygen and water vapour in air to form nitric acid which is a main contributor to acid rain. The gases also react with sunlight and other pollutants to form ozone. It further reacts with unburnt hydrocarbons to form petrochemical smog.
To reduce emission of these pollutants, cars have been fitted with catalytic converters. These remove about 95% of pollutants from exhaust gases. In the first half of the converter, oxides of nitrogen react with carbon monoxide as they pass through a catalyst to produce carbon dioxide and nitrogen gas, both of which are harmless. In the second half, air enters and oxidises unburnt hydrocarbon and carbon monoxide to form carbon dioxide and water, which are non polluting as well. These are then expelled through exhaust of car.
3) Carbon monoxide
Most of carbon dioxide comes from the incomplete combustion of carbon containing fuel in motor engines, which produce CO instead of carbon dioxide.
CO is dangerous to human health as it has the ability to bind irreversibly to haemoglobin (and has a higher affinity to haemoglobin than oxygen), which reduces the oxygen carrying capacity of the human body, and this can potentially result in death.
To reduce emission of CO, cars have been fitted with catalytic converters. These remove about 95% of pollutants from exhaust gases. In the first half of the converter, oxides of nitrogen react with carbon monoxide as they pass through a catalyst to produce carbon dioxide and nitrogen gas, both of which are harmless. In the second half, air enters and oxidises unburnt hydrocarbon and carbon monoxide to form carbon dioxide and water, which are non polluting as well. These are then expelled through exhaust of car.

4) Unburnt Hydrocarbons
Unburnt hydrocarbons come mainly from hydrocarbons in fuel that has not completely undergone combustion in vehicle engines (incomplete combustion).
These hydrocarbons react with sunlight and oxides of nitrogen, as well as other pollutants, to form petrochemical smog.
To reduce emission of this pollutant, cars have been fitted with catalytic converters. These remove about 95% of pollutants from exhaust gases. In the first half of the converter, oxides of nitrogen react with carbon monoxide as they pass through a catalyst to produce carbon dioxide and nitrogen gas, both of which are harmless. In the second half, air enters and oxidises unburnt hydrocarbon and carbon monoxide to form carbon dioxide and water, which are non polluting as well. These are then expelled through exhaust of car.
5) Methane
Methane is naturally present in atmosphere from bacterial decay of vegetation. However, amount of methane in air is rising due to human activities such as agriculture. Cows and other farm animals, as well as rice fields, produce methane. Methane also leaks from pipelines while being used as natural gas fuel.
Methane contributes to greenhouse effect as it is a greenhouse gas and traps heat, which eventually results in global warming.
6) Petrochemical smog
Petrochemical smog is formed when unburnt hydrocarbons react with oxides of nitrogen and other pollutants. The mixture of substances in this pollutant is more harmful than its separate components. It burns the eye and is dangerous to people with breathing or heat problems. It is poisonous to plants, and damages materials like rubber. To reduce petrochemical smog emission, reduce emission of unburnt hydrocarbon and NOx.
(c) Acid Rain
Acid rain is any rainfall that has an acidity level beyond what is expected in non polluted rainfall, usually of an acid level of 1 to 4.
Acid rain has 2 main components, namely sulfuric acid and nitric acid. Both acids are formed when their (di) oxides react with oxygen and water vapour in air to produce the respective acids. These acids dissolve in rainwater, making it acidic. Wind can also carry the air pollutants over long distances before they dissolve in rainwater, which makes the impact of acid rain all the more dangerous and harmful.
Below are some harmful effects of acid rain: • Makes soil too acidic for plant to grow, hence decreasing plant life • Corrodes buildings made of limestone and cement • Attacks metals and hence corrode metal structures instantly • Damage trees • Acidifies the water in water bodies, killing marine wildlife
As such, solutions must be employed to reduce effect of acid rain: • Burn fuels which contain little or no sulfur, such as natural gas • Reduce amount of pollutants in air (catalytic converters, flue gas desulfurisation) • Neutralise acids in lakes and soil
(d) Greenhouse effect
The greenhouse effect is a process by which thermal radiation from a planetary surface is absorbed by atmospheric greenhouse gases, and is re-radiated in all directions, as a result trapping heat in the atmosphere and preventing it from escaping into space. Main elements of greenhouse gases are carbon dioxide, methane and water vapour. The increase in greenhouse gases such as carbon dioxide and methane is causing a great increase in the greenhouse effect. This increase in greenhouse effect causes heat to be trapped in the atmosphere, which warms the Earth surface and causes Global Warming.
(e) Global warming
Global warming is a direct result of increase in greenhouse effect. It is defined as the continuing rise in the average temperature of Earth's atmosphere and oceans. Consequences of global warming are as follows: • Expansion of seawater and melting of polar ice caps/glaciers will cause drastic rise in sea levels, causing flooding of low lying land and disappearance of many islands. • Big changes in the global climate and unpredictable weather changes. • Increased occurrence of natural disasters such as cyclones, typhoons and hurricanes. • Loss of wildlife and species due to inability to adapt to change in environment
To solve the problem of global warming: • Generate electricity via cleaner sources such as wind • International agreements to reduce carbon dioxide emission (Kyoto Protocol)

11) Electrochemistry
(a) Redox reactions
Redox reactions are reactions involving a change in oxidation states of elements in reactants during the reaction. Alternatively, redox reactions can also be in terms of electron transfer, as well as gain or loss of oxygen or hydrogen.
Redox reaction, as the name suggests, involved 2 reactions, namely oxidation and reaction. Oxidation reaction involves the increase in the oxidation number of elements in the reactants during the reaction, while reduction reaction involved the decrease in oxidation number. Oxidation reaction also refers to the loss of electron in reactants, gain of oxygen and the loss of hydrogen. Reduction reaction refers to the gain in electron, loss of oxygen and gain in hydrogen.
When answering questions involving redox reactions, answer it in terms of change in oxidation state. For example, X is oxidised/reduced as the oxidation number of X increases/ decrease from (ON before reaction) in (reagent) to (ON after reaction) in (product).
(b) Rules to determine ON
1) ON of an element ( 0
2) ON of a simple ion ( Charge and magnitude of ion
3) ON of a compound ( 0
4) ON of a complex ion ( Charge and magnitude of ion
To find out the ON of an element in the molecule, write out the charge and magnitude of other ions. In a compound, ON is 0. Hence, simply subtract 0 by all other ONs of other ions.
(c) Redox Reagents
1) Oxidising Agents
Oxidising agents are compounds that oxidise other compounds, while it is reduced in the reaction. Below are 2 examples of oxidising agents:
Potassium Dichromate (VI), K2Cr2O7
When this is used as the oxidising agent, it would be reduced, hence it would turn from orange to green in colour. The half equation of this reaction is below:
Cr2O72- + 14H+ +6e ( 2Cr3+ + 7H2O
Potassium Manganate (VII), KMnO4
When this is used as the oxidising agent, it would be reduced, hence it would decolourise and turn from purple to colourless in colour. The half equation of this reaction is below:
MnO4- + 8H+ +5e ( Mn2+ + 4H2O

2) Reducing agents
Reducing agents are compounds that reduce other compounds, while it is oxidised in the reaction. Below is an example of reducing agents:
Potassium Iodide, KI
When this is used as the reducing agent, it would be oxidised, hence it would turn from colourless to brown solution or black precipitate. The half equation of this reaction is below:
2I- ( I2 + 2e
(d) Half Equations
To balance 2 half equations to become a redox reaction, the following steps are taken:
1) Find the LCM of the 2 numbers of ‘electrons’
2) Combine the 2 equations together by adding them and cancelling out the electrons on both sides
3) Cancel the repeating compounds that are found on both sides of the resultant equation
(e) Electrolysis
Ionic compounds in the solid state have cation and anions arranged in a lattice structure. The melting and boiling points of the ionic compounds are high, owing to the strong electrostatic forces of attraction between ions in the compound. The ions are pulled apart as a result of heat energy applied when the solid is melted or with the help of solvent molecules when the solid is dissolved. When an ionic solid is melted or dissolved in water, some of the substances dissociate into freely moving charged particles called ions. This is known as ionisation. The resultant mobile ions allow the conduction of electricity, and this is used in the process of electrolysis, which consists of:
1) A DC battery
This supplies the circuit with electricity required.
2) An anode (positive electrode) and a cathode (negative electrode)
These are connected to the battery. Ideally, they should be inert and do not react with the electrolyte nor undergo redox reactions within the cell.
3) Electrolyte
This is the molten or aqueous liquid so that ions can move and conduct electricity.
When an electric current is passed through an electrolyte, the mobile ions lose their random movements. Instead, the cations are attracted to the cathode and anions to the anode. When ions reach the electrodes, oxidation and reduction occurs at the anode and cathode respectively. The ions are said to be discharged.
Anode reaction
At the anode, anions lose their excess electrons to the surface of electrodes and become neutral atoms/molecules. They are said to be oxidised.
Cathode reaction
At the cathode, the cations accept excess electrons on the surface of the cathode and discharge. They are said to be reduced.
(f) Factors affecting electrolysis
1) State of the electrolyte
Electrolysis of molten salts is predictable since usually only 1 cation and 1 anion is present, that is, that of the molten electrolyte.
However, products of the electrolysis of aqueous solution are unpredictable as when water is present, water may decompose in place of the molten electrolyte. In aqueous solutions, the choice of the molecule to be decomposed depends on the factors below.
2) Position of ions in the electrochemical series
When there is more than 1 cation present, the ions of metals lower down in the electrochemical series are discharged in preference to those higher up. For aqueous solutions of salts above hydrogen ion in the electrochemical series, hydrogen gas is produced instead of the salt of cation in the cathode reaction:
2H+ + 2e ( H2
When there is more than 1 anion present, the anion that is preferably discharged is below. For anions of diatomic molecules, these ions would be discharged. From aqueous solutions of nitrates and sulphates during electrolysis, oxygen gas is produced in the anode reaction:
4OH- ( 2H2O + 4e + O2
The electrochemical series for both cations and anions is as follows:
|Cation | |Anion |
|Silver |Easiest to discharge |Hydroxide |
|Copper | |Iodine |
|Hydrogen |Dividing point (cation) |Bromine |
|Lead (II) | |Chloride |
|Tin (II) | |Nitrate |
|Iron (II) | |Sulphate |
|Zinc (II) | | |
|Aluminium | | |
|Magnesium | | |
|Sodium | | |
|Calcium | | |
|Potassium |Most difficult to discharge | |

3) Concentration of aqueous solution
The ease of discharge depends on the electrode potential which in turn may depend on the concentration of solute in solution. If a cation is in very high concentration, it may be discharged in preference to one below it in electrochemical series. Similarly, if an anion is in very high concentration, it may be discharged in preference to one which would normally be discharged first if their concentrations were similar.
4) Nature of electrode
The choice of electrodes depends on the purpose the electrolysis. If electrodes are to play no part in the electrolysis, an inert electrode such as graphite or Pt is used. In this case, the purpose of the electrodes is simple to serve as surface of electron transfer.
In some case, electrodes are deliberately chosen to take part in redox reactions. The presence of these electrolytes may make the products of electrolysis different.
5) Overall summary
In dilute solution, for the cation, hydrogen ion would be discharged if metal ion is above hydrogen in the electrochemical series. If metal ion is below hydrogen in the electrochemical series, metal ion would be discharged. For anion, the hydroxide ion would be discharged regardless of the electrochemical series.
In concentrated solution, for the cation, hydrogen ion is discharged if metal ion is above below hydrogen in the electrochemical series. If the cation is just above the hydrogen ion, such as tin and lead (II), the metal ion may be discharged. If the cation is below hydrogen ion, the metal ion would be discharged. For the anion, halides would be discharged. If oxoanions are present, hydroxide ions would be discharged.
(g) Extraction of metals – applications of electrolysis
1) Down’s Process
Sodium is produced in Down’s process by electrolysing molten sodium chloride.
|Electrolyte |Anode Reaction |Cathode Reaction |
|Molten NaCl |Graphite |Iron |
| |2Cl- ( Cl2 + 2e |Na+ + e ( Na |

2) Extraction of aluminium – Hall-Heroult Process
Aluminium is extracted from its ore bauxite. Bauxite melts at over 2000 degree Celsius, making the electrolysis extremely costly. However, bauxite dissolves in molten cryolite which melts at over 1000 degree Celsius to make electrolysis more money- efficient.
|Electrolyte |Anode Reaction |Cathode Reaction |
|Molten Al2O3 dissolved in cryolite |Graphite |Graphite |
| |2O2- ( O2 + 4e |Al3+ + 3e ( Al |

3) Electrolytic refining of copper
The extraction of copper uses an electrolytic cell with the impure copper anode as the anode and a pure copper plate as the cathode. The electrolyte may be any soluble copper salt.
|Electrolyte |Anode Reaction |Cathode Reaction |
|Aqueous Copper (II) Sulfate |Impure copper |Pure copper |
| |Cu ( CU2+ + 2e |CU2+ + 2e ( Cu |

4) Electroplating of objects (silver)
Electroplating is the use of electrochemical means to coat the surface of an object with a metal. The metals used in electroplating are using copper, nickel, chromium, silver or gold. As current is passed through the cell, the metal dissolves into the electrolyte. The cations in solution migrate to the spoon and is plated onto the spoon/
|Electrolyte |Anode Reaction |Cathode Reaction |
|Solution of metal ion |Metal to electroplate object |Object to be electroplated |
| |M ( M+ + e |M+ + e ( M |

5) Industrial Electrolysis of Brine (Kellner Solvay Cell)
The industrial electrolysis of brine produces NaOH, chlorine gas and hydrogen gas. The setup consist of a tank containing brine. The cathode is a layer of mercury flowing in on the base of the tank is same direction as brine. The Na atoms reduced at mercury electrode forms sodium amalgam (sodium solution dissolved in mercury), which is continuosly drained into a separate chamber and treated with water. A series of graphite rods act as the anode. These rods are dipped into brine solution, oxidation of chloride anions takes place forming chlorine.
|Electrolyte |Anode Reaction |Cathode Reaction |
|Concentrated NaCl (Brine) |Graphite |Mercury |
| |2Cl- ( Cl2 + e |Na+ + e ( Na (Hg) |

12) Thermochemistry
(a) Endothermic Process – Bond Breaking
Bond breaking is a process that requires an energy input. Energy is required to break a bond between 2 atoms in a covalent molecule. The energy is used to overcome the electrostatic forces of attraction between 2 atoms or ions so that they can be separated. The Bond Dissociation Energy is the energy required to break one mole of bonds of the same type in the gaseous state under standard room temperature and pressure.
For ionic compounds, the strength of the electrostatic forces of attraction between cations and anions are measured by the lattice energy. The lattice energy of a substance is the energy released into the surroundings when 1 mole of an ionic compounds if formed from its constituent gaseous ions. The lattice energy is proportional to the product of the ionic charges of both ions and inversely proportional to the ionic sizes.
(b) Exothermic Process – Bond Formation
When covalent bonds are formed between atoms, energy is given out. The amount of energy required to break the bond must be the same amount of energy released during bond formation.
(c) Calculating enthalpy change in reaction
1) Write down equation for reaction
2) Draw Lewis structure for reactants and products
3) Calculate energy required to break bonds of reactants
4) Calculate energy required for bond formation of products
5) Calculate enthalpy change by subtracting energy required from energy released
(d) Energy Profile Diagrams
The reaction energy profile diagram gives a representation of how the internal energy of the species involved in reaction changes as reactants change into products.
1) Activation energy
The activation energy is the energy barrier that reactant molecules are required to overcome before a reaction is started. For a reaction occurring in a fixed temperature, the activation energy is also fixed. Activation energy can also be defined as the minimum amount of energy reactant molecules must possess for a chemical reaction to occur.
2) Exothermic Reaction – energy profile diagrams
In this, the energy for reactant is higher than that of products since energy are released. In such a reaction, overall enthalpy change is negative. One obvious sign would be heat being liberated during the reaction (system warms). Neutralisation is an exothermic reaction.
3) Endothermic Reaction – energy profile diagrams
In this, the energy for the reactant is lower than that of products since net energy is required. In such a reaction, overall enthalpy change is positive. One obvious sign would be heat being adsorbed during the reaction (system cools).
4) Bond Dissociation Energy of covalent bonds
The bond dissociation energy is proportional to the strength of covalent bonds, which in turn is affected by the number of bonds (single, double or triple) as well as the bond length, determined by the atomic radii. The more the number of bonds, and the shorter the bond length, the higher the BDE.
13) Rates of reaction

(a) Measuring rate of reaction

1) Measure time taken for complete reaction to occur
Most suitable for reactions in which visible observation can be made so as to effectively determine if complete reaction has occured.

2) Measure mass loss of system as reaction goes by
End point - mass has reached stable equilibrium/ visible observation can be made (eg. effervescence stopped). Only suitable for reactions in which gas is produced so that gas can escape. In this case, cotton wool is required to keep other products/reactants from escaping

3) Measuring volume of as produced
End point - visible observation can be made (eg. effervescence stopped), gas syringe nozzle stops moving. Only suitable for reactions in which gas is produced. In this case, system has to be closed except for gas syringe where gas produced would move so gas won’t escape.

(b) Collision theory
The collision theory states that products are formed only under certain conditions, or when reactant molecules have met certain requirements for reactions to take place. Hence, the frequency of effective collision depends on the factors below:

1) Frequency of collision between reactant molecules
2) Percentage of reactant molecules possessing energy higher than activation energy
3) Probability of correct collision geometry

(c) Factors affecting rate of reaction

1) Temperature
Increasing temperature of a reaction system would cause reactant molecules to move in higher velocities within the system, as a result colliding more frequently with one another and increasing the frequency of collision. Also, due to higher kinetic energy, there is a greater percentage of reactant molecules that possess energy higher than activation energy. As a result of increase in frequency of collision between reactant molecules and percentage of reactant molecules possessing energy higher than activation energy, frequency of effective collision increase hence increasing rate of reaction.

To study this effect more closely, the Maxwell- Boltzmann Distribution of Energy is used.

2) Concentration of solution
For 2 substances to react, particles of 1 reactant must collide with that of another for reaction to occur. A more concentrated solution presents more particles per unit volume of solution, allowing for more frequent collision between reactant particles leading to higher frequency of effective collision, hence increasing rate of reaction.

3) Pressure for gases
When confined to smaller volume, gaseous particles would collide more frequently with one another, thus increasing frequency of effective collision and rate of reaction. This is achieved by applying higher pressure on the reaction.

4) Size of particles
Small pieces of solid have larger total surface area than those in large pieces of the same mass. Hence, with decrease in particle size, total surface area for reactant to collide with would be larger, thus allowing for more frequent collision between reactant particles leading to higher frequency of effective collision, hence increasing rate of reaction.

5) Catalyst
A catalyst is a substance that alters the rate of a chemical reaction while it remains unchanged chemically at the end of the reaction. Catalysts work by proving a lower energy pathway for reaction to occur (lowering activation energy so that there is a greater percentage of reactant molecules that possess energy higher than lowered activation energy), as well as providing a platform for collision between reactant particles to occur. Properties of catalysts include the ability to regenerate and be used again for another chemical reaction, and that a minute amount of catalyst can greatly speed up the rate of a chemical reaction. Poisoning of catalysts may occur when poisons (atoms or molecules containing lone pair of electrons) is adsorbed on catalyst more strongly than reactants, thus effectively preventing reactants’ access to catalytic sites.
Homogenous catalyst
Homogenous catalysts work by allowing reaction to proceed via different mechanism by forming an intermediate with one of the reactants, essentially lowering the activation energy required by reactants before reaction occurs.

Heterogenous catalyst
Homogenous catalysts work by providing sites for the reactants to meet and/or enable bonds to be broken and reformed at lower energy levels, essentially increasing probability of correct collision geometry, thus increasing rate of reaction.

Transition metals as good catalyst
Transition metals are good catalysts due to the availability of empty d –orbitals and the ability to exhibit variable oxidation states.

Enzymes as biological catalysts
Enzymes are biological catalysts in living systems. They are effective in minute amounts, operate within a small temperature and pH range, and are highly specific.

14) Chemical Equilibrium
(a) Reversible reactions and dynamic equilibrium
Reversible reactions are reactions that take place in both directions. Only such reactions can attain dynamic equilibrium. A dynamic equilibrium exists in a system which the forward and reverse reactions are taking place at the same rate, and thus the concentration of each species of reactant and product remains constant. The position of a chemical equilibrium is the relative ratio of the products to reactants, depending on the activation energy of the forward and reverse reactions. Hence, the position will affect the yield of the chemical reaction. There are 2 types of equilibria:
1) Homogenous equilibria
This is for any reversible system with its products and reactants in the same phase/state.
2) Heterogenous equilibria
This is for any reversible reactions with its products and reactants in different phases/states.
Below are some characteristics of dynamic equilibrium:
1) Forward and reverse reactions are occurring at the same rate.
2) Concentrations of reactants and products remain unchanged over time
3) Can be obtained from either direction
4) Can only be attained in a close system
(b) Le Chatelier’s Principle
This principle states that if an external stress, such as pressure is applied to a system at equilibrium, the system will adjust itself (its position) to partially offset the stress.
(c) Factors affecting system in dynamic equilibrium
Changes in reaction might affect both the rate and yield of reaction.
1) Concentration
When reactant concentration is increased, reaction will shift right to produce more products, hence increasing rate of forward reaction while reverse reaction remains constant. When reactant concentration is decreased, reaction will shift left to produce more reactants, hence increasing rate of reverse reaction while forward reaction remains constant. This applies when product concentration is altered as well. Such changes are due to the change in rate of collision of particles involved in reaction as a result of the change in concentration.
| |Equilibrium Position |Equilibrium Concentration |
|Reactant concentration increase |Shift right, produce more products |Product concentration increases, reactant|
| | |concentration decreases |
|Reactant concentration decrease |Shift left, produce more reactants |Reactant concentration increases, product|
| | |concentration decreases |
|Product concentration increase |Shift left, produce more reactants |Reactant concentration increases, product|
| | |concentration decreases |
|Product concentration decrease |Shift right, produce more products |Product concentration increases, reactant|
| | |concentration decreases |

2) Pressure (gaseous reactants/products only)
Pressure would increase the rate of both forward and reverse reactions. However, when pressure of volume is increased, equilibrium will shift to side with smaller total molar volume. When pressure of volume is decreased, equilibrium will shift to side with larger total molar volume. When there is no difference in molar volume of reactants and products, changes in pressure would not shift the position of the equilibrium, hence the yield stays the same while rate of reaction increases for both forward and reverse reactions. This is due to the change in rate of collision of particles involved in reaction when pressure changes, resulting in change in rate of effective collision. Hence, increase in pressure would increase rate of reaction, while decrease in pressure would decrease rate of reaction. However, changes in yield would depend on the molar volume of reactants and products.
| |Forward: Higher molar volume |Forward: Lower molar volume |
|Pressure |[Reactants] |[Products] |[Reactants] |[Products] |
|Increase |Decrease |Increase |Increase |Decrease |
|Decrease |Increase |Decrease |Decrease |Increase |

3) Temperature
An increase in temperature would increase the rate of both forward and reverse reactions, while decrease in pressure would decrease the rate of both reactions. This is as it changes the proportion of molecules possessing energy levels equal to or greater than the activation energy of both forward and reverse reactions.
A change in temperature will also change the equilibrium position depending on enthalpy change of reaction. Increasing the temperature of system would favour the endothermic reaction, while decreasing the temperature of system would favour the exothermic reaction.
|Reaction |Change in temp |[Reactant] |[Product] |
|Exothermic |Increase |Increase |Decrease |
| |Decrease |Decrease |Increase |
|Endothermic |Increase |Decrease |Increase |
| |Decrease |Increase |Decrease |

(d) Haber Process – Manufacturing of Ammonia
In this process, nitrogen and hydrogen gases are mixed together in the molar ratio of 1:3. The mixture is passed over an iron catalyst. This is a reversible and exothermic reaction. Conditions for the reaction are as such:
1) Temperature ( 450 degree Celsius
The Haber process is carried out at a moderately high temperature. For highest yield of ammonia product, the temperature of system should be kept as low as possible, which favours the yield of products in exothermic reactions as it decreases the proportion of molecules possessing energy equal to or higher than activation energy, hence according to LCP, equilibrium shift to right to produce more products. However, low temperature would mean a very slow rate of reaction, which is not efficient. Hence, 450 degree Celsius is used.
2) Pressure ( 200 atm
The Haber process is carried out at high pressure. For highest yield of ammonia product, the pressure of system should be kept as high as possible. This is as the molar volume of reactant to product is 2:1, thus increase in pressure would result in greater collision of reactant molecule according to LCP, so equilibrium shifts right, hence increasing product yield. However, higher pressures cannot be used as it is very expensive and not cost efficient.
15) States of Matter
(a) Kinetic Theory of Particulate Matters
[pic]
(b) Change of state
Below show processes whereby matter changes from one state to another, usually involving the addition or removal of heat energy.
[pic]
1) Melting
Melting is a process whereby state of matter change from solid to liquid. Heating imparts kinetic energy to particles of the solid. Particles vibrate more vigourously about their fixed positions, and gain rotational and translational motion. Particles gain enough energy to overcome their forces of attraction, and move out of their lattice arrangement, move freely within bulk of the matter.

2) Boiling

Boiling is a process whereby state of matter change from liquid to gas. Heating imparts kinetic energy to particles of the liquid. Particles of liquid gain more kinetic energy (ie. move at higher velocities). Particles gain enough energy to overcome the forces of attraction between the already mobile particles, and move further apart and finally gain enough energy to break out of the liquid surface.

3) Condensation

Condensation is a process whereby state of matter change from gas to liquid. Cooling removes kinetic energy from particles of the gas, and overall inter-particle distance decreases. Particles move more slowly as the lose kinetic energy. Particles becomes closer together as they collide with one another and become associated closely (as they do not have enough k.e. to move away).

4) Freezing

Freezing is a process whereby state of matter change from liquid to solid. Cooling removes kinetic energy from particles of the liquid. Particles move more slowly as the lose kinetic energy. Overall, inter-particle distance decreases, particles become more ordered in their arrangement in space. Particles becomes closer together as they collide with one another and become closely associated (as they do not have enough k.e. to move away). As more heat energy is lost, the particle translational movement slows to a stop.

5) Evaporation v/s boiling

Evaporation is a surface phenomenon (some molecules near the surface have enough k.e. to escape). When the container is closed, equilibrium may be reached (no. molecules escaping out of the surface is equal to the no. of molecules re-entering liquid bulk). We say the vapour is saturated.

When boiling occurs, the atmospheric pressure is equal to the saturated vapour pressure. Bubbles will form as the pressure of the vapour (inside the bubble) is equal to atmospheric pressure.

(c) Ideal Gas

An ideal gas model is a theoretical model based on Newtonian mechanics. It states that pressure of the gas arises from the collisions of particles with walls of container/vessel. Also, energy of the particles is measured in terms of the absolute temperature of gas.

1) Assumptions

• There is a large number of particles at all speeds and in random motion • Average kinetic energy of particles is proportional to absolute temperature only • Particles are of negligible size when compared to volume of vessel • Particles have negligible attractive forces between particles except at instance of collision • Collision of particles is perfectly elastic

2) Conversions

Volume ( 1 m3 = 1000 dm3 = 1000000cm3
Pressure ( 1 atm = 101325 Pa = 760 mmHg
Temperature ( X Kelvin = X degree Celsius + 273

(d) Gas Laws

1) Boyle’s Law
Boyle’s law states that at a constant temperature, the volume of any quantity of gas is inversely proportional to its volume.
[pic]

2) Charles’ Law
Charles’ Law states that at a constant pressure, the volume of any quantity of gas is proportional to its absolute temperature
[pic]

3) Gay-Lussac’s Law
Gay-Lussac’s Law states that when gases react with each other at a constant temperature and pressure, they combine in volumes that are related to each other as ratios of small whole numbers.

4) Avogadro’s Principle
Avogadro’s Principle states that equal volumes of different gases at the same temperature and pressure contain equal numbers of atoms or molecules.

5) Ideal Gas Equation

[pic]
T = Temperature in K n = no. of moles
For R = 8.31451/0.0820578,
P = Pressure in Pa/atm,
V = volume in m3/L

[pic]

(e) Real Gases

1) Real gases v/s Ideal gases

|Ideal Gas |Real Gas |
|Point-masses with no volume |Have a definite volume |
|No attractive forces between them except during instances of |Have inter-particle forces of attraction |
|collision | |
|Perfectly elastic collisions |Inelastic collision due to inter-particle forces of attraction |

2) Deviation of real gas from ideal gas
Real gases deviate significantly from ideal gases at high pressure and low temperature, due to the failure of two of the assumptions of the kinetic-molecular theory.

Molecular size
The ideal gas model assumes that particles are of negligible size and hence volume of gas is negligible when compared to volume of vessel. This is not valid for real gases which have a definite size. Thus, real gases are less compressible than ideal gases, since gas particles cannot be forced into a large volume of space occupied by other gas particles. This is more significant at high pressure, where volume becomes smaller. The larger the molecular size, the greater the deviation from ideality.

Intermolecular forces of attraction
The model also assumes that particles have negligible intermolecular forces of attraction between particles except at instance of collision. This is not valid for real gases as there are always intermolecular forces of attraction between gas particles, especially at high pressure and low temperature. At high pressure, gas volume is small, and gas particles are forced close enough for intermolecular forces of attraction to operate. At low temperature, average kinetic energy in particles is not high enough to overcome forces of attraction between them, hence particles are drawn more closely together, resulting in them occupying smaller volume than ideal. Hence, the stronger the intermolecular forces of attraction between gas particles, the greater the deviation from ideality.
16) Organic Chemistry

(a) Reasons for great variety of organic compounds

1) Carbon atom’s ability to form chains with other carbon atoms due to catenation
2) Valency of carbon – each carbon can form 4 covalent bonds
3) Different arrangement of same atoms produces different compounds (isomerism)

(b) General properties of OC

1) Structure and Bonding
Organic compounds exist as simple, discrete covalent molecules with weak van der Waals forces of attraction between most molecules. Due to low electronegativity difference between C and H atoms, organic compounds are largely non polar.

2) Solubility in water
Being largely non polar, most organic compounds are insoluble in water due to the insufficient interaction between the weak van der Waals forces of attraction and the relatively stronger hydrogen bonding between water molecules. Exceptions are polar compounds containing electronegative elements.

3) Solubility in non polar organic solvents
Due to similar interactions between organic compounds, organic compounds are soluble in non polar organic solvents since strong interactions exist between compounds.

4) Low melting and boiling point
Weak van der Waals forces of attraction exist between molecules, hence low amount of energy is required to overcome these forces, so low melting and boiling point

5) Thermal Instability
They are thermally unstable due to weak van der Waals forces, and decompose into simpler molecules when heated to temperatures above 500 degree Celsius in absence of oxygen.

6) Reactivity
Organic compounds tend to be less reactive than organic compounds, hence leading to slower rates of organic reactions, which is sped up via use of catalyst or heating.

7) Flammability
Organic compounds burn in oxygen to produce carbon dioxide and water, or carbon monoxide and water in limited oxygen supply (incomplete combustion).

(c) Formulae of OC

1) Empirical Formula
This is the simplest formula of the compound which indicates the relative numbers of each kind of atom in a molecule.

2) Molecular Formula
This indicates the actual number of each kind of atom in a molecule of a substance.

3) Displayed Formula/Full Structural Formula
It shows the covalent bonds between atoms in an organic molecule, and hence the arrangement of atoms.

4) Structural Formula/ Condensed Formula
It is the condensed displayed formula, in which the alkyl and function groups are arranged in the correct sequence. Brackets are used to indicate the atom(s) is attached to the carbon preceding the bracket.

(d) Definitions

1) Functional group
A functional group is an atom or group of atoms on an organic compound that gives common characteristics to a homologous series of organic compounds. It determines the main chemical properties of the series.

2) Homologous series
A homologous series is a family of organic compounds which follows a regular structural pattern and has a general molecular formula, in which each successive member differs in composition by the addition of a CH2 group.

3) Isomerism
Isomerism refers to the phenomenon whereby certain compounds, possessing the same molecular formula, exist in different forms due to different structural formula and hence different arrangement of atoms.

(e) Isomers

1) Stereoisomers

2) Structural isomers
Isomers which have the same molecular formula but different structural formula and hence different arrangement of atoms.

Chain Isomers
Chain isomers are different in their longest carbon chain ( Straight v/s branched

Positional Isomers
Position isomers differ in position of alkyl/functional groups ( Propan-1-ol and Propan-2-ol

Functional Group Isomers
Functional group isomers differ in functional groups present ( Ethanol, methoxymethane

(f) Breaking of covalent bonds

1) Homolytic fission
This occurs when the covalent bond is broken equally, such that each of bonding atom leaves with 1 of the bonding electron.

2) Heterolytic fission
This occurs when covalent bond is broken such that 1 of bonding atom carries both bonding electrons and becomes anion and other atom has no bonding electrons, becoming a cation.
(g) Classification of reactants in OR

1) Nucleophile
It is an atom, ion or molecule that has at least 1 lone pair of electrons, which are attracted to electron deficient sites. It is an electron pair donor and forms a new covalent bond by providing an electron pair to a Lewis acid. Examples include fluorine and chlorine anion.

Electron rich sites include atoms with a partial negative charge due to covalent bond with less electronegative atom, double bonds (or any degree of unsaturation), lone pairs and anions.

2) Electrophile
It is an electron deficient atom, ion or molecule that is attracted to regions of negative charge/ electron rich sits. It is a Lewis acid (electron pair acceptor) that forms a new covalent bond by accepting an electron pair from a Lewis acid. Examples include hydrogen ion.

Electron deficient sites include atoms with a partial positive charge due to covalent bond with more electronegative atom, and cations.

(h) Classification of OR types

Organic reactions fall into a few distinct types, of which other than combustion, most other reactions only involve part of molecule with functional groups, which contain reactive sites that are attacked by incoming chemical species. A reactive site is a region of higher or lower electron density compared to rest of organic molecule.

1) Substitution
This occurs when an atom/group of atoms is substituted by another in the organic molecule, leaving at least 2 products.

2) Addition
This occurs when an atom/group of atoms is inserted into an organic molecule to form a single substance, usually for unsaturated functional group.

3) Elimination
This occurs when a small molecule is removed from larger organic molecule.

4) Condensation
This involves the addition of 2 organic molecules followed by elimination a small molecule.

5) Hydrolysis
Hydrolysed reactions involve water, often catalysed by dilute acids or alkalis.

6) Redox Reactions
Oxidation reactions are those involving the addition of oxygen or removal of hydrogen to/from the molecule. Reduction reactions are those involving the removal of oxygen or addition of hydrogen from/to the molecule.

(i) Alkanes

Alkanes are saturated hydrocarbon and contain maximum number of hydrogen atoms per carbon atom. They are tetrahedral with respect to each carbon molecule.

1) Physical Properties

Melting and Boiling Points
Melting and boiling points increase with length of carbon chain. At room temperature and pressure, alkanes with 1 to 4 C atoms are gases, 5 to 16 are liquid, 17 onwards are solid. With increase in number of C atoms, electrons increase, hence weak van der Waals forces of attraction increase, more energy required to overcome forces of attraction.

Viscosity
Viscosity of liquid alkanes increases with number of C atoms. With increase in number of C atoms, electrons increase, hence weak van der Waals forces of attraction increase, thus decrease in surface area, hence increase in viscosity.

Solubility in water
Alkanes are insoluble in water due to great difference in intermolecular forces of attraction between water molecules (hydrogen bonding) and alkane molecules (weak van der Waals forces). Hence, insufficient interaction occurs, so alkanes are insoluble in water.

Density
Alkanes are less dense than water due to weak van der Waals forces which does not pull molecules together closely enough, hence volume increases and density decreases.

Flammability
Flammability of alkanes decreases as number of C atoms increase, due to the need to supply enough energy to break the numerous C-C bonds present in long chain alkane.

2) Chemical Properties and Organic Reactions

Alkanes have low reactivity due to non polar nature of C-H bonds, so alkanes have no reactive sites for electrophiles and nucleophiles to attack.

Combustion
Alkanes react with oxygen during complete combustion to produce carbon dioxide and water. With limited oxygen supply, incomplete combustion occurs and carbon monoxide is produced instead of carbon dioxide. Due to decrease in flammability with increase in number of C atoms in alkanes, incomplete combustion is more likely to occur with higher number of C atoms.

Free Radical Substitution
Alkanes react with chlorine and bromine easily in presence of strong sunlight to form halogenoalkanes. Since alkanes are saturated and non polar, they can only undergo substitution reactions involving free radicals. Radicals are reactive chemical species with unpaired electrons. Alkanes are non polar, hence only very reactive chemical species which can attract bonding electrons can break C-H bonds in alkanes. The presence of light is necessary for the homolytic cleavage of the Halogen-Halogen bond.

During the reaction, greenish yellow gas or reddish brown liquid decolourises rapidly, and formation of white fumes of HCl or HBr is observed.

This is split into 3 stages, namely initiation, propagation and termination.

Problems associated with this are that numerous products and impurities are formed from the uncontrolled pairing up of free radicals. Also, in the presence of excess halogen, further substation can occur resulting in multi substituted products causing impurities.

Cracking
Cracking refers to the process where C-C bonds in long chain alkane molecules are broken to produce smaller molecules of alkanes and alkenes. Cracking is a very important process in the petrochemical industry to break down the large hydrocarbons into smaller ones which burns more easily as fuel. Also, cracking produces branched chain alkanes which provide petrol with higher octane rating.

Cracking can either produce: • Shorter chain alkane and a short chain alkene • 2 short chain alkanes and Hydrogen gas

Two types of cracking: • Thermal cracking (Pyrolysis) – Alkanes are heated up to a temperature between 500 to 800 degree Celsius, causing strong C-C bonds to be broken in a free radical chain mechanism. It produces a mixture of products which can be further separated by fractional distillation. • Catalytic cracking – Alkanes are passed over catalyst at 400 to 500 degree Celsius. Catalyst used is usually aluminum oxide mixed with sulfur dioxide, or zeolites.

Catalytic cracking is preferred as it is highly specific and usually yields higher percentages of hydrocarbon with between 5 and 10 carbon atoms. It is also designed to produce high proportions of branched alkanes and aromatic hydrocarbons. Furthermore, catalytic cracking occurs in lower temperature, thus being more cost efficient.

(j) Alkenes

Alkenes are unsaturated hydrocarbons and each alkene contains at least 1 C=C bond, which consists of a sigma and a pi bond. It is Trigonal planar with respect to each of the C=C carbon. Due to double bond present, there is restricted rotation about C=C bond, meaning substituent groups are fixed with respect to each other ( cis-trans isomerism. Alkenes are usually obtained from cracking of longer chain alkanes.

1) Chemical Properties and Organic reactions
The electron rich C-C bond activates alkene molecules towards electrophiles. During reactions, pi bond is broken instead of sigma bonds, hence product molecules have 2 sigma bonds. In general, alkenes undergo addition reactions in which an unsaturated compound and an attacking reagent (electrophile) combine to form a single new compound without any other products. An addition reaction involves the direct addition of attacking reagent across double or triple bond of unsaturated compound to yield a saturated one.

Combustion
Alkenes tend to burn with a more luminous and and smoky flame than corresponding alkanes as they contain a larger percentage of carbon.

Reaction with halogens
Chlorine and bromine add on to alkenes easily at room temperature to form halogenated products. This is an example of an electrophilic addition reaction. In this reaction, greenish yellow gas or reddish brown liquid would decolourise rapidly.

The reaction with bromine is a test to differentiate between alkanes and alkenes. Although both reactions would decolourise the halogens, for alkane, substitution reaction would produce the hydrogen halide, hence white fumes which turn damp blue litmus paper red would be observed. For alkenes, addition reaction only produces 1 product, hence no white fumes would be observed.

|Feature |Alkanes |Alkenes |
|Requires sunlight? |Yes |No |
|White fumes observed? |Yes |No |
|Reaction type |Substitution |Addition |
|Number of products |Numerous |1 |
|Free radicals? |Yes |No |
|Common observation |Halogens decolourise |

Reaction with steam
Alkenes can react with pressurized steam in presence of phosphoric (V) acid catalyst to produce alcohols. Conditions required would be temperature of 300 degree Celsius, pressure of 60 to 70 atm and phosphoric (V) acid catalyst. Observation would be formation of a colourless liquid that decolourises purple acidified potassium permanganate or formation of colourless liquid turn turns orange acidified potassium dichromate green. This is as alcohol reduces these oxidizing agents.

Reaction with hydrogen gas
Alkenes can react with hydrogen gas to form an alkane. Unsaturated alkene reacts with hydrogen under high temperature and pressure conditions with nickel catalyst to form saturated alkane.

Addition Polymerisation
Polymerisation is a process whereby two or more simple molecules, monomers, link together to form a large molecule consisting of repeating units of monomer called polymer. Alkene monomers undergo addition polymerisation to form a polymer, and involves monomers linking together to form polymer without any gain or loss of atoms/molecules. It is initated by free radicals.

(k) Types of fuel

Natural Gas
|Source |Alkane |
|Advantages |Clean |
|Disadvantages |Limited source, non-renewable |
|Uses |Fuel, raw material for synthesis of organic products |
|Products of combustion |Fuel, carbon dioxide, water, carbon (soot), carbon monoxide, |
| |nitrogen dioxide |

Petroleum (Liquid Hydrocarbon)
|Source |Mainly alkene |
|Advantages |Relatively clean |
|Advantages (Obtaining) |Liquid, hence easy to transport and requires little energy |
|Disadvantages |Limited source, non-renewable, air pollution and impurities |
| |such as sulfur present |
|Disadvantages (Obtaining) |Obtaining of remainder amount is difficult as remainder is |
| |often scattered around the well, thus requiring more energy to |
| |collect |
|Uses |Fuel, raw material for synthesis of organic products |
|Products of combustion |Fuel, carbon dioxide, water, carbon (soot), carbon monoxide, |
| |nitrogen dioxide, sulfur dioxide |

Gasoline
|Source |Derived from petroleum: |
| |Fractional distillation followed by cracking to give gasoline |
| |range (5 to 12 carbon) |
|Disadvantages |Dirty – when used in inefficient engines, air pollution occurs |
| |with production of pollutant |
|Uses |Fuel |

Improvements made to improve octane rating of gasoline:

• Catalytic Cracking to give branched alkanes • Alkylation which combine small hydrocarbon to give right gasoline range • Lead (Tetraethyllead) – Minute, cheap amounts can improve octane rating greatly

(l) Problems of hydrocarbons as fuels
Causes of problems • Inefficient car engines • Rapid heating-cooling cycle leading to incomplete oxidation of fuel • High temperature promoting oxidation of nitrogen to nitrogen dioxide
Oxides of nitrogen
At high temperature, nitrogen and oxygen in air combine to form nitrogen monoxide. The nitrogen monoxide can also combine with more oxygen to become nitrogen dioxide. These reactions occur naturally in lightning. However, more significant sources to these oxides are the engines of vehicles, as well as burners in power stations, factories and incinerators.
These gases irritate and damage the lungs. Nitrogen dioxide is also a cause of acid rain, as it rises and reacts with oxygen and water vapour in air to form nitric acid which is a main contributor to acid rain. The gases also react with sunlight and other pollutants to form ozone. It further reacts with unburnt hydrocarbons to form petrochemical smog. Last but not least, it attacks and depletes the ozone layer.
Carbon monoxide
Most of carbon dioxide comes from the incomplete combustion of carbon containing fuel in motor engines, which produce CO instead of carbon dioxide. CO is dangerous to human health as it has the ability to bind irreversibly to haemoglobin (and has a higher affinity to haemoglobin than oxygen), which reduces the oxygen carrying capacity of the human body, and this can potentially result in death.
Unburnt Hydrocarbons
Unburnt hydrocarbons come mainly from hydrocarbons in fuel that has not completely undergone combustion in vehicle engines (incomplete combustion). These hydrocarbons react with sunlight and oxides of nitrogen, as well as other pollutants, to form petrochemical smog.
Petrochemical Smog
Petrochemical smog is formed when unburnt hydrocarbons react with oxides of nitrogen and other pollutants. The mixture of substances in this pollutant is more harmful than its separate components. It burns the eye and is dangerous to people with breathing or heat problems. It is poisonous to plants, and damages materials like rubber.
(m) Solutions to these problems • Adjust fuel rating supply • Adjust fuel/air ratio – with a ratio of 15, there is decrease in hydrocarbon and CO formed due to increase in complete oxidation, but an increase in nitrogen dioxide levels. This shows that adjusting the ratio only trades 1 set of pollutant for another. • Adjust fuel spark timing • Adjust fuel compression ratio – Decrease in ratio decreases temperature of engine, which decreases production of oxides of nitrogen. However, this reduces the efficiency of car engines. • Tetraethyllead – Cheap and effective in minute amounts in improving octane rating of fuel, but is toxic, poisonous and renders catalytic converters ineffective. • Catalytic converter - These remove about 95% of pollutants from exhaust gases. In the first half of the converter, oxides of nitrogen react with carbon monoxide as they pass through a catalyst to produce carbon dioxide and nitrogen gas, both of which are harmless. In the second half, air enters and oxidises unburnt hydrocarbon and carbon monoxide to form carbon dioxide and water, which are non polluting as well. These are then expelled through exhaust of car. • Alternative, more efficient engines
(n) Alcohols
Alcohols are organic compounds containing carbon, hydrogen and oxygen atoms.
1) Physical Properties
Boiling point
Alcohols with less than 12 carbons exist as liquids at room temperatures, and boiling point of alcohols increase with number of C atoms. This is due to the highly polar nature of the hydroxyl group leading to possibility of hydrogen bonding between alcohol molecules. Hence, energy is required to break the hydrogen bonds, thus boiling point is relatively high.
Solubility in water
Alcohols of 1 to 3 carbon atoms are infinitely soluble in water, 4 and 5 are sparingly soluble, 6 or more are insoluble. In general, alcohol molecules can displace water molecules in the hydrogen bonding of water molecules. However, alcohol with longer carbon chain can form fewer hydrogen bonds than numerous water molecules they displace, and enthalpy considerations make longer chain alcohol less soluble than short ones, thus the trend in solubility in water.
2) Sources of Alcohol
Industrial Preparation
The industrial preparation of alcohols is from hydration of the corresponding alkenes using pressurised steam at about 60 to 70 atm, 300 degree Celsius and phosphoric (V) acid catalyst.
Fermentation of glucose by yeast – Manufacture of ethanol
Fermentation is the slow decomposition of organic compounds induced by micro organisms such as yeast and bacteria. Ethanol can be obtained from fermentation of glucose by yeast to produce ethanol and carbon dioxide. The underlying reaction is the catalytic conversion of glucose into ethanol by enzyme zymase in yeast.
The setup is maintained at around 37 degree Celsius as enzymes are higher sensitive to temperature – yeast would be destroyed, and enzymes denatured at too high temperatures, while enzymes are inactivated at too low temperatures.
3) Chemical Properties and Organic reactions
The chemical reactions of alcohols can be classified into reactions of hydroxyl group and reactions of -H on hydroxyl group. The oxygen of the hydroxyl group attached to the saturated carbon creates a partial positive centre on carbon thus making it susceptible to nucleophilic attack
Combustion
Alcohols react with oxygen to produce carbon dioxide and water.
Oxidation by acidified KMnO4 or K2Cr2O7
In this reaction, primary alcohols are oxidised to corresponding carboxylic acids. Secondary alcohols are oxidised to their corresponding ketones. In both reactions, observations are similar in that purple acidified KMnO4 would decolourise, and orange K2Cr2O7 turns green. Tertiary alcohols cannot be oxidised unless subjected to vigourous conditions.
Oxidation by oxygen
Alcohols can be oxidised by oxygen in air with aid of bacteria found in air in the process of aerobic oxidation.
Esterification
Alcohols react with carboxylic acids very slowly to produce ester and water. This is a reversible reactions carried out under reflux and concentrated sulphuric acid as catalyst. The purpose of the concentrated sulphuric acid is to acid-catalyse the process – in this case, sulphuric acid is acting as provider of hydrogen ions. Also, concentrated sulphuric acid is a dehydrating agent, hence absorbing water produced from reaction, pushing equilibrium to the right, thus forming more esters.
4) Uses of alcohols • Solvent • Fuel • Anti freeze
(o) Carboxylic Acids
All organic acids contain carboxyl group as their functional group.
1) Source of carboxylic acids
Oxidation of primary alcohols
The preparation of carboxylic acids is from oxidation of corresponding alcohols by reflux with acidified potassium dichromate (VI) solution/ acidified potassium manganate (VII) solution.
2) Physical properties
Boiling point
Carboxylic acids exist as liquids at room temperatures, and boiling point of carboxylic acids increase with number of C atoms. This is due to the highly polar nature of the OH bond, in addition to oxygen (C=O), leading to possibility of hydrogen bonding between alcohol molecules. Hence, energy is required to break the hydrogen bonds, thus boiling point is relatively high.
Boiling point of carboxylic acid is higher than that of alcohol due to extra C=O bond which allows H atom from another carboxylic acid to form hydrogen bond with it, increasing the strength of hydrogen bonding between molecules.
Solubility in water
Carboxylic acids are usually soluble in water as they dissociate in water to give hydrogen ions and corresponding carboxylate anion. In general, carboxylic acid molecules can displace water molecules in the hydrogen bonding of water molecules. However, as with alcohol, carboxylic acids with longer carbon chain can form fewer hydrogen bonds than numerous water molecules they displace, and enthalpy considerations make longer chain carboxylic acids less soluble than short ones, thus the trend in solubility in water.
3) Chemical Properties and organic reactions
The chemical reactions of alcohols can be classified into reactions of hydroxyl (-OH) group and reactions of -H on carboxyl OH in acid base reactions. The oxygen of the hydroxyl group attached to the saturated carbon creates a partial positive centre on carbon thus making it susceptible to nucleophilic attack
Acid Base Reactions
Most carboxylic acids are monobasic acids as they only contain 1 hydrogen ion per carboxyl functionality. Carboxylic acids are weak acids, as only a small percentage of acid molecules dissociate in water to give hydrogen ions. Solutions of carboxylic acids exhibit same chemical properties as other acids.
- Reaction of acid with metals
Most dilute acids react with metals above hydrogen in the reactivity series to produce hydrogen gas and a metal salt. The validity of this test can be confirmed by testing for hydrogen gas produced. If hydrogen gas is produced, burning splint put near mouth of test tube in which reaction is occurring would extinguish with a ‘pop’ sound.
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- Reaction of acid with carbonate
In this reaction, acid reacts with a carbonate to produce the resultant salt, water and carbon dioxide. The validity of this test can be confirmed by testing for carbon dioxide gas produced, in which a white precipitate would be formed when gas is bubbled into limewater. Reaction of acid with hydrogen-carbonates is the intermediate step of this reaction.
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- Reaction of acid with metal oxides and hydroxides
In this reaction, acid reacts with metal oxides and hydroxides to produce the resultant salt and water. This is a relatively slow process.
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Esterification
Alcohols react with carboxylic acids very slowly to produce ester and water. This is a reversible reactions carried out under reflux and concentrated sulphuric acid as catalyst. The purpose of the concentrated sulphuric acid is to acid-catalyse the process – in this case, sulphuric acid is acting as provider of hydrogen ions. Also, concentrated sulphuric acid is a dehydrating agent, hence absorbing water produced from reaction, pushing equilibrium to the right, thus forming more esters.
This is a condensation reaction, which involves the addition of carboxylic acid and alcohol molecules before the elimination of the water molecule which is derived from –OH group of carboxylic acid and H atom of hydroxyl group on alcohol.
Esterification is not an acid base neutralisation reaction, as the acid base reaction is between ions and thus proceeds to completion almost instantaneously, while esterification occurs between molecules very slowly as some covalent bonds in molecules need to be broken for reaction to take place.

(p) Polymers and polymerisation
A macromolecule is a long chain molecule that contains hundreds and thousands of atoms joined together by covalent bonds. It is formed by polymerisation. Polymerisation is a process whereby two or more simple molecules, monomers, link together to form a large molecule consisting of repeating units of monomer called polymer. Conditions required are high temperature and pressure as well as use of a catalyst.
1) Addition polymerisation
This occurs when many small unsaturated molecules with sae kind of functional group link together with no loss of materials.
Examples:
• Polyethene • Polychloroethene (PVC) • Polystyrene (Polyphenylethene) • Teflon (Polytetrafluoroethene)
2) Condensation polymerisation
This occurs when many small molecules which possess at least 2 functional groups in each monomer link together with the elimination of small molecules like water. This reaction does not need a double bond in reacting molecules, but require monomers to have at least 2 functional groups.
Polyester
Polyesters are made by condensation polymerisation involving 2 types of monomers, namely diol and dicarboxylic acid. Polyesters are joined together by ester linkages. 1 example is Terylene.
Polyamide
Polyamides are made by condensation polymerisation involving 2 types of monomers, namely diamine and dicarboxylic acid. Polyamides are joined together by amide linkages. 1 example is nylon, which is used to make sewing threads as it is strong.
3) Advantages and disadvantages of polymers
Due to C=C bonds being very strong, durable and difficult to break, using polymers would mean that the material would be very strong and durable.
However, the polymer is non biodegradable, and burning it could release toxic chemicals and pollutants which is harmful to the environment. Also, polymers containing chlorine gives off toxic hydrogen chloride when burning. In addition, burying it in landfills would waste precious land space.